Radiological and Chemical Fact Sheets 

to Support Health Risk Analyses for Contaminated Areas 

Prepared by 

Argonne National Laboratory 

Environmental Science Division 

John Peterson, Margaret MacDonell, 

Lynne Haroun, and Fred Monette 

In collaboration with 

U.S. Department of Energy 

Richland Operations Office 

R. Douglas Hildebrand

and 

Chicago Operations Office 

Anibal Taboas

March 2007 

3⁄4 Selenium 

3⁄4 Strontium 

3⁄4 Technetium 

3⁄4 Thorium 

3⁄4 Tin 

3⁄4 Tritium 

3⁄4 Uranium 

3⁄4 Depleted uranium (DU) 

 complements uranium sheet

3⁄4 Zirconium 

3⁄4 Carbon tetrachloride

3⁄4 Chloroform

3⁄4 Polychlorinated biphenyls (PCBs)

3⁄4 Trichloroethane

3⁄4 Trichloroethylene 

3⁄4 Cyanide 

3⁄4 Nitrate/nitrite

3⁄4 Iridium 

3⁄4 Krypton 

3⁄4 Neptunium 

3⁄4 Nickel 

3⁄4 Plutonium 

3⁄4 Polonium 

3⁄4 Potassium-40 

3⁄4 Protactinium 

3⁄4 Radium 

3⁄4 Samarium

Radiological and Chemical Fact Sheets 

to Support Health Risk Analyses for Contaminated Areas 

These fact sheets summarize health-related information for contaminants present in the environment as a 

result of past industrial activities and other releases. The objective is to provide scientific context for risk 

analyses to guide health protection measures. Geared toward an audience familiar with basic risk 

concepts, they were originally developed for the U.S. Department of Energy (DOE) Richland and 

Chicago Operations Offices to serve as an information resource for people involved in environmental 

programs. The initial set was expanded to address evolving homeland security concerns, and these 

51 radiological and chemical fact sheets also serve as a scientific information resource for the public. 

Twenty-nine radionuclide-specific fact sheets have been prepared: 

Four companion fact sheets present basic radiological concepts: 

3⁄4 Ionizing radiation 

3⁄4 Natural decay series (radium, thorium, and uranium) 

3⁄4 Transuranic radionuclides and decay series (plutonium and others) 

3⁄4 Radioactive properties, internal distribution, and risk coefficients 

Two additional sheets directly address the homeland security context for a subset of radionuclides: 

3⁄4 Radiological dispersal device 

3⁄4 Health-based radionuclide concentrations in drinking water and air

Several of the radiological fact sheets also include information for chemical toxicity, notably those for 

cadmium, chlorine, nickel, selenium, strontium, and uranium. In addition to the radiological fact sheets, 

chemical-specific fact sheets have been prepared for 7 nonradioactive metals; 5 organic compounds 

(including one class of compounds); and 2 inorganic ions (with one covering a pair of anions): 

3⁄4 Arsenic 

3⁄4 Beryllium 

3⁄4 Chromium 

3⁄4 Copper 

3⁄4 Lead (includes link to radiological information)

3⁄4 Mercury 

3⁄4 Zinc 

For cumulative risk analyses, the evaluation of interactions is framed by two fact sheets for mixtures: 

3⁄4 Mixtures concepts 

3⁄4 Arsenic, cadmium, chromium, and lead 

3⁄4 Americium 

3⁄4 Cadmium 

3⁄4 Californium 

3⁄4 Carbon-14 

3⁄4 Cesium

3⁄4 Chlorine

3⁄4 Cobalt

3⁄4 Curium 

3⁄4 Europium 

3⁄4 Iodine 
These 51 fact sheets are provided in alphabetical order in this report except DU, which follows the 

general sheet for uranium, and the two mixtures fact sheets, which are provided at the end. Each 

contaminant-specific fact sheet presents brief information on: 

3⁄4 Key properties, origin, and use 

3⁄4 General environmental levels 

3⁄4 Distribution in the body 

3⁄4 Primary health effects 

3⁄4 Values for estimating risk 

Additional information provided separately includes a summary of radionuclide morbidity and mortality

risk coefficients for key isotopes (Table 1), a summary of the radioactive properties of these isotopes 

(Table 2), and source references for the radiological fact sheets (Table 3). For the chemicals, parallel 

information (chemical toxicity values, illustration of organs affected, and references) is included within 

the individual fact sheets. 

(Note: These fact sheets have been prepared at different times to support various program needs; the 

preparation date is identified in the header of each. To open a given fact sheet, click on the link provided 

within the list on the previous page.) 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Radioactive Properties of Key Americium Isotopes 

and Associated Radionuclides 

Radiation Energy (MeV)

Isotope HalfLife 

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ)

Am-241 430 yr 3.5 α 5.5 0.052 0.033 

Am-242m 150 yr 9.8 IT 0.025 0.044 0.0051 

Am-242 16 hr 820,000 β, EC - 0.18 0.018 

Am-243 7,400 yr 0.20 α 5.3 0.022 0.055 

Np-239 2.4 days 230,000 β - 0.26 0.17 

IT = isomeric transition, EC = electron capture, Ci = curie, g = gram, and 

MeV = million electron volts; a dash means that the entry is not applicable. 

(See the companion fact sheet on Radioactive Properties, Internal 

Distribution, and Risk Coefficients for an explanation of terms and 

interpretation of radiation energies.) Americium-242 decays by two means: 

by emitting a beta particle (83%) and by electron capture (17%). Certain 

properties of americium-242 and neptunium-239 are included here because 

these radionuclides accompany the americium decays. Values are given to 

two significant figures. 

Symbol: Am

Atomic Number: 95 

(protons in nucleus)

Atomic Weight: - 

(not naturally occurring)

What Is It? Americium is a malleable, silvery white metal that tarnishes slowly in dry air at room 

temperature. Americium does not occur naturally but is produced artificially 

by successive neutron capture reactions by plutonium isotopes. There are 

sixteen known isotopes of americium and all of them are radioactive. 

(Isotopes are different forms of an element that have the same number of 

protons in the nucleus but a different number of neutrons.) Americium-241 

was first produced in 1944 in a nuclear reactor at the University of Chicago. 

Dr. Glenn Seaborg gave the new element its name in 1946 in honor of the 

continent on which it was discovered. 

Of the sixteen radioactive isotopes, only three have half-lives long enough to warrant concern at Department 

of Energy (DOE) environmental management sites: americium-241, americium-242m, and americium-243. 

The half-lives of these three 

isotopes range from 150 to 7,400 

years, while those of the other 

isotopes are less than a day. 

Americium-241 is generally the 

most prevalent isotope at DOE 

sites such as Hanford. It has a 

half-life of 430 years and decays 

by emitting an alpha particle 

with attendant gamma radiation. 

The other two isotopes typically 

represent less than a few percent 

of the total americium inventory 

at a site. Americium-242m (the 

“m” means metastable) has a 

half-life of 150 years, and it 

decays by isomeric transition. 

Americium-243 is generally not 

a major concern at DOE sites 

given its low abundance relative 

to americium-241 and low specific activity. 

Where Does It Come From? Americium is a byproduct of plutonium production activities and results 

from the successive capture of neutrons by plutonium. The most common isotope is americium-241, a decay 

product of plutonium-241. When plutonium-239 absorbs two neutrons it produces plutonium-241, which 

decays by emitting a beta particle with a fairly short half-life of 14 years to generate americium-241. 

Americium-243 is produced in a similar manner from the decay of plutonium-243, which decays by emitting 

a beta particle with a half-life of 5 hours. Successive neutron absorptions of the isotope americium-241can 

produce both americium-242m and americium-243. 

How Is It Used? The most common use of americium is in smoke detectors. These detectors rely on the 

alpha particle associated with the decay of americium-241 to ionize the air in a gap between two electrodes, 

causing a very small electrical current to flow between them. When smoke enters the space between the 

electrodes, the alpha radiation is absorbed by the soot particles, the current is interrupted, and the alarm is 

sounded. Alpha particles from smoke detectors do not themselves pose a health hazard, as they are absorbed 

in a few centimeters of air or by the structure of the detector. Americium is also used as a portable source for 

gamma radiography, for crystal research, and as target material in nuclear reactors or particle accelerators to 

produce even heavier elements. A common neutron source is composed of americium-241 and beryllium. 

The alpha particle given off during the radioactive decay of americium-241 is absorbed by beryllium-9, 

producing carbon-12 and a neutron. Such devices can be used for the nondestructive testing of machinery 

and equipment and for other industrial applications. 

Americium 
What’s in the Environment? Atmospheric testing of nuclear weapons, which ceased worldwide by 1980, 

generated most environmental americium. Accidents and other releases from weapons 

production facilities have caused localized contamination. Americium oxide is the most 

common form in the environment. Average americium-241 levels in surface soil are about 

0.01 picocuries (pCi)/g. Americium is typically quite insoluble, although a small fraction can 

become soluble through chemical and biological processes. It adheres very strongly to soil, 

with americium concentrations associated with sandy soil particles estimated to be 1,900 times higher than in 

interstitial water (the water in the pore spaces between the soil particles); it binds more tightly to loam and 

clay soils so those concentration ratios are even higher. At DOE sites such as Hanford, americium can be 

present in areas that contain waste from the processing of irradiated fuel. 

What Happens to It in the Body? Americium can be taken into the body by eating food, drinking water, 

or breathing air. Gastrointestinal absorption from food or water is a likely source of internally deposited 

americium in the general population. After ingestion or inhalation, most americium is excreted from the 

body within a few days and never enters the bloodstream; only about 0.05% of the amount taken into the 

body by ingestion is absorbed into the blood. After leaving the intestine or lung, about 10% clears the body. 

The rest of what enters the bloodstream deposits about equally in the liver and skeleton where it remains for 

long periods of time, with biological retention half-lives of about 20 and 50 years, respectively (per 

simplified models that do not reflect intermediate redistribution). The amount deposited in the liver and 

skeleton depends on the age of the individual, with fractional uptake in the liver increasing with age. 

Americium in the skeleton is deposited uniformly on cortical and trabecular surfaces of bones and slowly 

redistributes throughout the volume of mineral bone over time. 

What Are the Primary Health Effects? Americium is generally a health hazard only if it is taken into 

the body, although there is a small risk associated with the gamma rays emitted by neptunium-239, a 

radioactive decay product of americium-243. The main means of exposure are ingestion of food and water 

containing americium isotopes and inhalation of americium-contaminated dust. Ingestion is generally the 

exposure of concern unless there is a nearby source of contaminated airborne dust. Because americium is 

taken up in the body much more readily if inhaled rather than ingested, both exposure routes can be 

important. The major health concern is tumors resulting from the ionizing radiation emitted by americium 

isotopes deposited on bone surfaces and in the liver. 

 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including americium (see box at 

right). While ingestion is generally the most 

common type of exposure, the risk coefficients for 

this route are much lower than those for inhalation. 

As for other nuclides, the coefficient for tap water 

is about 80% of that shown for dietary ingestion. 

In addition to risks from internal exposures, there 

is an external gamma exposure risk associated with 

americium-243. To estimate a lifetime cancer 

mortality risk, if it is assumed that 100,000 people 

were continuously exposed to a thick layer of soil 

with an initial average concentration of 1 pCi/g 

americium-243, then 3 of these 100,000 people 

would be predicted to incur a fatal cancer. (This is 

in comparison to the 20,000 people from the group 

predicted to die of cancer from all other causes per 

the U.S. average.) This risk is largely associated 

with the gamma ray emitted by its short-lived decay product neptunium-239. The external risk for the other 

two americium isotopes is less than 10% of that for americium-243. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Recommended default 

absorption types were used for inhalation, and dietary 

values were used for ingestion. These values include the 

contributions from the short-lived americium decay 

products. Risks are for lifetime cancer mortality per unit 

intake (pCi), averaged over all ages and both genders 

(10-9 is a billionth, and 10-12 is a trillionth). Other 

values, including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Americium-241 2.4 × 10-8 9.5 × 10-11

Americium-242m 1.3 × 10-8 6.8 × 10-11

Americium-243 2.3 × 10-8 9.8 × 10-11

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Radioactive Properties of Key Cadmium Isotopes 

Radiation Energy (MeV)

Isotope Half-Life 

(yr) 

Natural 

Abundance 

(%)

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma

(γ) 

Cd-109 1.3 <<1 2,600 EC - 0.083 0.026 

Cd-113 9.3 million

billion 12 0.34 trillionth β - 0.093 - 

Cd-113m 14 <<1 240 β - 0.019 - 

EC = electron capture, Ci = curie, g = gram, and MeV = million electron volts; a dash means 

the entry is not applicable. (See the companion fact sheet on Radioactive Properties, Internal 

Distribution, and Risk Coefficients for an explanation of terms and interpretation of radiation 

energies.) Values are given to two significant figures. 

Symbol: Cd 

Atomic Number: 48

(protons in nucleus)

Atomic Weight: 112

(naturally occurring)

What Is It? Cadmium is a soft, silvery gray metal that is malleable and ductile, similar to 

zinc. When heated, it burns in air with a bright light to form the oxide CdO. In nature, 

essentially all cadmium exists as seven stable isotopes and one radioactive isotope. 

(Isotopes are different forms of an element that have the same number of protons in the 

nucleus, but a different number of neutrons.) The seven stable isotopes and their 

approximate abundances are cadmium-106 (1.3%), cadmium-108 (0.9%), cadmium-110 

(12%), cadmium-111 (13%), cadmium-112 (24%), cadmium-114 (29%), and cadmium116 (7.5%). The primary radioactive isotope, cadmium-113, comprises about 12% of 

natural cadmium and has an extremely long half-life. 

 

Nine major radioactive isotopes of cadmium exist, of which only three – cadmium-109, cadmium-113, and cadmium-113m 

– have half-lives long enough to warrant potential concern. The half-lives of the other six are less than 45 days. Cadmium109 decays by 

electron capture with 

a half-life of 

1.3 years, so any that 

was produced more 

than 20 years ago has 

long since decayed 

away. The other two 

cadmium isotopes 

decay by emitting a 

beta particle. The 

very low specific 

activity of cadmium113 limits its 

radioactive hazards. 

Thus, cadmium113m (the “m” means metastable) is the isotope of most concern at Department of Energy environmental management sites 

such as Hanford. Cadmium-113m decays by emitting a beta particle with no gamma radiation. 

Where Does It Come From? Cadmium is found in rare ores such as sphalerite and greenockite, and it is formed as a 

byproduct during production of zinc, copper, and lead. The United States is among the top ten producers, refining over 

1,000 metric tons of cadmium a year. The majority of cadmium that enters the environment is from mining, smelting, oil 

and coal combustion, and waste incineration. Cadmium-113m is produced by neutron activation of the stable isotope 

cadmium-112 and as a fission product. When a fissile nuclide such as uranium-235 fissions, it generally splits 

asymmetrically into two large fragments – fission products with mass numbers in the range of about 90 and 140 – and two 

or three neutrons. In a nuclear reactor, these neutrons can cause additional fissions (producing a chain reaction), escape 

from the reactor, or irradiate nearby materials. Cadmium is used in reactor components such as control rods and shields to 

absorb neutrons, resulting in the formation of various isotopes including cadmium-113m. The fission yield of cadmium113m is very low (about 0.0002%). This radioactive isotope is present in spent nuclear fuel and radioactive wastes 

associated with operating nuclear reactors and fuel reprocessing plants. 

 

How Is It Used? Most cadmium in the United States (about 75%) is used in nickel-cadmium batteries. It has also been 

used as an anticorrosive coating for steel and cast iron, and it is a component of certain specialty alloys. Cadmium is used 

in semiconductors (such as cadmium selenide and telluride), in dyes and pigments, as a stabilizer in plastics such as 

polyvinyl chloride, and as a neutron absorber in nuclear reactor control rods and shields. Its use in the United States has 

recently decreased by about 50% in response to environmental concerns. 

What’s in the Environment? Nonradioactive cadmium is present in U.S. soil at an average 

concentration of about 0.15 milligram per kilogram (mg/kg). Trace amounts of cadmium-113m are 

present in soil around the world from radioactive fallout due to past atmospheric weapons tests. It may 

also be present at certain nuclear facilities, such as reactors and facilities that process spent nuclear fuel. 

Cadmium is usually relatively immobile, with concentrations in sandy soil estimated to be 80 times 

higher than in interstitial water (water in the pore space between the soil particles); it is even less mobile 

Cadmium
in clay soils, with estimated concentration ratios above 500. The typical ratio of the concentration of cadmium in plants to 

that in soil is estimated at 0.15 (or 15%). (Cadmium-113m is generally not a major contaminant in groundwater at DOE 

sites like Hanford because of its relatively low concentration in wastes as a result of its low fission yield.) 

What Happens to It in the Body? Cadmium can be taken into the body by eating food, drinking water, breathing air, or 

smoking a cigarette. Children, and to a lesser extent adults, can also be exposed by ingesting soil. Gastrointestinal 

absorption from food or water is the principal source of internally deposited cadmium in the general population. 

Gastrointestinal absorption is generally quite low, with only about 5% of the amount ingested being transferred to the 

bloodstream; the unabsorbed cadmium is excreted in the feces. Cadmium absorption via inhalation exposure is higher, 

estimated at 30 to 60%. Little metabolic conversion of cadmium compounds occurs in the body, but cadmium does bind to 

proteins and other components of macromolecules. Thirty percent of cadmium that reaches the blood deposits in the liver, 

another 30% deposits in the kidneys, and the rest distributes throughout all other organs and tissues of the body (per 

simplified models that do not reflect intermediate redistribution). Absorbed cadmium is excreted primarily in the urine, but 

because daily excretion is only about 0.01% of the total body burden, the biological half-life is about 25 years. Dermal 

contact is not usually of concern because the amount absorbed is small (about 0.1%). 

What Are the Primary Health Effects? Cadmium-113m is a health hazard only if it is taken into the body. It does not 

pose an external hazard because it decays by emitting a relatively low-energy beta particle with no gamma radiation. While 

it concentrates in the liver and kidneys, cadmium can also deposit in other organs and tissues depending on its chemical 

form. The main concern is cancer induction from the beta particles associated with its radioactive decay, but cadmium also 

exhibits chemical toxicity, with health effects following inhalation exposure to high levels that can include damage to the 

respiratory system (bronchial and pulmonary irritation), headache, chest pains, muscular weakness, pulmonary edema, and 

death. Chronic exposure may result in emphysema and chronic bronchitis. Repeated low exposures may also cause 

permanent kidney damage, leading to kidney stones and other health problems. In its narrative for the cancer weight of 

evidence, the U.S. Environmental Protection Agency (EPA) states that occupational studies of cadmium smelter workers 

developing lung cancer provide limited evidence for the carcinogenicity of cadmium in humans following inhalation 

exposure, and that there is sufficient evidence of carcinogenicity in rats and mice by inhalation and intramuscular and 

subcutaneous injection. Cadmium was classified as a probable human carcinogen under the EPA 1996 cancer guidelines. 

In contrast, cadmium has not been shown to cause cancer when ingested. Information on the joint toxicity of cadmium with 

other metals is provided in the companion chemical mixtures fact sheet. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including cadmium (see box at right). While 

the coefficients for ingestion are lower than for inhalation, 

ingestion is generally the most common means of entry into 

the body. Similar to other radionuclides, the risk 

coefficients for tap water are about 80% of those for dietary 

ingestion. The EPA has also developed toxicity values to 

estimate the risk of developing cancer or other adverse 

health effects as a result of inhaling or ingesting cadmium. 

The toxicity value for estimating the risk of cancer for 

inhalation exposure is called a unit risk (UR), which is an 

estimate of the chance that a person will get cancer from 

continuous exposure to a chemical in air at a concentration 

of 1 milligram per cubic meter (mg/m3

). The value for the 

non-cancer effect following oral exposure is a reference 

dose (RfD), which is an estimate of the highest dose that can 

be taken in every day without causing an adverse noncancer effect. The UR is based on studies of 

humans exposed to cadmium in the workplace. Oral RfDs 

for food and water, shown at right, were developed using a 

toxicokinetic model that relates cadmium intake to 

concentrations in the kidney. To illustrate how the RfD is 

applied, a 150-pound (lb) person could safely ingest food 

containing 0.068 mg or drink water containing 0.034 mg of 

cadmium every day without expecting any adverse effects (2.2 lb = 1 kg, or 1,000 g, or 1 million mg). In contrast to the 

RfD, which represents a “safe daily dose” (and so is compared to the amount an individual takes in, as a ratio), the UR is 

multiplied by the air concentration to estimate the cancer risk. Using the UR, the EPA estimates that a person would have a 

one-in-a-million chance of developing cancer if exposed to air containing 0.0006 microgram (μg)/m3

 every day over a 

lifetime. (A microgram is one millionth of a gram.) 

Radiological Risk Coefficients 

The following table provides selected radiological risk 

coefficients for inhalation and ingestion. Maximum 

values are given for inhalation (no default absorption 

types were provided), and dietary values were used for 

ingestion. Risks are for lifetime cancer mortality per 

unit intake (picocurie, pCi), averaged over all ages and 

both genders (10-9 is a billionth, and 10-12 is a trillionth). 

Other values, including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1)

Ingestion 

(pCi-1)

Cadmium-109 2.0 × 10-11 4.2 × 10-12 

Cadmium-113 8.1 × 10-11 2.0 × 10-11 

Cadmium-113m 9.3 × 10-11 2.5 × 10-11 

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 

Chemical Toxicity Values 

Cancer Risk Non-Cancer Effect 

Inhalation UR Oral RfD: Water Oral RfD: Food 

1.8 per mg/m3 0.0005 mg/kg-d 0.001 mg/kg-d
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Cf 

Atomic Number: 98 

(protons in nucleus)

Atomic Weight: - 

(not naturally occurring)

Radioactive Properties of Key Californium Isotopes

Radiation Energy (MeV)

Isotope 

HalfLife 

(yr)

Specific

Activity

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Cf-248 0.91 1,600 α 6.3 0.0060 0.0013 

Cf-249 350 4.1 α 5.8 0.044 0.33 

Cf-250 13 110 α 6.0 0.0057 0.0012 

Cf-251 900 1.6 α 5.8 0.20 0.13 

Cf-252 2.6 540 α 5.9 0.0056 0.0012 

Ci = curie, g = gram, and MeV = million electron volts. (See the 

companion fact sheet on Radioactive Properties, Internal Distribution, 

and Risk Coefficients for an explanation of terms and interpretation of 

radiation energies.) About 3% of the decays of californium-252 are by 

spontaneous fission (SF), with alpha-particle decay occurring 97% of the 

time. A very small fraction (<0.1%) of the decays of californium-248, 

californium-249, and californium-250 are also by SF. Values are given to 

two significant figures. 

What Is It? Californium is a silvery-white or gray metal with a density 

somewhat greater than that of lead. Californium does not occur naturally 

but is produced artificially in nuclear reactors and particle accelerators. 

Ten isotopes of californium are known to exist and all are radioactive. 

(Isotopes are different forms of an element that have the same number of 

protons in the nucleus but a different number of neutrons.) Californium 

was the sixth transuranic element to be discovered and was first produced 

in 1950 in a cyclotron at the University of California at Berkeley by bombarding curium-242 with helium 

ions. Californium was named for the state and the University of California. 

Of the ten known isotopes, only five have half-lives long enough to warrant concern: californium-248, 

californium-249, californium-250, californium-251, and californium-252. The half-lives of these isotopes 

range from 0.91 to 900 years, 

while those of the other isotopes 

are less than two months. All 

five of these isotopes decay by 

emitting an alpha particle, and 

all but californium-248 also 

decay by spontaneous fission 

(SF), a process in which the 

atom self-disintegrates into two 

smaller atoms accompanied by 

a burst of neutrons and a release 

of energy. About 3% of the 

radioactive decays of 

californium-252 are by SF, 

while only a very small fraction 

of the decays of the other three 

isotopes are by SF. Californium 

is not a major radionuclide at 

Department of Energy legacy 

sites such as Hanford because it was not produced in large quantities. 

Where Does It Come From? Californium is a byproduct of plutonium production and can be formed 

by various neutron capture and radioactive decay routes. Californium-249 results from the beta decay of 

berkelium-249, while the heavier (higher-numbered) isotopes are produced by intense neutron irradiation, 

typically in a nuclear reactor. Californium can also be produced in particle accelerators. Although it has 

never been detected in nature, very minute amounts of californium might exist in some uranium ores and 

it has been suggested (but not confirmed) that it could be produced in stellar explosions (supernova). 

How Is It Used? The only californium isotope that has a commercial use is californium-252. Because 

this radionuclide is only available in very small quantities its uses are quite limited. Californium-252 is a 

very strong neutron emitter, with one microgram emitting 170 million neutrons per minute. Thin foils 

containing californium-252 can be used as a source of fission fragments for research purposes. 

Californium-252 can also be used as a portable neutron source to identify gold or silver ores through 

neutron activation analysis, and it can be used in moisture gauges to locate water and oil-bearing layers in 

oil wells. In addition, californium-252 is used in brachytherapy to treat various types of cancer. 

What’s in the Environment? Atmospheric testing of nuclear weapons, which ceased worldwide by 

1980, would have generated a small amount of environmental californium but fallout levels are 

extremely low. Californium is typically quite insoluble and adheres well to soil. The 

concentration in soil particles is estimated to be about 500 times higher than in interstitial water. 

Californium 
What Happens to It in the Body? Californium can be taken into the body by eating food, drinking 

water, or breathing air. Gastrointestinal absorption from food or water is a likely source of internally 

deposited californium in the general population. After ingestion or inhalation, most californium is 

excreted from the body within a few days and never enters the bloodstream; only about 0.05% of the 

amount taken into the body by ingestion is absorbed into the blood. After leaving the intestine or lung, 

about 65% of the californium that does enter the bloodstream deposits in the skeleton, 25% deposits in the 

liver, and the rest deposits in other organs or is excreted, primarily in urine. The biological half-lives in 

the skeleton and liver are about 50 and 20 years, respectively. (This information is per simplified models 

that do not reflect intermediate redistribution.) Californium in the skeleton is deposited on bone surfaces 

and slowly redistributes throughout the bone volume over time. 

 

What Is the Primary Health Effect? Californium is generally a health hazard only if it is taken into 

the body, although there is an external risk associated with the gamma rays emitted by californium-249 

and californium-251. The main means of exposure are ingestion of food and water containing 

californium isotopes and inhalation of californium-contaminated dust. Ingestion is generally the exposure 

of concern unless there is a nearby source of contaminated airborne dust. Because californium is taken up 

in the body much more readily if inhaled rather than ingested, both exposure routes can be important. 

The major health concern is cancer resulting from the ionizing radiation emitted by californium isotopes 

deposited on bone surfaces and in the liver. 

What Is the Risk? Lifetime cancer 

mortality risk coefficients have been 

calculated for nearly all radionuclides, 

including californium (see box at right).

While ingestion is generally the most 

common route of exposure, the risk 

coefficients for this route are much lower 

than those for inhalation. Similar to other 

radionuclides, the risk coefficients for tap 

water are about 80% of those for dietary 

ingestion. In addition to risks from internal 

exposures, there is an external gamma risk 

associated with exposure to californium-249 

and californium-251. To estimate a lifetime 

cancer mortality risk, if it is assumed that 

100,000 people were continuously exposed to 

a thick layer of soil with an initial average 

concentration of 1 pCi/g, then 6 of these 

100,000 people would be predicted to incur a 

fatal cancer if the soil contained 

californium-249, and 2 if it contained 

californium-251. (This is in comparison to 

the 20,000 people from the group predicted 

to die of cancer from all other causes per the 

U.S. average.) 

Californium-252, with about 3% of the decays by spontaneous fission, is a significant source of neutrons 

and gamma rays. As noted in Federal Guidance Report 12, no detailed information has been compiled on 

the radiation field (both neutron and gamma rays) associated with distributed sources of radionuclides that 

decay by spontaneous fission, such as californium-252 in soil. This is because radionuclides with a 

significant spontaneous fission yield are not typically present in the environment, or if present they are at 

extremely low concentrations. Hence, no standard values are currently available to estimate the dose or 

risk from external exposure to californium-252 that accounts for the spontaneous fission dose component.

Radiological Risk Coefficients 

This table provides selected risk coefficients for inhalation 

and ingestion. Maximum values are given for inhalation as 

no default absorption types were provided, and dietary 

values were used for ingestion. Risks are for lifetime 

cancer mortality per unit intake (pCi), averaged over all 

ages and both genders (10-9 is a billionth, and 10-12 is a 

trillionth). Other values, including for morbidity, are also 

available. Because risk values are not available for 

californium-252, these were estimated by multiplying the 

risk factors for californium-250 by the ratio of the dose 

conversion factors for californium-252 to californium-250. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Californium-248 2.4 × 10-8 3.8 × 10-11

Californium-249 4.0 × 10-8 1.2 × 10-10

Californium-250 3.5 × 10-8 8.0 × 10-11

Californium-251 4.1 × 10-8 1.3 × 10-10

Californium-252 2.1 × 10-8 4.1 × 10-11

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Symbol: C(-14) 

Atomic Number: 6 

(protons in nucleus)

Atomic Weight: 12 

(naturally occurring)

Radioactive Properties of Carbon-14

Radiation Energy (MeV)

Isotope 

HalfLife 

(yr)

Natural 

Abundance 

(%)

Specific 

Activity 

(Ci/g)

Decay

Mode Alpha 

(α)

Beta

(β)

Gamma 

(γ)

C-14 5,700 <<1 4.5 β - 0.049 - 

Ci = curie, g = gram, and MeV = million electron volts; a dash means the entry is not 

applicable. (See companion fact sheet on Radioactive Properties, Internal Distribution, 

and Risk Coefficients for explanation of terms and interpretation of radiation energies.) 

Values are given to two significant figures. 

What Is It? Carbon-14 is a naturally occurring radioactive isotope of 

carbon. (An isotope is a different form of an element that has the 

same number of protons in the nucleus but a different number of 

neutrons.) Carbon is widely distributed in nature and is present in all 

organic compounds. Natural forms include diamonds and graphite, 

which are among the hardest and softest minerals known, respectively. 

The nucleus of a carbon-14 atom contains six protons and eight 

neutrons. There are two stable (nonradioactive) isotopes of carbon: carbon-12, which has six protons and 

six neutrons, and carbon-13, which has six protons and seven neutrons. Carbon-12 comprises most 

(about 99%) of naturally occurring carbon, and carbon-13 accounts for about 1.1%. Naturally occurring 

carbon contains an extremely small fraction (about two trillionths) of radioactive carbon-14. 

There are several radioactive isotopes of carbon in addition to carbon-14. These isotopes are very shortlived – with half-lives ranging from 20 minutes for carbon-11 to less than a second – so they are not a 

health concern for 

Department of 

Energy (DOE) 

environmental 

management sites. 

The half-life of 

carbon-14 is 

about 5,700 years, 

and it decays by 

emitting a beta 

particle with no 

attendant gamma radiation to produce nitrogen-14. Carbon-14 is an important radionuclide in the lowlevel radioactive wastes previously disposed of at Hanford. 

Where Does It Come From? Carbon-14 is produced naturally in the upper atmosphere by the reaction 

of neutrons originating from cosmic rays with nitrogen and, to a lesser extent, with oxygen and carbon. 

The natural steady-state inventory of carbon-14 in the biosphere is about 300 million Ci, most of which is 

in the oceans. Large amounts of carbon-14 have also been released to the atmosphere as a result of 

nuclear weapons testing. Weapons testing through 1963 added about 9.6 million Ci, an increase of 3% 

above natural steady-state levels. Carbon-14 is also made commercially for use in medical or biological 

tracer research. Carbon-14 is produced in nuclear reactors by the capture of neutrons by nitrogen, carbon, 

or oxygen present as components of the fuel, moderator, or structural hardware. The contribution to the 

carbon-14 global inventory from commercial nuclear reactors and DOE facilities in the United States has 

been less than 600 Ci per year, or less than 1/500,000th of the natural steady-state level. Carbon-14 was 

produced at Hanford by neutron activation of carbon in graphite-moderated plutonium-production 

reactors in the 100 Area. Carbon-14 is present in the graphite moderator of these shutdown reactors and 

in certain wastes associated with previous reactor operations, as well as in wastes from ongoing 

decommissioning activities, including for spent graphite. 

How Is It Used? Two main uses of carbon-14 are in diagnostic medical procedures and radiocarbon 

dating to determine the age of previously living animals and plants. In medicine, carbon-14 can be 

injected to study abnormalities of metabolism that underlie diabetes, gout, anemia, and acromegaly (adult 

“gigantism”), and to trace the metabolism of new drugs. However, its main use to date has been to 

determine the age of fossils and other dead organic material. All living organisms absorb carbon from the 

environment, which contains carbon-12 and carbon-14 in a fixed ratio. When an organism dies, it no 

longer takes in carbon through respiration so the amount of carbon-14 will decrease at a constant rate due 

to radioactive decay, resulting in a lower ratio of carbon-14 to carbon-12 over time. Because this is 

constant in all living organisms, one can determine when an organism died by measuring the ratio of these 

Carbon-14 
 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. The recommended default 

absorption type was used for inhalation as an 

organic particulate, and the dietary value was used 

for ingestion. Risks are for lifetime cancer mortality 

per unit intake (pCi), averaged over all ages and 

both genders (10-12 is a trillionth). Other values, 

including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Carbon-14 6.5 × 10-12 1.4 × 10-12

For more information, see the companion fact sheet 

on Radioactive Properties, Internal Distribution, and 

Risk Coefficients and the accompanying Table 1. 

two isotopes. Radiocarbon dating is considered one of the most reliable means of determining the age of 

artifacts containing plant or animal matter, including some prehistoric materials up to 50,000 years old. 

What’s in the Environment? Carbon-14 is present in the atmosphere, oceans, and all organic material, 

and it behaves in the environment in the same manner as other carbon isotopes. The largest source is in 

the upper atmosphere where nitrogen interacts with neutrons from cosmic rays, with about 38,000 Ci of 

carbon-14 being produced by this process each year. Carbon-14 occurs in the ratio of 6 picocuries (pCi) 

of carbon-14 per gram of total carbon, and it is assimilated into tissues of all plants and 

animals just like other carbon isotopes. The atmospheric inventory is estimated at 

13 million Ci, and it is generally present as carbon dioxide with less than 1% in the form 

of carbon monoxide, methane, formaldehyde, and other molecules. Carbon-14 distributes 

throughout the atmosphere and surface ocean waters over a period of several years. 

Transfer to deep ocean waters proceeds much more slowly, taking hundreds to thousands of years. The 

carbon-14 concentration in the troposphere has been reported to be 3.4 pCi per kilogram of air, and its 

concentration in soil is about 0.2 pCi/g. Carbon-14 can be present at sites with graphite-moderated 

reactors. It is not a major contaminant in site groundwater due to its low leachability from graphite waste 

and limited presence in soil. Concentrations in sandy soil are estimated to be 5 times higher than in the 

interstitial water (in the pore spaces between the soil particles). Thus, carbon-14 that does leach from 

solids to soil can move downward fairly quickly with percolating water to groundwater. 

What Happens to It in the Body? Carbon-14 can be taken into the body by drinking water, eating 

food, or breathing air. Carbon-14 is present in the human body at a level of about 0.1 microcurie (or 

100,000 pCi) in adults, and it behaves in the same manner as other carbon isotopes. Most carbon-14 is 

almost completely absorbed upon ingestion, moving quickly from the gastrointestinal tract to the 

bloodstream. However, some carbon-containing compounds in food, such as cholesterol, fat-soluble 

vitamins, cellulose, and polysaccharides, may be less completely absorbed. The fractional uptake of 

carbon-14 by inhalation is strongly dependent on its chemical form. For carbon dioxide gas and organic 

compounds, essentially all inhaled carbon-14 is absorbed into the bloodstream, while for carbon 

monoxide gas the absorption fraction is about 40%. The absorption fraction for carbon-14 on inorganic 

particulate aerosols is significantly lower. The carbon-14 that enters the bloodstream after either 

ingestion or inhalation is quickly distributed to all organs and tissues of the body, as for other isotopes of 

carbon. Carbon-14 is eliminated from the body with a biological half-life of 40 days. 

What Is the Primary Health Effect? Carbon-14 poses a health hazard only if it is taken into the body, 

because it decays by emitting a weak beta particle with no gamma radiation. The beta particle emitted by 

carbon-14 has low energy and cannot penetrate 

deeply into tissue or travel far in air. Carbon-14 

behaves the same as ordinary carbon, both in the 

environment and in the human body. Hence, a 

significant fraction of the carbon-14 taken in by 

either ingestion or inhalation is absorbed into the 

bloodstream, where it is transferred to all organs of 

the body. The health hazard of carbon-14 is 

associated with cell damage caused by the ionizing 

radiation that results from radioactive decay, with 

the potential for subsequent cancer induction. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including carbon-14 (see box at 

right). Additional values are also available, 

including for inhalation of carbon-14 as a gaseous 

oxide, i.e., as carbon monoxide and carbon dioxide. 

As for other radionuclides, the risk coefficient for tap water is about 80% of that for dietary ingestion. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Cs

Atomic Number: 55 

(protons in nucleus)

Atomic Weight: 133 

(naturally occurring)

Radioactive Properties of Key Cesium Isotopes 

and an Associated Radionuclide 

Radiation Energy (MeV)

Isotope Half-Life 

Specific

Activity

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Cs-134 2.1 yr 1,300 β - 0.16 1.6 

Cs-135 2.3 million yr 0.0012 β - 0.067 - 

Cs-137 30 yr 88 β - 0.19 - 

Ba-137m

 (95%) 2.6 min 540 million IT - 0.065 0.60 

IT = isomeric transition, Ci = curie, g = gram, and MeV = million electron 

volts; a dash indicates that the entry is not applicable. (See the companion fact 

sheet on Radioactive Properties, Internal Distribution, and Risk Coefficients for 

an explanation of terms and interpretation of radiation energies.) Certain 

properties of barium-137m are included here because this radionuclide 

accompanies the cesium decays. Values are given to two significant figures. 

What Is It? Cesium is a soft, silvery white-gray metal that occurs in nature 

as cesium-133. The natural source yielding the greatest quantity of cesium is 

the rare mineral pollucite. American ores of pollucite, found in Maine and 

South Dakota, contain about 13% cesium oxide. Although it is a metal, 

cesium melts at the relatively low temperature of 28o C (82o F), so like mercury 

it is liquid at moderate temperatures. This most alkaline of metals reacts 

explosively when it comes in contact with cold water. 

There are 11 major radioactive isotopes of cesium. (Isotopes are different forms of an element that have the 

same number of protons in the nucleus but a different number of neutrons.) Only three have half-lives long 

enough to warrant concern: cesium-134, cesium-135 and cesium-137. Each of these decays by emitting a 

beta particle, and their half-lives range from about 2 to 2 million years. The half-lives of the other cesium 

isotopes are less than two 

weeks. Of these three, the 

isotope of most concern for 

Department of Energy (DOE) 

environmental management 

sites and other areas is 

cesium-137 which has a halflife of 30 years. Its decay 

product, barium-137m (the 

“m” means metastable) 

stabilizes itself by emitting 

an energetic gamma ray with 

a half-life of about 

2.6 minutes. It is this decay 

product that makes cesium an 

external hazard (that is, a 

hazard without being taken 

into the body). Cesium-135 

and cesium-134 are typically of less concern because of their radiological decay characteristics. The very 

long half-life of cesium-135 means it has a very low specific activity, and the slow decay rate combined with 

its low decay energy contribute to its low hazard. Cesium-134 has a half-life of 2.1 years and decays by 

emitting a beta particle. The relatively small amount of cesium-134 produced more than 20 years ago would 

essentially all be gone today due to radioactive decay. 

Where Does It Come From? Cesium is naturally present as the isotope 133 in various ores and to a lesser 

extent in soil. The three radioactive cesium isotopes identified above are produced by nuclear fission. When 

an atom of uranium-235 (or other fissile nuclide) fissions, it generally splits asymmetrically into two large 

fragments – fission products with mass numbers in the range of about 90 and 140 – and two or three 

neutrons. (The mass number is the sum of the number of protons and neutrons in the nucleus of the atom.) 

Cesium radionuclides are such fission products, with cesium-135 and cesium-137 being produced with 

relatively high yields of about 7% and 6%, respectively. That is, about 7 atoms of cesium-135 and 6 atoms 

of cesium-137 are produced per 100 fissions. Cesium-137 is a major radionuclide in spent nuclear fuel, highlevel radioactive wastes resulting from the processing of spent nuclear fuel, and radioactive wastes associated 

with the operation of nuclear reactors and fuel reprocessing plants. 

How Is It Used? Cesium metal is used in photoelectric cells and various optical instruments, and cesium 

compounds are used in the production of glass and ceramics. Cesium-137 is also used in brachytherapy to 

treat various types of cancer. (Brachytherapy is a method of radiation treatment in which sealed sources are 

used to deliver a radiation dose at a distance of up to a few centimeters by surface, intracavitary, or 

interstitial application.) 

Cesium
What’s in the Environment? Cesium-133 exists naturally as a stable isotope. The concentration of 

cesium in the earth’s crust is 1.9 milligrams per kilogram (mg/kg), and the concentration in seawater is about 

0.5 micrograms/kg. Cesium has been shown to biomagnify in aquatic food chains. Radioactive cesium is 

present in soil around the world largely as a result of fallout from past atmospheric nuclear 

weapons tests. The concentration of cesium-137 in surface soil from fallout ranges from 

about 0.1 to 1 picocurie (pCi)/g, averaging less than 0.4 pCi/g (or 0.3 billionth of a 

milligram per kilogram soil). Cesium is also present as a contaminant at certain locations, 

such as nuclear reactors and facilities that process spent nuclear fuel. 

Cesium is generally one of the less mobile radioactive metals in the environment. It preferentially adheres 

quite well to soil, and the concentration associated with sandy soil particles is estimated to be 280 times 

higher than in interstitial water (water in the pore space between soil particles); concentration ratios are much 

higher (about 2,000 to more than 4,000) in clay and loam soils. Thus, cesium is generally not a major 

contaminant in groundwater at DOE sites or other locations. 

What Happens to It in the Body? Cesium can be taken into the body by eating food, drinking water, or 

breathing air. After being taken in, cesium behaves in a manner similar to potassium and distributes 

uniformly throughout the body. Gastrointestinal absorption from food or water is the principal source of 

internally deposited cesium in the general population. Essentially all cesium that is ingested is absorbed into 

the bloodstream through the intestines. Cesium tends to concentrate in muscles because of their relatively 

large mass. Like potassium, cesium is excreted from the body fairly quickly. In an adult, 10% is excreted 

with a biological half-life of 2 days, and the rest leaves the body with a biological half-life of 110 days. 

Clearance from the body is somewhat quicker for children and adolescents. This means that if someone is 

exposed to radioactive cesium and the source of exposure is removed, much of the cesium will readily clear 

the body along the normal pathways for potassium excretion within several months. 

What Are the Primary Health Effects? Cesium-137 presents an external as well as internal health 

hazard. The strong external gamma radiation associated with its short-lived decay product barium-137m 

makes external exposure a concern, and shielding is often needed to handle materials containing large 

concentrations of cesium. While in the body, cesium poses a health hazard from both beta and gamma 

radiation, and the main health concern is associated with the increased likelihood for inducing cancer. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including cesium (see box at right).

While the coefficients for ingestion are somewhat 

lower than for inhalation, ingestion is generally the 

most common means of entry into the body. Similar 

to other radionuclides, the risk coefficients for tap 

water are about 80% of those for dietary ingestion. 

In addition to risks from internal exposures, there is a 

risk from external gamma exposure. Using the 

external gamma risk coefficient to estimate a lifetime 

cancer mortality risk, if it is assumed that 100,000 

people were continuously exposed to a thick layer of 

soil with an initial average concentration of 1 pCi/g 

cesium-137, then 6 of these 100,000 people would be 

predicted to incur a fatal cancer. (This is in 

comparison to about 20,000 people from the group 

predicted to die of cancer from all other causes per 

the general U.S. average.) This risk is largely 

associated with the gamma ray from barium-137m. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Recommended default 

absorption types were used for inhalation, and dietary 

values were used for ingestion. The cesium-137 values 

include the contribution from the decay product 

barium-137m. (See text at left for information on the 

risk from external exposure.) Risks are for lifetime 

cancer mortality per unit intake (pCi), averaged over 

all ages and both genders (10-12 is a trillionth). Other 

values, including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Cesium-134 1.1 × 10-11 3.5 × 10-11

Cesium-135 1.3 × 10-12 4.0 × 10-12

Cesium-137 8.1 × 10-12 2.5 × 10-11

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1.
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Cl

Atomic Number: 17 

(protons in nucleus)

Atomic Weight: 35 

(naturally occurring)

Radioactive Properties of the Key Chlorine Isotope

Radiation Energy (MeV)

Isotope 

HalfLife 

(yr)

Specific

Activity

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Cl-36 300,000 0.033 β, EC - 0.027 < 

EC =electron capture, Ci = curie, g = gram, and MeV = million 

electron volts; a dash means the entry is not applicable, and a ”<” 

means the radiation energy is less than 0.001 MeV. (See the companion 

fact sheet on Radioactive Properties, Internal Distribution, and Risk 

Coefficients for an explanation of terms and interpretation of radiation 

energies.) Chlorine-36 decays by both emitting a beta particle (98%) 

and electron capture (2%). Values are given to two significant figures. 

What Is It? Chlorine in its pure form is a greenish-yellow gas with a 

disagreeable, suffocating odor. About two and a half times denser than air, 

chlorine occurs in nature as two stable isotopes. (Isotopes are different forms 

of an element that have the same number of protons in the nucleus but a 

different number of neutrons.) Chlorine-35 is the most prevalent of these 

stable forms; it comprises about 76% of natural chlorine with chlorine-37 

accounting for the rest. 

Of the seven radioactive chlorine isotopes, only one – chlorine-36 – has a half-life long enough to warrant 

concern. The half-lives of all other chlorine isotopes are less than 1 hour. Chlorine-36 decays with a halflife of 300,000 years by emitting a 

beta particle and electron capture; 

most of the decays (98%) are by 

beta-particle emission. Chlorine-36 

is present at the Hanford Site as a 

contaminant in former plutoniumproduction reactors that are currently 

being decommissioned. The long 

half-life of chlorine-36 (with its 

subsequent low-specific activity) 

combined with the relatively low 

energy of its beta particle and small 

amount of gamma radiation limit the 

hazards associated with this radionuclide. 

Where Does It Come From? Chlorine is a very reactive element that does not occur uncombined in 

nature. It is usually found bound with elements such as sodium, potassium, and magnesium. The most 

common chlorine compound is sodium chloride (table salt), which is present in seawater, salt wells, and large 

salt deposits, often in association with other chlorides. Chlorine is produced commercially by the electrolysis 

of sodium chloride, and it can also be produced by oxidizing the hydrogen chloride in hydrochloric acid. 

Chlorine-36 is generated in the atmosphere by the spallation of argon-36 with cosmic ray protons, and in soil 

and rock by neutron activation of chlorine-35. Large amounts of chlorine-36 were produced by irradiating 

seawater during certain nuclear weapons tests conducted between 1952 and 1958. 

The graphite used as a neutron moderating material in the plutonium-production reactors in the 100 Area of 

the Hanford Site was treated with chlorine gas at high temperatures to remove the impurity boron (which is a 

good neutron absorber). A small amount of chlorine remained in the graphite that was charged to the nuclear 

reactors. When a fissile nuclide such as an atom of uranium-235 fissions, it generally splits asymmetrically 

into two large fragments – fission products with mass numbers in the range of about 90 and 140 – and two or 

three neutrons. (The mass number is the sum of the number of protons and neutrons in the nucleus of the 

atom.) These neutrons can cause additional fissions (producing a chain reaction), escape from the reactor, or 

irradiate nearby materials. The chlorine-35 remaining in the graphite moderator absorbed neutrons to 

become chlorine-36. Thus, chlorine-36 is present in the graphite moderator of these shutdown reactors and 

in certain wastes associated with previous reactor operations as well as in wastes from ongoing 

decommissioning activities, including spent graphite. 

How Is It Used? Chlorine has a number of industrial and commercial uses. Its primary use is as a 

disinfectant, and it has been used for this purpose against a wide range of life-threatening infections, viruses, 

and bacteria for more than 150 years. Chlorine is used to disinfect water for both drinking and swimming, 

and to kill harmful levels of bacteria such as Salmonella and E. coli during food processing in restaurants and 

in meat and poultry packaging plants. It is also used in the manufacture of bleaching powder, household 

cleansers, dyes, explosives, textiles, pharmaceuticals, synthetic rubber, paper and petroleum products, 

Chlorine 
plastics, and poisonous gases. Chlorinated hydrocarbons have been used extensively as pesticides, and their 

long persistence has made them troublesome environmental pollutants. The freon refrigerants are 

hydrocarbons that have been reacted with chlorine and fluorine, and carbon tetrachloride and 

trichloroethylene are two common solvents. 

What’s in the Environment? Chlorine is present in crustal rock at a concentration of about 

170 milligrams per kilogram (mg/kg), and its concentration in seawater is about 20 g/liter. In pure water, 

chlorine forms elemental chlorine (Cl2), chloride ion (Cl-

) and hypochlorous acid (HOCl). Although it does 

preferentially adhere to soil particles, it is one of the more mobile radionuclides in soil and 

can move downward with percolating water to underlying soil and groundwater. 

Chlorine-36 is generally not a major contaminant in groundwater at contaminated sites, 

mainly because of its limited presence in waste and soil. Its mobility combined with its 

long half-life makes chlorine-36 a radionuclide of potential concern for long-term 

management for wastes associated with decommissioning plutonium-production reactors 

(such as at the DOE Hanford site; chlorine-36 is associated with the graphite-moderated reactors). 

What Happens to It in the Body? Chlorine can be taken into the body by eating food, drinking water, or 

breathing air; small amounts could be absorbed through the skin. Gastrointestinal absorption from food or 

water is the principal source of internally deposited chorine in the general population. Chorine as chloride is 

an essential nutrient in the human diet and is necessary for healthy nervous and digestive systems. Once 

taken in, chlorine-36 behaves in the body the same way as other chlorine isotopes. Chlorine is almost 

completely absorbed upon ingestion, moving quickly from the gastrointestinal tract into the bloodstream. 

The chlorine-36 that enters the bloodstream after ingestion or inhalation is quickly distributed to all organs 

and tissues of the body. Chlorine-36 is eliminated from the body with a biological half-life of 10 days. 

 

What Are the Primary Health Effects? Chlorine is a health hazard only if it is taken into the body. 

External gamma exposure is not a concern because chlorine-36 decays by emitting a relatively low-energy 

beta particle with only a small amount of gamma 

radiation. While in the body, chlorine poses a hazard 

from the beta particles and gamma radiation. The main 

concern is the increased likelihood of cancer induction. 

In addition to its radiological effects, chlorine is a very 

toxic gas, and acute exposures to high levels can cause 

respiratory distress and death. Chronic ingestion of 

chlorine has been shown to decrease organ and body 

weights in animals, notably at high doses. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including chlorine (see box at right). 

While the coefficients for ingestion are lower than for 

inhalation, ingestion is generally the most common 

means of entry into the body. Similar to other 

radionuclides, the risk coefficients for tap water are about 

75% of those for dietary ingestion. The EPA uses a 

reference dose (RfD) for estimating the non-cancer 

health effects from oral exposure, which is an estimate of 

the highest dose that can be ingested every day over a 

lifetime without causing an adverse effect. The RfD for 

chlorine, shown at right, was developed by analyzing the biological effects of test animals given relatively 

large amounts of chlorine and then adjusting and normalizing those results to a mg/kg-day basis for humans. 

A toxicity value for non-cancer effects following chronic inhalation exposure to chlorine is not defined. 

Radiological Risk Coefficients

This table provides selected risk coefficients for 

inhalation and ingestion. Maximum values are 

given for inhalation (no default absorption types 

were provided), and dietary values were used for 

ingestion. Risks are for lifetime cancer mortality 

per unit intake (picocuries, pCi), averaged over all 

ages and both genders (10-9 is a billionth, and 10-12

is a trillionth). Other values, including for 

morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Chlorine-36 9.6 × 10-11 2.9 × 10-12

For more information, see the companion fact sheet 

on Radioactive Properties, Internal Distribution, 

and Risk Coefficients and accompanying Table 1. 

Chemical Toxicity Value 

Non-Cancer Effect: Oral RfD

0.1 mg/kg-day
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Radioactive Properties of Key Cobalt Isotopes 

Radiation Energy (MeV) Isotope HalfLife 

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Co-57 270 days 8,600 EC - 0.019 0.13 

Co-60 5.3 yr 1,100 β - 0.097 2.5 

EC = electron capture, Ci = curie, g = gram, and MeV = million electron 

volts; a dash means the entry is not applicable. (See the companion fact 

sheet on Radioactive Properties, Internal Distribution, and Risk Coefficients 

for an explanation of terms and interpretation of radiation energies.) Values 

are given to two significant figures. 

Symbol: Co

Atomic Number: 27 

(protons in nucleus)

Atomic Weight: 59 

(naturally occurring)

What Is It? Cobalt is a hard, silvery-white metal that occurs in nature as 

cobalt-59. Cobalt is a constituent of the minerals cobaltite, smaltite, erythrite, 

and other ores, and it is usually found in association with nickel, silver, lead, 

copper, and iron. Pure cobalt metal is prepared by reducing its compounds 

with aluminum, carbon, or hydrogen. It is similar to iron and nickel in its 

physical properties. Cobalt has relatively low strength and little ductility at 

normal temperatures and is a component of several alloys. 

There are nine major 

radioactive cobalt isotopes. 

(Isotopes are different forms 

of an element that have the 

same number of protons in 

the nucleus but a different 

number of neutrons.) Of 

these, only cobalt-57 and 

cobalt-60 have half-lives 

long enough to warrant 

concern. The half-lives of all 

other isotopes are less than 

80 days. Cobalt-57 decays 

with a half-life of 270 days by electron capture and cobalt-60 decays with a half-life of 5.3 years by 

emitting a beta particle with two energetic gamma rays; the combined energy of these two gamma rays is 

2.5 MeV (one has an energy of 1.2 MeV and the other has an energy of 1.3 MeV). Cobalt-60 is the 

isotope of most concern at Department of Energy (DOE) environmental management sites such as 

Hanford, for the cobalt-57 produced more than 20 years ago has long since decayed away. The two 

energetic gamma rays that accompany the radioactive decay of cobalt-60 make this isotope an external 

hazard (that is, it can be hazardous without being taken into the body). 

Where Does It Come From? Cobalt is naturally present as the isotope 59 in various ores and to a 

lesser extent in soil. Cobalt-60 is produced by neutron activation of components in nuclear reactors; it 

can also be produced in a particle accelerator. When an atom of uranium-235 (or other fissile nuclide) 

fissions, it generally splits asymmetrically into two large fragments – fission products with mass numbers 

in the range of about 90 and 140 – and two or three neutrons. (The mass number is the sum of the number 

of protons and neutrons in the nucleus of the atom.) These neutrons can cause additional fissions 

(producing a chain reaction), escape from the reactor, or irradiate nearby materials. A number of reactor 

components are made of various alloys of steel that contain chromium, manganese, nickel, iron and 

cobalt, and these elements can absorb neutrons to produce radioactive isotopes, including cobalt-60. 

Cobalt-60 is a radionuclide of concern in spent nuclear fuel (as a component of the fuel hardware) and in 

the radioactive wastes associated with nuclear reactors and fuel reprocessing plants. 

How Is It Used? Cobalt is used as a component of several alloys, including carboloy and stellite that 

are used to make very hard cutting tools. Cobalt is also used in some stainless steels. Alnico, an alloy of 

aluminum, nickel, cobalt, and other metals, is used to make high-strength, permanent magnets. Cobalt is 

also used in electroplating to give a hard surface that is resistant to oxidation, and as a blue colorant in 

pottery enamels and glass. High-energy gamma rays emitted during the radioactive decay of cobalt-60 

can be used to detect flaws in metal components and in brachytherapy to treat various types of cancer. 

(Brachytherapy is a method of radiation treatment in which sealed sources are used to deliver a radiation 

dose at a distance of up to a few centimeters by surface, intracavitary, or interstitial application.) 

Cobalt
What’s in the Environment? Cobalt-59 is present in soil as a stable isotope at a concentration of about 

1 to 2 milligram per kilogram (mg/kg). Trace amounts of cobalt-60 are also present around the globe 

from radioactive fallout as a result of past atmospheric weapons tests. It may also be present 

as a contaminant at certain locations, such as in nuclear reactors and facilities that process 

spent nuclear fuel, principally in the hardware associated with the spent fuel. Transport of 

cobalt in the environment is strongly influenced by its chemical form. It is generally one of 

the less mobile radioactive metals in soil, although certain forms can move downward with 

percolating water into underlying layers of soil. Under certain conditions, mobility can be enhanced 

based on the form and environmental setting, for example, when complexes form with cyanide or 

ferrocyanide. In other settings, cobalt has been found to preferentially adhere to soil. For sandy soil, the 

concentration in soil particles is estimated to be about 60 times higher than in water between the soil 

particles, and cobalt binds even more tightly to loam where the estimated concentration ratio is 1,300. 

What Happens to It in the Body? Cobalt can be taken into the body by eating food, drinking water, or 

breathing air. Gastrointestinal absorption from food or water is the principal source of internally 

deposited cobalt in the general population. Estimates of the gastrointestinal absorption of cobalt range 

from 5 to 30%, depending on the chemical form and amount ingested; 10% is a typical value for adults 

and 30% for children. Cobalt is an essential element found in most body tissues, with the highest 

concentration in the liver. Vitamin B12 is a cobalt-containing vitamin essential for red blood cell 

formation in humans, and the intestinal absorption of cobalt in this vitamin is high. Fifty percent of 

cobalt that reaches the blood is excreted right away, mainly in urine; 5% deposits in the liver, and the 

remaining 45% deposits evenly in other tissues of the body. Of the cobalt that deposits in the liver and 

other tissues, 60% leaves the body with a biological half-life of 6 days and 20% clears with a biological 

half-life of 60 days; the last 20% is retained much longer, with a biological half-life of 800 days. On the 

basis of animal studies, retention of cobalt was determined to be the same for all age groups. Inhaled 

cobalt oxide moves from the lung to body tissues quite readily. 

What Are the Primary Health Effects? Cobalt-60 poses both an internal and external hazard, and the 

main health concern is associated with the increased likelihood of cancer. External exposure is a concern 

because of the strong external gamma radiation, and shielding is often needed to handle wastes and other 

materials with high concentrations of this isotope. 

Inside the body, cobalt presents a hazard from 

both beta and gamma radiation. 

 

What Is the Risk? Lifetime cancer mortality 

risk coefficients have been calculated for nearly 

all radionuclides, including cobalt (see box at 

right). While the coefficients for ingestion are 

somewhat lower than for inhalation, ingestion is 

generally the most common means of entry into 

the body. Similar to other radionuclides, the risk 

coefficients for tap water are about 70% of those 

for dietary ingestion. In addition to the risk from 

internal exposure, there is a risk from external 

gamma exposure. Using the external gamma risk 

coefficient to estimate a lifetime cancer mortality 

risk, if it is assumed that 100,000 people were 

continuously exposed to a thick layer of soil with 

an initial average concentration of 1 pCi/g cobalt-60, then 6 of these 100,000 people would be predicted 

to incur a fatal cancer. (This is compared to the 20,000 people from the group predicted to die of cancer 

from all other causes per the U.S. average.) The external risk for cobalt-57 is less than 1% of this risk. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Recommended default 

absorption types were used for inhalation, and dietary 

values were used for ingestion. (See text at left for 

information on the external gamma exposure.pathway) 

Risks are for lifetime cancer mortality per unit intake 

(pCi), averaged over all ages and both genders (10-9 is 

a billionth, and 10-12 is a trillionth). Other values, 

including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Cobalt-57 1.7 × 10-12 9.0 × 10-13

Cobalt-60 3.0 × 10-11 1.4 × 10-11

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Symbol: Cm

Atomic Number: 96 

(protons in nucleus)

Atomic Weight: - 

(not naturally occurring)

Radioactive Properties of Key Curium Isotopes 

and Associated Radionuclides 

Radiation Energy (MeV)

Isotope Half-Life 

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Cm-242 160 days 3,400 α 6.1 0.010 0.0018 

Cm-243 29 yr 52 α 5.8 0.14 0.13 

Cm-244 18 yr 82 α 5.8 0.086 0.0017 

Cm-245 8,500 yr 0.17 α 5.4 0.065 0.096 

Cm-246 4,700 yr 0.31 α 5.4 0.0080 0.0015 

Cm-247 16 million yr 0.000094 α 4.9 0.021 0.32 

Pu-243 5.0 hr 2.6 million β - 0.17 0.026 

Cm-248 340,000 yr 0.0043 α 4.7 0.0060 0.0012 

Cm-250 6,900 yr 0.21 α, β 1.3 0.0016 - 

Pu-246 

 (25%) 11 days 49,000 β - 0.13 0.14 

Bk-250 

 (14%) 3.2 hr 3.9 million β - 0.29 0.89 

Am-246 

 (25%) 39 min 20 million β - 0.66 0.70 

Ci = curie, g = gram, and MeV = million electron volts; a dash indicates the entry is 

not applicable. (See the companion fact sheet on Radioactive Properties, Internal 

Distribution, and Risk Coefficients for an explanation of terms and interpretation of 

radiation energies.) About 8% of the decays of curium-248 are by spontaneous 

fission (SF), with alpha-particle decay occurring 92% of the time. 61% of the 

curium-250 decays are by SF, with alpha-particle emission (25%) and beta-particle 

emission (14%) accounting for the remainder. Certain properties of plutonium-243, 

plutonium-246, berkelium-250, and americium-246 are included here because these 

radionuclides accompany the curium. Values are given to two significant figures. 

What Is It? Curium is a hard, brittle, silvery metal that tarnishes slowly in dry 

air at room temperature. Curium does not occur naturally; it is typically 

produced artificially in nuclear reactors through successive neutron captures by 

plutonium and americium isotopes. Sixteen isotopes of curium are known to 

exist, and all are radioactive. (Isotopes are different forms of an element that 

have the same number of protons in the nucleus but a different number of 

neutrons.) Curium was first produced in 1944 by bombarding plutonium-239 

with alpha particles in a cyclotron at the University of California at Berkeley. 

Curium was isolated in visible amounts as the hydroxide in 1947 and is named in honor of Pierre and Marie 

Curie, who pioneered the study of radioactivity. 

Eight of the sixteen curium 

isotopes have half-lives 

greater than one month. 

Curium-243 and curium-244 

are the two isotopes of most 

concern at Department of 

Energy (DOE) environmental 

management sites such as 

Hanford. The curium-242 

produced more than 20 years 

ago has essentially all decayed 

away, and the low specific 

activities of the other curium 

isotopes limit their 

radiological hazards. In 

addition, the longer-lived 

isotopes typically represent 

much less than 1% of the 

curium inventory at a site. 

Curium generally decays to 

plutonium by emitting an 

alpha particle; gamma 

radiation is associated with 

some of these decays. A 

relatively small percentage 

(14%) of curium-250 decays 

are by beta-particle emission 

to berkelium-250. Curium248 and curium-250 also 

decay by spontaneous fission 

(SF), a process in which the atom self-disintegrates into two smaller atoms accompanied by a release of energy. 

(A very small fraction of curium-242, curium-244, and curium-246 decays are also by SF.) 

Where Does It Come From? Although the presence of natural curium has never been detected, minute 

amounts may exist in some uranium ores. Curium is a byproduct of plutonium production activities and results 

from the successive capture of neutrons by plutonium and americium, generally in nuclear reactors. 

How Is It Used? Curium has few uses outside of research activities, and it is only available in extremely small 

quantities. Curium isotopes can be used without heavy shielding as sources of thermoelectric power in satellites 

and crewless space probes. Curium-242 has been used in isotopic power generators because it produces about 

3 watts of heat energy (from radioactive decay) per gram. Curium-242 was used on lunar missions to bombard 

the moon’s soil with alpha particles to determine what it was made of. Instruments analyzed the characteristics 

of the scattered alpha particles from the moon’s surface, from which it was determined that lunar soil was 

Curium 
similar in composition to basalt, a common terrestrial volcanic rock. The high specific activity of curium-242 

coupled with its low external hazard made this isotope an ideal choice for such an application. 

What’s in the Environment? Atmospheric testing of nuclear weapons, which ceased worldwide by 1980, 

generated most environmental curium. Accidents and other releases from weapons production 

facilities have caused localized contamination. Curium oxide is the most common form in the 

environment. Curium is typically quite insoluble and adheres very tightly to soil particles. The 

concentration of curium in sandy soil particles is estimated to be about 4,000 times higher than in 

interstitial water (in pore spaces between soil particles), and it binds even more tightly to loam 

soil where concentration ratios are even higher (18,000). 

What Happens to It in the Body? Curium can be taken into the body by eating food, drinking water, or 

breathing air. Gastrointestinal absorption from food or water is the most likely source of any internally 

deposited curium in the general population. After ingestion, most curium is excreted from the body within a few 

days and never enters the bloodstream; only about 0.05% of the amount ingested is absorbed into the 

bloodstream. Of the curium that reaches the blood, about 45% deposits in the liver where it is retained with a 

biological half-life of 20 years, and 45% deposits in bone where it is retained with a biological half-life of 

50 years (per simplified models that do not reflect intermediate redistribution). Most of the remaining 10% is 

directly excreted. Curium in the skeleton is deposited mainly on the endosteal surfaces of mineral bone and 

only slowly redistributes throughout the bone volume. 

 

What Are the Primary Health Effects? Curium is generally a health hazard only if it is taken into the body; 

however, a small external risk is associated with curium-243, curium-245, curium-247, and curium-250 (note 

that the SF contribution from curium-250 is not quantified). People can be exposed by ingesting contaminated 

food or water or by inhaling contaminated dust. Ingestion is generally the exposure route of concern unless a 

nearby source of dust contamination exists. Because curium is absorbed within the body much more readily if 

inhaled rather than ingested, both exposure routes can be important. The main health concern is bone tumors 

resulting from ionizing radiation emitted by curium 

isotopes deposited on bone surfaces. 

 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including curium (see box at right).

Although ingestion is generally the most common 

exposure, these risk coefficients are much lower than 

those for inhalation. As for other nuclides, the risk 

coefficient for tap water is about 80% of that shown 

for dietary ingestion. In addition to risks from internal 

exposures, a risk from external gamma exposure is 

associated with curium-243, curium-245, curium-247, 

and curium-250. Using the external gamma risk 

coefficients to estimate lifetime cancer mortality risks, 

if it is assumed that 100,000 people were continuously 

exposed to a thick layer of soil with an initial average 

concentration of 1 pCi/g, then 1 person would be 

predicted to incur a fatal cancer if the soil contained 

curium-243, 1 if it contained curium-245, 6 if it 

contained curium-247, and 7 if it contained 

curium-250 (from its short-lived decay products). 

(This compares to 20,000 people from the group 

predicted to die of cancer from all other causes per the 

U.S. average.) The external risk coefficients for the 

other curium isotopes are less than 1% of those for 

these four. About 8% of curium-248 decays and 61% of curium-250 decays are by SF. No detailed information 

has been compiled on the SF radiation field (neutron and gamma rays) associated with distributed sources of 

radionuclides in soil. This is because radionuclides with a significant SF yield are not usually found in the 

Radiological Risk Coefficients

This table provides selected risk coefficients for inhalation 

and ingestion. Recommended default absorption types 

were used for inhalation, and dietary values were used for 

ingestion. Risks are for lifetime cancer mortality per unit 

intake (picocurie, pCi), averaged over all ages and both 

genders (10-9 is a billionth, and 10-12 is a trillionth). Other 

values, including for morbidity, are also available. 

Because risk values are not available for curium-248 and 

curium-250, these were estimated by multiplying the risk 

factors for curium-246 by the ratios of the dose conversion 

factors for curium-248 and curium-250 to curium-246. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Curium-242 1.4 × 10-8 3.2 × 10-11

Curium-243 2.4 × 10-8 8.5 × 10-11

Curium-244 2.3 × 10-8 7.5 × 10-11

Curium-245 2.4 × 10-8 9.5 × 10-11

Curium-246 2.4 × 10-8 9.3 × 10-11

Curium-247 2.2 × 10-8 9.1 × 10-11

Curium-248 8.8 × 10-8 3.4 × 10-10

Curium-250 5.0 × 10-7 2.0 × 10-9

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
environment, and if present they are at extremely low levels. Thus, no standard values are available to estimate 

the dose or risk from external exposure to curium-248 or curium-250 that accounts for the SF dose component. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Eu

Atomic Number: 63 

(protons in nucleus)

Atomic Weight: 152 

(naturally occurring)

Radioactive Properties of Key Europium Isotopes 

Radiation Energy (MeV)

Isotope Half-Life

(yr)

Specific

Activity

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Eu-150 34 70 EC - 0.044 1.5 

Eu-152 13 180 β, EC - 0.14 1.2 

Eu-154 8.8 270 β - 0.29 1.2 

Eu-155 5.0 470 β - 0.063 0.061 

EC = electron capture, Ci = curie, g = gram, and MeV = million electron 

volts; a dash means the entry is not applicable. (See the companion fact sheet 

on Radioactive Properties, Internal Distribution, and Risk Coefficients for an 

explanation of terms and interpretation of radiation energies.) Europium-152 

decays both by emitting a beta particle (28%) and electron capture (72%). A 

second isotope of europium-150 having a half-life of 13 hours also exists. 

Values are given to two significant figures.

What Is It? Europium is a silvery-white metal. It is the softest, least 

dense, and most volatile member of the lanthanide series, and it ignites in 

air at high temperatures (150 to 180°C). In nature, europium occurs as two 

stable isotopes. (Isotopes are different forms of an element that have the 

same number of protons in the nucleus but a different number of neutrons.) 

Europium-153 accounts for 52% of natural europium, and europium-151 

makes up the remaining 48%. 

Of the fourteen major radioactive isotopes, only four have half-lives long enough to warrant potential 

concern. The half-lives of all other europium isotopes are less than four months. Of these four longerlived isotopes, three (europium-152, europium-154, and europium-155) are produced by the fissioning of 

uranium and plutonium and 

are present at Department 

of Energy (DOE) environmental management sites 

such as Hanford. These 

three isotopes have halflives ranging from 5 to 

13 years, and they decay by 

emitting a beta particle. A 

fraction (28%) of the 

europium-152 decays is by 

electron capture. A 

significant amount of 

energy in the form of 

gamma rays accompanies 

the decays of europium-152 

and europium-154. 

Where Does It Come From? Europium is found in a variety of ores, primarily bastnasite, monazite, 

and xenotime. These ores contain different mixes of rare earth metals, which are the elements from 

lanthanum through lutetium in the periodic table. Europium generally makes up less than 0.2% of this 

mixture. China currently produces the vast majority of the rare earth metals, totaling about 70,000 metric 

tons (MT) annually. The United States comes in a distant second, producing around 5,000 MT annually, 

and europium makes up only a small fraction of this amount. 

While europium-152, europium-154, and europium-155 are produced primarily as fission products, 

europium-152 can also be produced by neutron activation of nuclear reactor control rods. When a fissile 

nuclide such as an atom of uranium-235 fissions, it generally splits asymmetrically into two large 

fragments – which can include the three europium isotopes – and two or three neutrons. The fission yield 

of europium-155 is about 0.03% while the yield of the other two isotopes is much lower. That is, about 3 

atoms of europium-155 are produced per 10,000 fissions. In order to control this fission reaction, 

isotopes that can absorb excess neutrons are used in nuclear reactor control rods. Because europium-151 

is a very good neutron absorber, it is often used in these control rods. Neutron activation of this stable 

isotope produces europium-152. 

How Is It Used? The primary use of europium is in nuclear reactor control rods, because of its 

effectiveness in absorbing neutrons. Other uses have been limited because it is rare and thus very 

expensive. Europium-doped plastics have been used as laser materials, and europium oxide serves as a 

phosphor activator. For example, europium has been used to activate yttrium vanadate for its use in the 

red phosphors of color television tubes. 

Europium 
What’s in the Environment? Europium is present in the earth’s crust at a concentration of about 

1.8 milligrams per kilogram (mg/kg), while its concentration in seawater is about 0.00013 micrograms per 

liter. Trace amounts of europium-152, europium-154, and europium-155 are present in 

soil around the globe from radioactive fallout. It can also be present at certain nuclear 

facilities, such as reactors and spent fuel reprocessing plants. Europium is generally 

one of the more immobile radioactive metals in the environment. It preferentially 

adheres fairly tightly to soil, and the concentration associated with soil particles is 

estimated to be about 240 times higher than in interstitial water (the water in the pore 

space between soil particles). Thus, europium is generally not a major contaminant in groundwater at 

DOE sites. 

What Happens to It in the Body? Europium can be taken into the body by eating food, drinking 

water, or breathing air. Gastrointestinal absorption from food or water is the principal source of internally 

deposited europium in the general population. Europium is not well absorbed into the body after intake, 

with only about 0.05% of the amount ingested being absorbed into the bloodstream through the digestive 

tract. Of the europium that reaches the blood, 40% is deposited in the liver, and another 40% is deposited 

on the surface of the bone, where it can irradiate the bone-forming cells; this deposited europium is 

retained in the body with a biological half-life of almost 10 years (3,500 days); an additional 6% of the 

absorbed europium is deposited in the kidneys, where it is retained with a short biological half-life of 

10 days (per simplified models that do not reflect intermediate redistribution). The remainder of the 

absorbed europium is excreted. 

 

What Are the Primary Health Effects? Europium poses an external as well as an internal health 

hazard. The strong gamma radiation associated with europium-152 and europium-154 makes external 

exposure to these two isotopes a concern. While in the body, europium poses a health hazard from both 

the beta particles and gamma rays, and the main health concern is associated with the increased likelihood 

of inducing cancer in the liver and bone. 

What Is the Risk? Lifetime cancer mortality 

risk coefficients have been calculated for nearly 

all radionuclides, including europium (see box at 

right). While the coefficients for ingestion are 

somewhat lower than for inhalation, ingestion is 

generally the most common means of entry into 

the body. Similar to other radionuclides, the risk 

coefficients for tap water are about 70% of those 

for dietary ingestion. 

In addition to the risks from internal exposure, a 

risk from external gamma exposure is associated 

with europium-152 and europium-154. If it is 

assumed that 100,000 people were continuously 

exposed to a thick layer of soil with an initial 

average concentration of 1 picocurie (pCi)/g, then 

the number of fatal cancers estimated for these 

100,000 people is 7 for exposure to europium-152 

and 5 for exposure to europium-154. (This can 

be compared to the 20,000 people from this group 

who would be predicted to die of cancer from all 

other causes per the U.S. average.) 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Maximum values are given 

for inhalation since no default absorption types were 

provided, and dietary values were used for ingestion. 

Risks are for lifetime cancer mortality per unit intake 

(pCi), averaged over all ages and both genders (10-9 is 

a billionth, and 10-12 is a trillionth). Other values, 

including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(per pCi-1) 

Ingestion 

(per pCi-1) 

Europium-150 2.1 × 10-10 3.6 × 10-12 

Europium-152 1.5 × 10-10 5.0 × 10-12 

Europium-154 1.7 × 10-10 8.5 × 10-12 

Europium-155 1.7 × 10-11 1.6 × 10-12 

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Symbol: I

Atomic Number: 53 

(protons in nucleus)

Atomic Weight: 127 

(naturally occurring)

Radioactive Properties of Key Iodine Isotopes 

Radiation Energy (MeV) Isotope Half-Life Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

I-129 16 million yr 0.00018 β - 0.064 0.025 

I-131 8.0 days 130,000 β - 0.19 0.38 

EC = electron capture, Ci = curie, g = gram, and MeV = million electron volts; a 

dash means the entry is not applicable. (See the companion fact sheet on Radioactive 

Properties, Internal Distribution, and Risk Coefficients for an explanation of terms 

and interpretation of radiation energies.) Values are given to two significant figures.

What Is It? Iodine is a bluish-black, lustrous solid that mainly occurs in 

nature as stable iodine-127. A small amount of radioactive iodine-129 is 

produced naturally in the upper atmosphere by the interaction of highenergy particles with xenon. Iodine volatilizes at ambient temperatures into 

a pretty blue-violet gas with an irritating odor. Iodine exhibits some metallike properties and is only slightly soluble in water. It occurs in nature as 

iodide ions, and it is in this form that it is taken into our bodies. 

Of the fourteen major radioactive isotopes of iodine, only iodine-129 has a half-life sufficiently long to 

warrant concern for Department of Energy (DOE) environmental management sites such as Hanford. 

(Isotopes are different 

forms of an element 

that have the same 

number of protons in 

the nucleus but a 

different number of 

neutrons.) Iodine-129 

decays by emitting a 

beta particle with a 

half-life of about 

16 million years; the 

half-lives of all other 

iodine radionuclides are less than 60 days. The very long half-life of iodine-129 (with its subsequent lowspecific activity) combined with the low energy of its beta particle and minimal gamma radiation limit the 

hazards of this radionuclide. Iodine-131 has a short half-life (8 days) and is not generally a major isotope 

of concern for DOE environmental management sites. However, information is included here for this 

radionuclide because it was released during past operations of nuclear reactors at Hanford. 

Where Does It Come From? Stable iodine (iodine-127) is naturally present in seaweeds, sponges, and 

other materials. Radioactive isotopes of iodine are produced by nuclear fission. When an atom of 

uranium-235 (or other fissile nuclide) fissions, it generally splits asymmetrically into two large fragments 

– fission products with mass numbers in the range of about 90 and 140 – and two or three neutrons. (The 

mass number is the sum of the number of protons and neutrons in the nucleus of the atom.) Iodine-129 

and iodine-131 are two such products. The fission yield of iodine-129 is about 1% and the yield of 

iodine-131 is close to 3%. That is, about one atom of iodine-129 and three atoms of iodine 131 are 

produced per 100 fissions. Iodine-129 is present in spent nuclear fuel, high-level radioactive wastes 

resulting from processing spent nuclear fuel, and radioactive wastes associated with the operation of 

nuclear reactors and fuel reprocessing plants. 

How Is It Used? Iodine is used to treat cuts and scrapes on the skin as a tincture of iodine, which is a 

dilute mixture of alcohol and iodine. Iodine is also used in photography and lasers (silver iodide), in 

dyes, and as a nutrient added to table salt. Iodine-131 is used for a number of medical procedures, 

including to monitor and trace the flow of thyroxin from the thyroid. With its short half-life of 8 days, it 

is essentially gone in less than three months. Iodine-129 has no important commercial uses. 

What’s in the Environment? Iodine is present in nature in various materials, with soil, rock, and all 

living organisms containing low concentrations. Iodine is assimilated by seaweeds and sponges (from 

which it may be recovered) and is found in Chilean saltpeter, caliche, and brine associated with salt 

Iodine 
deposits. The ratio of stable iodine-127 to radioactive iodine-129 in the environment is more than 

10 million to 1. The human body contains 10 to 20 milligrams of iodine, of which more 

than 90% is contained in the thyroid gland. Iodine-129 is present in soil around the 

world as a result of fallout from past atmospheric nuclear weapons tests; any iodine-131 

that may have been present in soil from fallout has long since decayed away. Iodine may 

also be found as a contaminant at facilities where spent nuclear fuel was processed. 

Iodine-129 is one of the more mobile radionuclides in soil and can move downward with percolating 

water to groundwater. Iodine concentrations in sandy soil are about the same as in interstitial water (in 

the pore spaces between soil particles). It binds more preferentially to loam, where the concentration in 

soil is estimated to be 5 times higher than in interstitial water. 

What Happens to It in the Body? Iodine can be taken into the body by eating food, drinking water, or 

breathing air. It is a constituent of thyroid hormone and as such is a required element for humans. Iodine 

is readily taken into the bloodstream from both the lungs and the gastrointestinal tract (essentially 100%) 

after inhalation and ingestion. Upon entering the bloodstream, 30% is deposited in the thyroid, 20% is 

quickly excreted in feces, and the remainder is eliminated from the body within a short time (per 

simplified models that do not reflect intermediate redistribution). Clearance from the thyroid is agedependent, with biological half-lives ranging from 11 days in infants to 23 days in a five-year-old child 

and 80 days in adults. 

What Are the Primary Health Effects? Iodine is an essential component of the human diet, and lack 

of dietary iodine is a cause of goiter. Elemental iodine (I2) can be toxic, and its vapor irritates the eyes 

and lungs. While iodine is generally a health hazard only if it is taken into the body in substantial doses, 

iodine-131 emits fairly high-energy beta particles and a number of gamma rays. The gamma rays are of 

sufficient energy to be measured outside the body if deposited in tissue such as the thyroid. Because 

iodine selectively deposits in the thyroid, the primary health hazard for iodine is thyroid tumors resulting 

from ionizing radiation emitted by iodine-129 and iodine-131. Historically, the major exposure pathway 

has been ingestion of milk from cows grazing on iodine-contaminated crops. Other pathways include 

ingestion of fruits and vegetables and inhalation. 

What Is the Risk? Lifetime cancer mortality 

risk coefficients have been calculated for nearly 

all radionuclides, including iodine (see box at 

right). Additional values are also available, 

including for inhalation of iodine vapor and 

methyl iodide. Similar to other radionuclides, 

the risk coefficients for tap water are about 

75% of those shown for dietary ingestion. 

Thyroid cancer is the main risk associated with 

radioactive iodine. Based on epidemiological 

studies for external radiation, children are more 

susceptible than adults to cancer from thyroid 

irradiation. Data available for iodine-131 have 

not shown it to be carcinogenic in the human 

thyroid. While certain studies have indicated a 

potential effect from exposures to iodine-131 

(e.g., at Chernobyl, where external radiation 

was also quite high), others have not. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Recommended default 

absorption types were used for inhalation of 

particulates, and milk consumption values were used for 

ingestion. Risks are for lifetime cancer mortality per 

unit intake (pCi), averaged over all ages and both 

genders (10-9 is a billionth, and 10-12 is a trillionth). 

Other values, including for morbidity, are also 

available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Iodine-129 6.2 × 10-12 3.3 × 10-11

Iodine-131 2.1 × 10-12 1.4 × 10-11

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory/EVS Human Health Fact Sheet, March 2005 

Symbol: Ir

Atomic Number: 77 

(protons in nucleus)

Atomic Weight: 192 

(naturally occurring)

Radioactive Properties of Key Iridium Isotopes

Radiation Energy (MeV)

Isotope Half-Life 

Specific 

Activity 

(Ci/g)

Decay

Mode Alpha 

(α)

Beta

(β)

Gamma 

(γ)

Ir-192 74 days 9,200 β, EC - 0.22 0.82 

Ir-192m 240 yr 7.8 IT - - 0.16 

Ir-194m 170 days 4,000 β - 0.16 2.3 

EC = electron capture, IT = isomeric transition, Ci = curie, g = gram, and 

MeV = million electron volts; a dash means the entry is not applicable. (See 

the companion fact sheet on Radioactive Properties, Internal Distribution, 

and Risk Coefficients for an explanation of terms and interpretation of 

radiation energies.) Iridium-192 decays by two means: emitting a beta 

particle (95%) and electron capture (5%). Values are given to two significant 

figures. 

 

What Is It? Iridium is a silvery white metal named after the Latin word for rainbow 

because its salts are highly colored. Iridium is hard and brittle with low ductility, 

which makes it very difficult to machine and form. It is quite dense, about twice as 

dense as lead, and occurs in nature as two stable isotopes. (Isotopes are different forms 

of an element that have the same number of protons in the nucleus but a different 

number of neutrons.) Iridium-193 is the most prevalent form, comprising about 63% 

of natural iridium, with iridium-191 accounting for the rest. 

Of the 15 major radioactive iridium 

isotopes, only three have half-lives 

longer than a month and warrant 

concern; these are shown in the box 

at right. The half-lives of the other 

isotopes are less than 2 weeks. 

Iridium-192 has a half-life of 74 days, 

and it decays to stable platinum-192 

and osmium-192 by emitting a beta 

particle and by electron capture; most 

of these decays (95%) are by beta 

emission. Iridium-192m (the “m” 

means metastable) has a half-life of 

240 years, and it decays to 

iridium-192 by isomeric transition 

and by emitting a relatively lowenergy gamma ray. Iridium-194m has a half-life of 170 days, and it decays to stable platinum-194 by emitting a 

beta particle and several gamma rays. Of these three radioactive isotopes, iridium-192m presents the lowest acute 

or short-term hazard from external exposures as a result of its relatively low specific activity and gamma radiation; 

the other two iridium isotopes have high specific activities and significant gamma radiation. Iridium-192 is the 

isotope of most concern based on general availability; it is used in a number of industrial and medical applications. 

Where Does It Come From? Iridium is found uncombined in nature with platinum and other metals in the 

platinum group, and it can be obtained from platinum ores and as a byproduct of nickel mining. Naturally 

occurring iridium alloys include osmiridium and iridiosmium, which are mixtures of iridium and osmium. The 

highest iridium concentration on earth is at its core, and it can be released to the earth’s surface in volcanoes. 

Iridium-192 is produced by neutron activation of iridium metal, usually in nuclear reactors. The strength of an 

iridium-192 source is related to the amount of neutron irradiation. Natural iridium contains 37% iridium-191, and 

when this isotope absorbs a neutron it produces iridium-192. Neutron absorption by iridium-193, the other 

naturally occurring isotope, produces iridium-194. Iridium-194 has a short half-life of 19 hours, and it decays to 

stable platinum-194 in a few days; very little iridium-192m and iridium-194m are produced in this process. 

How Is It Used? Iridium is used as an alloying agent with a number of other metals to produce composites that 

are extremely hard and have good corrosion resistance. Its principal use is as a hardening agent in platinum alloys. 

Iridium is used for high-temperature applications, including in crucibles and thermocouples and as electrodes in 

spark plugs for severe operating conditions, such as those experienced by jet engine igniters. Radioactive 

iridum-192 is used industrially as a radiotracer in the oil industry and in gamma radiography to identify flaws in 

metal castings and welded joints. These radiographic sources are constructed of metal discs or pellets in a welded 

stainless steel capsule, and their activity levels range from less than one curie to several hundred curies. 

Iridium-192 is also used medically in brachytherapy to treat various types of cancer. (Brachytherapy is a method of 

radiation treatment in which sealed sources are used to deliver a radiation dose at a distance of up to a few 

centimeters by surface, intracavitary, or interstitial application.) Iridium-192 implants are used especially in the 

head and breast. They are produced in wire form and are introduced through a catheter to the target area. After 

Iridium 
being left in place for the time required to deliver the desired dose, the implant wire is removed. This procedure is 

very effective at providing localized radiation to the tumor site while minimizing the patient’s whole-body dose. 

What’s in the Environment? Iridium is naturally present in the earth’s crust at a concentration of about 

1 microgram per kilogram (μg/kg), or part per billion. In contrast, the concentration of iridium in some meteorites 

is about 500 μg/kg. A thin layer of iridium exists worldwide in a layer of deep sediment that was 

put down at the end of the Cretaceous period (roughly 150 million years ago), which is seen as 

evidence that the earth was struck by a large meteor or asteroid at that time. Dust from that impact 

would have spread around the globe, depositing iridium. As a very corrosion-resistant metal, 

iridium is quite insoluble in water. This means the iridium that is present in soil and rock tends to 

stay in place rather than being mobilized and transported in solution toward groundwater. There is 

no evidence that the three radioactive isotopes are present in soil around the world. These isotopes are not fission 

products so they would not exist in radioactive fallout from past atmospheric weapons tests. Only one isotope 

(iridium-192m) has a half-life longer than a year, and that isotope is not produced in any significant quantities. 

 

What Happens to It in the Body? Iridium can be taken into the body by eating food, drinking water, or breathing 

air. Gastrointestinal absorption from food or water is the likely source of internally deposited iridium in the general 

population. After ingestion or inhalation, most iridium is excreted from the body and never enters the bloodstream; 

only about 1% of the amount taken into the body by ingestion is absorbed into the blood. Twenty percent of the 

iridium that reaches the blood is excreted right away, 20% deposits in the liver, 4% deposits in the kidney, 2% 

deposits in the spleen, and the remaining 54% is evenly distributed among other organs and tissues of the body. Of 

the iridium that deposits in any organ or tissue, 20% leaves the body with a biological half-life of 8 days and 80% 

clears with a biological half-life of 200 days. On the basis of animal studies, retention of iridium was determined to 

be the same for all age groups. Most inhaled iridium compounds appear to clear the lungs quite rapidly. 

 

What Are the Primary Health Effects? The three radioactive iridium isotopes pose both an internal and 

external hazard, and the main health concern is associated with the increased likelihood of cancer. External 

exposure is a concern because of the strong gamma radiation (especially for iridium-192 and iridium-194m), and 

shielding is needed to handle iridium-192 radiographic and medical sources. Iridium can concentrate in several 

organs depending on its chemical form, so while there is no dominant organ of health concern the liver is a main 

organ of deposition. Inside the body, these iridium isotopes can pose a hazard from both beta and gamma radiation. 

 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including iridium (see box at right). While 

ingestion is generally the most common means for entry into 

the body, the risk coefficients for that route are lower than 

for inhalation. Similar to other radionuclides, the risk 

coefficients for tap water are about 70% of those for dietary 

ingestion. 

In addition to the risks from internal exposures, a risk from 

external gamma radiation is associated with these iridium 

isotopes. To estimate the lifetime cancer mortality risk, if it 

is assumed that 100,000 people were continuously exposed 

to a thick layer of soil with an initial average concentration 

of 1 pCi/g iridium-192m, then 10 of these people would be 

predicted to incur a fatal cancer. (This is in comparison to 

the 20,000 people from this group predicted to die of cancer 

from all other causes, per the U.S. average.) While the 

gamma radiation for iridium-192 and iridium-194m is much 

greater than for iridium-192m, these radionuclides will 

decay to very low levels in a few years. Because of this 

decay, less than one additional cancer fatality is predicted to occur among 100,000 people exposed to soil with an 

initial concentration of 1 pCi/g of either isotope. However, the respective external gamma radiation risks for acute 

exposures to iridium-192 and iridium-194m are 6 and 19 times higher than for iridium-192m. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Maximum values are given 

for inhalation (no default absorption types were 

provided), and dietary values were used for ingestion. 

Risks are for lifetime cancer mortality per unit intake 

(pCi), averaged over all ages and both genders (10-9 is 

a billionth and 10-12 is a trillionth). Other values, 

including for morbidity, are also available. 

Lifetime Cancer Mortality 

Risk Isotope 

Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Iridium-192 2.1 × 10-11 6.0 × 10-12

Iridium-192m 1.7 × 10-11 8.7 × 10-13

Iridium-194m 4.0 × 10-11 7.3 × 10-12

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Kr

Atomic Number: 36

(protons in nucleus)

Atomic Weight: 84

Radioactive Properties of Key Krypton Isotopes

Radiation Energy (MeV)

Isotope 

HalfLife 

(yr)

Specific

Activity

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Kr-81 210,000 0.021 EC - 0.0051 0.012 

Kr-85 11 400 β - 0.25 0.0022 

EC = electron capture, Ci = curie, g = gram, and MeV = million electron 

volts; a dash means the entry is not applicable. (See the companion fact 

sheet on Radioactive Properties, Internal Distribution, and Risk 

Coefficients for an explanation of terms and interpretation of radiation 

energies ) Values are given to two significant figures

What Is It? Krypton is a colorless, odorless, tasteless gas about three 

times heavier than air. It was discovered in 1898 by Sir William Ramsay 

and Morris Travers in the residue left after evaporating water, oxygen, 

nitrogen, helium, and argon from a sample of liquid air. The name comes 

from the Greek work kryptos, meaning hidden. As a noble gas, krypton is 

generally inert and forms very few chemical compounds. It occurs in 

nature as six stable isotopes. (Isotopes are different forms of an element that have the same number of 

protons in the nucleus but a different number of neutrons.) Krypton-84 is the most prevalent, comprising 

about 57% of natural krypton. The other five stable isotopes and their relative abundances are krypton-78 

(0.4%), krypton-80 (2.3%), krypton-82 (12%), krypton-83 (11%), and krypton-86 (17%). 

Eleven major radioactive isotopes of krypton exist of which only two – krypton-81 and krypton-85 – have 

half-lives long enough to warrant concern. Krypton-81 has a half-life of 210,000 years, and krypton-85 

has a half-life of 11 years; the 

half-lives of the other krypton 

isotopes are less than two days. 

Krypton-85 is the isotope of 

concern at Department of 

Energy (DOE) environmental 

management sites such as 

Hanford. It is produced by the 

fissioning of uranium and 

plutonium and is present in 

spent nuclear fuel. The low 

specific activity of krypton-81 

limits its radioactive hazards. 

Where Does It Come From? Krypton is naturally present in meteorites and minerals in trace 

quantities. It exists naturally in the atmosphere at a concentration of about 1 cubic centimeter per cubic 

meter (cm3

/m3). Radioactive krypton-85 is present in the natural environment in minute quantities due to 

the spontaneous and neutron-induced fission of uranium and other actinides. Krypton-81 and krypton-85 

are both present in the atmosphere due to neutron capture reactions from cosmic ray neutrons interacting 

with stable krypton isotopes. Krypton can be obtained as a byproduct from the liquefaction and 

separation of air. 

 

The major source of krypton-85 is nuclear fission. When an atom of uranium-235 (or other fissile nuclide) 

fissions, it generally splits asymmetrically into two large fragments – fission products with mass numbers 

in the range of about 90 and 140 – and two or three neutrons. (The mass number is the sum of the 

number of protons and neutrons in the nucleus of the atom.) Krypton-85 is one such fission product with 

a fission yield of about 0.3%. That is, three atoms of krypton-85 are produced per 1,000 fissions. An 

estimated 5 million curies of krypton-85 were released to the atmosphere as a result of nuclear weapons 

tests from 1945 through 1962. A large commercial nuclear power plant produces about 300,000 curies of 

krypton-85 per year, essentially all of which is retained within the fuel elements. This gaseous 

radionuclide is a component of spent nuclear fuel and is generally released to the atmosphere when the 

fuel is reprocessed. About 50,000 curies of krypton-85 were released to the atmosphere as a result of the 

accident at Three Mile Island in which a large number of fuel elements ruptured. 

How Is It Used? Krypton has a number of industrial and medical applications. It is used alone or in 

combination with argon and neon in fluorescent lights. It emits a characteristic bright orange-red color 

and is used in lights at airports because the red light is visible for long distances and penetrates fog and 

haze to a greater extent than ordinary light. Krypton is also used in tungsten-filament projection lamps 

for home movies and slide projectors. A krypton gas laser produces a very intense and concentrated light, 

Krypton 
and these lasers are used for medical applications such as surgery on the retina of the eye. The intense 

krypton laser light causes the blood to clot during the surgery, thus preventing further bleeding with 

subsequent loss of vision, and the laser is so accurate that surrounding tissues are not damaged. 

Krypton is also used as a standard because the spectral lines of its isotopes are very sharp. In 1960, the 

International Commission on Weights and Measures defined the length of the standard meter as exactly 

1,650,763.73 wavelengths (in a vacuum) of the orange-red line in the emission spectrum of krypton-86. 

This unit was redefined in October 1983 as the path length of light in a vacuum during a time interval of 

1/299,792,458 of a second. Radioactive krypton-85 is used to detect leaks from sealed containers, with 

the escaping atoms being identified through their radiation. Krypton-85 is also used to excite phosphors 

in light sources with no external source of energy and in medicine to detect abnormal heart openings. 

What’s in the Environment? The highest concentrations of krypton are in the atmosphere. Krypton is 

present in air at a concentration of about 1 cm3

/m3, or parts per million by volume. On 

a mass basis, the concentration is about 3 mg/kg. For comparison, the krypton 

concentration in the atmosphere of Mars is about 1/3 this amount (0.3 cm3

/m3). 

Krypton is naturally present in the earth’s crust at a concentration of about 

0.15 micrograms per kilogram (μg/kg), and its concentration in seawater is about 

0.21 μg/liter. Krypton-85 has been released to the atmosphere during nuclear fuel 

reprocessing activities and as a result of past aboveground nuclear weapons tests. In 1970, the 

concentration of krypton-85 in the atmosphere reached about 10 picocuries (pCi)/m3

 (or 10 trillionths of a 

curie per m3

), mainly from nuclear weapons tests and plutonium production activities. The concentration 

is significantly lower now due to the relatively short half-half of this radionuclide, the cessation of 

aboveground nuclear weapons tests worldwide by 1980, and the shutdown of plutonium production 

facilities at DOE sites. Neither the oceans nor the land surfaces act as significant sinks for this 

radionuclide. Krypton-85 is in spent nuclear fuel stored at certain sites (such as the DOE Hanford Site). 

What Happens to It in the Body? As a noble gas, krypton does not generally participate in any 

biological processes. After being taken into the body, a very small amount can be dissolved in the 

bloodstream and distributed to organs and tissues throughout the body. Nevertheless, the tissue of most 

concern from exposure to a cloud of krypton-85 gas is generally the skin, with most of the dose resulting 

from the beta particles associated with its radioactive decay. 

 

What Are the Primary Health Effects? The main health concern is the increased likelihood for 

cancer induction, and the exposure pathway of most concern is external exposure in a cloud of gas. The 

radiation dose for krypton-85 (the primary isotope of concern) from an external cloud of gas is more than 

130 times higher than the dose from any gas in the lungs and more than 200 times higher than that from 

any gas in body organs and tissues after being taken into the body. For kypton-81, most of the dose is 

associated with gamma rays that will irradiate all tissues and organs of the body. In contrast, much of the 

dose for kypton-85 is from beta particles, and the skin is the primary tissue of concern. 

What Is the Risk? Radiation doses from inhaling or ingesting krypton are small compared to the dose 

from external radiation, such as could occur in a cloud of krypton gas. In contrast to most other 

radionuclides, lifetime cancer mortality risk coefficients have not been developed for the inhalation and 

ingestion of krypton isotopes. The only pathway for which cancer mortality risk coefficients have been 

developed is external exposure. External gamma risk coefficients for krypton-81 and krypton-85 were 

used to estimate lifetime cancer mortality risks for submersion in krypton clouds. If it is assumed that 

krypton releases occurred and 100,000 people were continuously exposed to a cloud of air with an 

average concentration of 1 pCi/cm3

 over a period of one year, then the estimated number of fatal cancers 

in this group of 100,000 would be 2 for krypton-81 and less than 1 for krypton-85. (This is in comparison 

to the 20,000 people from this group who would be predicted to die of cancer from all other causes per 

the U.S. average.) This risk is due to the beta particles and gamma rays associated with the two krypton 

isotopes. (For more information, see the companion fact sheet on Radioactive Properties, Internal 

Distribution, and Risk Coefficients and the accompanying Table 1.) 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Np

Atomic Number: 93 

(protons in nucleus)

Atomic Weight: - 

(not naturally occurring)

Radioactive Properties of Key Neptunium Isotopes 

and Associated Radionuclides 

Radiation Energy (MeV) Isotope Half-Life 

(yr)

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma

(γ) 

Np-235 1.1 1,400 EC < 0.010 0.0071 

Np-236 120,000 0.013 β, EC - 0.21 0.14 

Pu-236 

 (9%) 

2.9 540 α 5.8 0.013 0.0021 

Np-237 2.1 million 0.00071 α 4.8 0.070 0.035 

Pa-233 0.074 21,000 β - 0.20 0.20 

EC = electron capture, Ci = curie, g = gram, and MeV = million electron 

volts; a ”<” means the radiation energy is less than 0.001 MeV, and a dash 

means the entry is not applicable. (See the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk Coefficients for 

explanation of terms and interpretation of radiation energies.) About 

0.001% of the neptunium-235 decays are by alpha-particle emission. The 

isotope neptunium-236 decays both by emitting a beta particle (9%) and by 

electron capture (91%); another isotope of neptunium-236 with a half-life of 

23 hours also exists. Certain properties of plutonium-236 and protactinium233 are included here because these radionuclides accompany the neptunium 

decays. Values are given to two significant figures. 

What Is It? Neptunium is a ductile, silver-colored metal about twice as 

dense as lead. It does not occur naturally but is produced artificially by 

neutron capture reactions by uranium. There are seventeen known isotopes 

of neptunium, and all are radioactive. (Isotopes are different forms of an 

element that have the same number of protons in the nucleus but a different 

number of neutrons.) The first neptunium isotope to be identified was 

neptunium-239, which has a half-life of 2.4 days. This isotope was first 

produced in 1940 in a cyclotron at the University of California at Berkeley by 

bombarding uranium-238 with high-energy neutrons. Neptunium was the first transuranic element to be 

formed and was named for the planet Neptune. 

Of the seventeen neptunium isotopes, only three have half-lives long enough to warrant concern at 

Department of Energy (DOE) environmental management sites: neptunium-235, neptunium-236, and 

neptunium-237. The half-lives 

of these three isotopes range 

from 1.1 to 2.1 million years, 

while those of the other isotopes 

are less than five days. Of the 

three, neptunium-237 is the most 

prevalent isotope at DOE sites 

such as Hanford. It has a halflife of 2.1 million years and 

decays by emitting an alpha 

particle. The other two isotopes 

typically represent less than a 

few percent of the total 

neptunium inventory at a site. 

Neptunium-235 has a half-life of 

1.1 years and decays by electron 

capture; essentially all of this 

isotope that was produced more 

than 20 years ago has long since 

decayed away. Neptunium-236 

has a half-life of 120,000 years 

and decays by emitting a beta 

particle and electron capture. 

Where Does It Come From? Neptunium is a byproduct of plutonium production activities and results 

from the capture of neutrons by uranium isotopes, usually in a nuclear reactor. Neptunium isotopes can be 

formed by a variety of neutron capture and radioactive decay routes. Neptunium is present in spent nuclear 

fuel, high-level radioactive wastes resulting from the processing of spent nuclear fuel, and radioactive wastes 

associated with the operation of reactors and fuel reprocessing plants. Although neptunium is essentially not 

naturally present in the environment, very minute amounts may be associated with uranium ores. 

How Is It Used? There are no major commercial uses of neptunium, although neptunium-237 is used as a 

component in neutron detection instruments. Neptunium-237 can also be used to make plutonium-238 (by 

absorption of a neutron). Neptunium is considered useable in nuclear weapons, although no country is 

known to have used it to make a nuclear explosive device. 

What’s in the Environment? Atmospheric testing of nuclear weapons, which ceased worldwide by 1980, 

generated most environmental neptunium. The level of neptunium in soil from fallout is quite low; for 

Neptunium 
 

example, the concentration of neptunium-237 is less than 1% of that for plutonium-239 (on the order of 

0.0001 picocuries per gram, pCi/g). Accidents and other releases from weapons production facilities have 

caused localized contamination. Neptunium typically occurs in the environment as an oxide, although other 

forms can be present. It is generally more mobile than other transuranic elements such as plutonium, 

americium, and curium, and it can move down with percolating water to underlying layers of 

soil. Neptunium preferentially adheres to soil particles, with the concentration associated with 

sandy soil particles estimated to be about 5 times higher than in interstitial water (water in pore 

spaces between the soil particles); it bonds more tightly to clay soils, where concentration ratios 

are typically higher (55). Neptunium is readily taken up by plants, and plant concentrations are 

typically similar to soil concentrations. 

What Happens to It in the Body? Neptunium can be taken into the body by eating food, drinking water, 

or breathing air. Gastrointestinal absorption from food or water is a likely source of internally deposited 

neptunium in the general population. After ingestion or inhalation, most neptunium is excreted from the 

body within a few days and never enters the bloodstream; only about 0.05% of the amount taken into the 

body by ingestion is absorbed into the blood. After leaving the intestine or lung, about 50% of the 

neptunium that does enter the bloodstream deposits in the skeleton, about 10% deposits in the liver, about 5% 

deposits in other soft tissues, and the rest is excreted, primarily in urine. The biological half-lives in the 

skeleton and liver are about 50 and 20 years, respectively. (This information is per simplified models that do 

not reflect intermediate redistribution.) The amount deposited in the liver and skeleton depends on the age of 

the individual, with fractional uptake in the liver increasing with age. Neptunium in the skeleton is deposited 

on bone surfaces and slowly redistributes throughout the bone volume over time. 

What Is the Primary Health Effect? Neptunium is generally a health hazard only if it is taken into the 

body, although there is an external risk associated with the gamma rays emitted by neptunium-236 and 

neptunium-237 and its short-lived decay product protactinium-233. The main means of exposure are 

ingestion of food and water containing neptunium isotopes and inhalation of neptunium-contaminated dust. 

Ingestion is generally the exposure of concern unless there is a nearby source of contaminated airborne dust. 

Because neptunium is taken up in the body much more readily if inhaled rather than ingested, both exposure 

routes can be important. The major health concern is cancer resulting from the ionizing radiation emitted by 

neptunium isotopes deposited on bone surfaces and in the liver. 

What Is the Risk? Lifetime cancer mortality 

risk coefficients have been calculated for nearly 

all radionuclides, including neptunium (see box 

at right). While ingestion is generally the most 

common route of exposure, the risk coefficients 

for this route are much lower than those for 

inhalation. Similar to other radionuclides, the 

risk coefficients for tap water are about 70 to 

75% of those for dietary ingestion. 

In addition to risks from internal exposures, there 

is an external gamma risk associated with 

exposure to neptunium-236 and neptunium-237. 

To estimate a lifetime cancer mortality risk, if it 

is assumed that 100,000 people were 

continuously exposed to a thick layer of soil with 

an initial average concentration of 1 pCi/g, then 2 

of these 100,000 people would be predicted to 

incur a fatal cancer if the soil contained 

neptunium-236, and 4 if it contained neptunium-237. (This is in comparison to the 20,000 people from the 

group predicted to die of cancer from all other causes per the U.S. average.) The external risk for 

neptunium-237 is largely due to its short-lived decay product protactinium-233. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for inhalation 

and ingestion. Recommended default absorption types were 

used for inhalation, and dietary values were used for 

ingestion. Values include contributions from the short-lived 

neptunium decay products. Risks are for lifetime cancer 

mortality per unit intake (pCi), averaged over all ages and 

both genders (10-9 is a billionth, and 10-12 is a trillionth). 

Other values, including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Neptunium-235 1.0 × 10-12 2.8× 10-13

Neptunium-236 2.6 × 10-9 1.5 × 10-11

Neptunium-237 1.5 × 10-8 5.8 × 10-11

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Radioactive Properties of Key Nickel Isotopes

Radiation Energy (MeV)

Isotope 

HalfLife 

(yr)

Specific

Activity

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Ni-59 75,000 0.082 EC - 0.0046 0.0024 

Ni-63 96 60 β - 0.017 - 

EC =electron capture, Ci = curie, g = gram, and MeV = million electron 

volts; a dash means the entry is not applicable. (See the companion fact 

sheet on Radioactive Properties, Internal Distribution, and Risk 

Coefficients for an explanation of terms and interpretation of radiation 

energies.) Values are given to two significant figures. 

Symbol: Ni

Atomic Number: 28 

(protons in nucleus)

Atomic Weight: 59 

(naturally occurring)

What Is It? Nickel is a hard, silvery-white metal that is malleable and ductile. It 

occurs in nature as five stable isotopes. (Isotopes are different forms of an element that 

have the same number of protons in the nucleus but a different number of neutrons.) 

Nickel-58 is the most prevalent form, comprising about two-thirds of natural nickel. 

The other four stable isotopes and their relative abundance are nickel-60 (26%), 

nickel-61 (1.1%), nickel-62 (3.6%), and nickel-64 (0.9%). 

Of the six major radioactive isotopes, only two – nickel-59 and nickel-63 – have half-lives long enough to warrant 

concern. The half-lives of all other nickel isotopes are less than six days. Nickel-59 decays with a half-life of 

75,000 years by electron capture, and 

nickel-63 decays with a half-life of 

96 years by emitting a beta particle. 

Both isotopes are present in wastes 

resulting from the reprocessing of 

spent nuclear fuel. Nickel-63 is 

generally the isotope of most concern 

at U.S. Department of Energy (DOE) 

environmental management sites such 

as Hanford. The long half-life of 

nickel-59 (with its subsequent low 

specific activity) combined with its 

low decay energy limits the 

radioactive hazards associated with 

this isotope. 

Where Does It Come From? Nickel is naturally present in various ores and to a lesser extent in soil. It occurs in 

minerals such as garnierite, millerite, niccolite, pentlandite, and pyrrhotite, with the latter two being the principal ores. 

It is also found in most meteorites and often serves as one of the criteria for distinguishing a meteorite from other 

minerals. Most of the world’s supply of nickel is mined in Canada; other sources include Cuba, the former Soviet 

Union, China, and Australia. The United States has no large deposits of nickel and accounts for less than 1% of the 

annual world output. Most nickel used in the United States is imported, and about 30% of the annual consumption is 

from recycled sources. 

The two radioactive isotopes of potential concern are produced by neutron activation of components in nuclear reactors. 

When a fissile nuclide such as an atom of uranium-235 fissions, it generally splits asymmetrically into two large 

fragments – fission products with mass numbers in the range of about 90 and 140 – and two or three neutrons. (The 

mass number is the sum of the number of protons and neutrons in the nucleus of the atom.) These neutrons can cause 

additional fissions (producing a chain reaction), escape from the reactor, or irradiate nearby materials. A number of 

reactor components are made of various alloys of steel that contain chromium, manganese, nickel, iron, and cobalt, and 

these elements can absorb neutrons resulting in radioactive isotopes, including nickel-59 and nickel-63. Nickel-59 and 

nickel-63 are radionuclides of concern in spent nuclear fuel (as a component of the fuel hardware) and the radioactive 

wastes associated with operating nuclear reactors and fuel reprocessing plants. Nickel-63 is present in much higher 

concentrations than nickel-59. 

How Is It Used? Nickel is used in various coins and as a component of several alloys, including nichrome and 

permalloy, and in some stainless steels. Alnico, an alloy of aluminum, nickel, cobalt, and other metals, is used to make 

high-strength, permanent magnets. Nickel alloy steels are used in heavy machinery, manufacturing, armaments, tools, 

and high-temperature equipment, including gas turbines and environmental devices used to control emissions such as 

scrubbers. Nickel is also used as a protective and ornamental coating for metals susceptible to corrosion, particularly 

iron and steel. The nickel plate is deposited by electrolysis in a nickel solution, and the coating can be highly polished. 

What’s in the Environment? Nickel is present in crustal rock at a concentration of about 

90 milligrams per kilogram (mg/kg). Its concentration in seawater is about 2 mg per liter (mg/L). Trace 

amounts of nickel-59 and nickel-63 are present around the globe from radioactive fallout. It can also be 

present at certain nuclear facilities as a contaminant from operating reactors and processing spent fuel. 

Nickel is generally one of the less mobile radioactive metals in the environment. The typical ratio of the 

concentration of nickel in plants to that in soil is low, estimated at 0.06 (or 6%). It also adheres quite well to soil. The 

Nickel 
concentration of nickel associated with sandy soil particles is typically about 400 times higher than in interstitial water 

(in the pore spaces between the soil particles); it binds even more tightly to clay soil where concentration ratios can 

exceed 600. Thus, nickel is generally not a major contaminant in groundwater at DOE sites. 

What Happens to It in the Body? Nickel can be taken into the body by eating food, drinking water, or breathing air. 

Children, and to a lesser extent adults, can also be exposed by ingesting soils. Gastrointestinal absorption from food or 

water is the principal source of internally deposited nickel in the general population. About 5% of the amount ingested 

is absorbed into the bloodstream through the intestines, while 20 to 35% of inhaled nicked is absorbed through the 

lungs. Of the nickel that reaches the blood, 68% is rapidly excreted in urine, while 2% remains in the kidneys with a 

very short biological half-life of 0.2 days (about 5 hours). The remaining 30% is evenly distributed to all remaining 

tissues of the body, including the kidneys, and clears with a biological half-life of more than 3 years (1,200 days). (This 

information is based on simplified models that do not reflect intermediate redistribution.) Nickel can be absorbed into 

the skin where it may stay, instead of being absorbed into the blood. 

 

What Are the Primary Health Effects? Nickel is a radiogenic health hazard only if it is taken into the body. 

External gamma exposure is not a concern because nickel-63 and nickel-59 do not emit significant gamma radiation. 

Nickel-63 decays by emitting a beta particle and nickel-59 decays by electron capture, in which low-energy gamma 

radiation is emitted. While in the body, radioactive nickel presents a health hazard from the beta particles and gamma 

radiation; the main health concern is associated with the increased likelihood of inducing cancer. Nickel also exhibits 

chemical toxicity. The most common effect is an allergic reaction of the skin, with about 10 to 15% of the population 

sensitive to nickel (e.g., in jewelry). Less frequently, nickel induces asthmatic attacks. Both human and animal studies 

indicate that the respiratory system is the primary target following acute inhalation of high concentrations of nickel. 

Acute toxicity following ingestion of high concentrations includes effects on the gastrointestinal system, blood, and 

kidneys. Effects reported in studies of workers chronically exposed to airborne nickel dusts include chronic bronchitis, 

reduced lung function, and cancer of the lung and nasal sinus. The U.S. Environmental Protection Agency (EPA) has 

classified nickel subsulfide (a relatively insoluble form of nickel) as a known human carcinogen. 

What Is the Risk? Lifetime cancer mortality risk coefficients have been calculated for nearly all radionuclides, 

including nickel (see box at right). While the coefficients for ingestion are somewhat lower than for inhalation, 

ingestion is generally the most common means of 

entry into the body. Similar to other radionuclides, 

the risk coefficients for tap water are about 70% of 

those for dietary ingestion. The EPA has developed 

toxicity values to estimate the risk of getting cancer 

or other adverse health effects associated with the 

chemical toxicity of nickel (see box at right). The 

toxicity value for estimating cancer risk following 

inhalation exposure is called a unit risk (UR), which 

is an estimate of the chance that a person will get 

cancer from continuous exposure to a chemical in 

air at a concentration of 1 milligram per cubic meter 

(mg/m3

). For example, using the inhalation UR, the 

EPA estimates that a person would have a one-in-amillion chance of developing cancer if exposed 

daily over a lifetime to air containing 0.002 

microgram per cubic meter (μg/m3

) nickel 

subsulfide. The toxicity value for estimating noncancer effects is a reference dose (RfD), which is an 

estimate of the highest dose that can be taken in 

every day without causing an adverse effect. The 

RfD for nickel was developed by studying test 

animals given relatively high doses over their 

lifetimes, then adjusting and normalizing those 

results to a mg per kg-day basis for humans. EPA 

has not derived a toxicity value for evaluating noncancer effects of nickel following inhalation 

exposure. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for inhalation and 

ingestion. Recommended default absorption types were used 

for inhalation, and dietary values were used for ingestion. 

Risks are for lifetime cancer mortality per unit intake (pCi), 

averaged over all ages and both genders (10-12 is a trillionth). 

Other values, including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Nickel-59 3.6 × 10-13 2.3 × 10-13

Nickel-63 1.4 × 10-12 5.7 × 10-13

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 

Chemical Toxicity Values 

Cancer Risk Non-Cancer Effect Form of Nickel 

Inhalation UR Oral RfD

Soluble salts None established 0.02 mg/kg-day

Nickel carbonyl None established None established 

Nickel refinery 

dust 

2.4 × 10-1 

per mg/m3 None established 

Nickel subsulfide 4.8 × 10-1 

per mg/m3 None established 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Pu 

Atomic Number: 94 

(protons in nucleus)

Atomic Weight: - 

(not naturally occurring)

Radioactive Properties of Key Plutonium Isotopes

Radiation Energy (MeV)

Isotope Half-Life Specific Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma

(γ) 

Pu-236 2.9 yr 540 α 5.8 0.013 0.0021 

Pu-238 88 yr 17 α 5.5 0.011 0.0018 

Pu-239 24,000 yr 0.063 α 5.1 0.0067 < 

Pu-240 6,500 yr 0.23 α 5.2 0.011 0.0017 

Pu-241 14 yr 100 β < 0.0052 < 

Pu-242 380,000 yr 0.0040 α 4.9 0.0087 0.0014 

Pu-244 83,000,000 yr 0.000018 α 4.6 0.0071 0.0012 

U-240 14 hr 940,000 β - 0.14 0.0076 

Np-240m 7.4 min 110 million β - 0.68 0.34 

Ci = curie, g = gram, and MeV = million electron volts; a “<” means the radiation 

energy is less than 0.001 MeV, and a dash means the entry is not applicable. (See the 

companion fact sheet on Radioactive Properties, Internal Distribution, and Risk 

Coefficients for an explanation of terms and interpretation of radiation energies.) A 

very small fraction (about 0.002%) of the plutonium-241 decays are by alpha-particle 

emission. Certain properties of uranium-240 and neptunium-240m are included 

here because these radionuclides accompany the plutonium-244 decays. Values are 

given to two significant figures. 

What Is It? Plutonium in its pure form is a very heavy, silver-colored, radioactive metal 

about twice as dense as lead. Essentially all the plutonium on earth has been created within 

the past six decades by human activities involving fissionable materials. Several plutonium 

isotopes exist, all of which are radioactive. (Isotopes are different forms of an element that 

have the same number of protons in the nucleus but a different number of neutrons.) 

The main plutonium isotopes at 

Department of Energy (DOE) 

environmental management sites 

are plutonium-238, plutonium-239, 

plutonium-240, and plutonium-241. 

Except for plutonium-241, these 

isotopes decay by emitting an alpha 

particle. Plutonium-241 decays by 

emitting a low-energy beta particle 

to americium-241, an alphaemitting radionuclide with a halflife of 430 years that is much more 

radiotoxic than its parent. The 

maximum activity of 

americium-241 is about 3% of the 

initial activity of plutonium-241 

and occurs 73 years later. An 

extremely small fraction of the 

decays of plutonium-236, 

plutonium-238, plutonium-240, and 

plutonium-242, are by spontaneous 

fission (SF), as are about 0.1% of 

the plutonium-244 decays. 

Plutonium-242 and plutonium-244 

are generally present in relatively 

minute activity concentrations. 

Where Does It Come From? Plutonium was first made in large quantities by American scientists in the 1940s as part of the 

Manhattan Project to create the atomic bomb, and this production continued through the Cold War. Plutonium is formed when 

the nucleus of a uranium atom captures one or more neutrons, changing the atomic structure and creating a new element. This 

process occurs in nuclear reactors and mainly involves transforming uranium-238 into plutonium. (Extremely small quantities 

of plutonium were created naturally in sustained underground nuclear reactions estimated to have occurred about 1.9 billion 

years ago in Gabon, Africa. This phenomenon occurred because concentrations of uranium-235 were much higher at that time. 

The current uranium-235 concentration, about 0.72%, will not sustain such natural reactions.) 

How Is It Used? The nuclear properties of plutonium-239, as well as our ability to produce large amounts of nearly pure 

plutonium-239, led to its use in nuclear weapons and nuclear power. The fissioning of uranium-235 in the reactor of a nuclear 

power plant produces two to three neutrons, and these neutrons can be absorbed by uranium-238 to produce plutonium-239 and 

other isotopes. Plutonium-239 can also absorb neutrons and fission along with the uranium-235. Plutonium fissions provide 

about one-third of the total energy produced in a typical commercial nuclear power plant. The use of plutonium in power 

plants occurs without it ever being removed from the nuclear reactor fuel, i.e., it is fissioned in the same fuel rods in which it is 

produced. Another isotope, plutonium-238, is used as a heat source in radiothermal generators to produce electricity for a 

variety of purposes including unmanned spacecraft and interplanetary probes. The United States recovered or acquired about 

110,000 kilograms (kg) of plutonium between 1944 and 1994, and about 100,000 kg remains in inventory. Of this amount, 

over 80% is in the form of weapons-grade plutonium, primarily plutonium-239. Plutonium was generated in production 

reactors at DOE’s Hanford and Savannah River sites, and weapons components were produced at the Rocky Flats facility. 

Surplus plutonium is currently stored at the Pantex Plant and other sites. 

What’s in the Environment? Atmospheric testing of nuclear weapons, which ceased worldwide by 1980, 

generated most environmental plutonium. About 10,000 kg were released to the atmosphere during these tests. 

Average plutonium levels in surface soil from fallout range from about 0.01 to 0.1 picocurie per gram (pCi/g). 

Plutonium 
 

Plutonium metal. Plutonium isotopes are 

primarily alpha-emitters so they pose little 

risk outside the body. Here the plastic bag, 

gloves, and outer (dead) layer of skin would 

each alone stop the emitted alpha particles 

from getting into the body. 

Accidents and other releases from weapons production facilities have caused greater localized contamination. The most 

common form in the environment is plutonium oxide. Plutonium is typically very insoluble, with the oxide being less soluble 

in water than ordinary sand (quartz). It adheres tightly to soil particles and tends to remain in the top few centimeters of soil as 

the oxide. In aquatic systems, plutonium tends to settle out and adhere strongly to sediments, again remaining in upper layers. 

Typically one part of plutonium will remain in solution for every 2,000 parts in sediment or soil. A small fraction of plutonium 

in soil can become soluble through chemical or biological processes, depending on its chemical form. While plutonium can 

bioconcentrate in aquatic organisms, data have not indicated that it biomagnifies in aquatic or terrestrial food chains. 

What Happens to It in the Body? When plutonium is inhaled, a 

significant fraction can move from the lungs through the blood to other 

organs, depending on the solubility of the compound. Little plutonium 

(about 0.05%) is absorbed from the gastrointestinal tract after ingestion, 

and little is absorbed through the skin following dermal contact. After 

leaving the intestine or lung, about 10% clears the body. The rest of what 

enters the bloodstream deposits about equally in the liver and skeleton 

where it remains for long periods of time, with biological retention halflives of about 20 and 50 years, respectively, per simplified models that do 

not reflect intermediate redistribution. The amount deposited in the liver 

and skeleton depends on the age of the individual, with fractional uptake 

in the liver increasing with age. Plutonium in the skeleton deposits on the 

cortical and trabecular surfaces of bones and slowly redistributes 

throughout the volume of mineral bone with time. 

What Is the Primary Health Effect? Plutonium generally poses a health hazard only if it is taken into the body because all 

isotopes except plutonium-241 decay by emitting an alpha particle, and the beta particle emitted by plutonium-241 is of low 

energy. Minimal gamma radiation is associated with these radioactive decays. However, there is an external gamma radiation 

hazard associated with plutonium-244 from it short-lived 

decay product neptunium-240m. Inhaling airborne 

plutonium is the primary concern for all isotopes, and 

cancer resulting from the ionizing radiation is the health 

effect of concern. The ingestion hazard associated with 

common forms of plutonium is much lower than the 

inhalation hazard because absorption into the body after 

ingestion is quite low. Laboratory studies with experimental 

animals have shown that exposure to high levels of 

plutonium can cause decreased life spans, diseases of the 

respiratory tract, and cancer. The target tissues in those 

animals were the lungs and associated lymph nodes, liver, 

and bones. However, these observations in experimental 

animals have not been corroborated by epidemiological 

investigations in humans exposed to lower levels. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including plutonium (see box at right). While 

ingestion is generally the most common route of exposure, 

the risk coefficients for this route are much lower than those 

for inhalation. As for other radionuclides, the risk 

coefficients for tap water are about 80% of those for dietary 

ingestion. In addition to risks from internal exposures, 

there is an external gamma exposure risk associated with 

plutonium-244. To estimate a lifetime cancer mortality risk, 

if it is assumed that 100,000 people were continuously 

exposed to a thick layer of soil with an initial concentration 

of 1 pCi/g of plutonium-244, then 7 of these 100,000 people 

would be predicted to incur a fatal cancer. (This is in comparison to the 20,000 people from this group predicted to die of 

cancer from all other causes per the U.S. average.) This risk is largely associated with the gamma rays emitted by its shortlived decay product neptunium-240m. The external risk for the other plutonium isotopes is less than 1% of that for 

plutonium-244. As a note, for inhalation (the exposure of highest risk), breathing in 5,000 respirable plutonium particles of

about 3 microns each is estimated to increase an individual’s risk of incurring a fatal cancer about 1% above the U.S. average.

Radiological Risk Coefficients 

This table provides selected risk coefficients for inhalation 

and ingestion. Recommended default absorption types were 

used for inhalation, and dietary values were used for 

ingestion. Risks are for lifetime cancer mortality per unit 

intake (pCi), averaged over all ages and both genders (10-9

is a billionth, and 10-12 is a trillionth). Other values, 

including for morbidity, are also available. Because values 

are not available for plutonium-244, these were estimated by 

multiplying the risk factors for plutonium-242 by the ratio of 

the dose conversion factors for plutonium-244 to 

plutonium-242. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Plutonium-236 2.1 × 10-8 6.9 × 10-11

Plutonium-238 3.0 × 10-8 1.3 × 10-10

Plutonium-239 2.9 × 10-8 1.3 × 10-10

Plutonium-240 2.9 × 10-8 1.3 × 10-10

Plutonium-241 2.8 × 10-10 1.9 × 10-12

Plutonium-242 2.8 × 10-8 1.3 × 10-10

Plutonium-244 2.7 × 10-8 1.3 × 10-10

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Symbol: Po

Atomic Number: 84 

(protons in nucleus)

Atomic Weight: 210 

(naturally occurring)

Radioactive Properties of Key Polonium Isotopes

Radiation Energy (MeV)

Isotope Half-Life 

Specific 

Activity 

(Ci/g)

Decay

Mode Alpha 

(α)

Beta

(β)

Gamma 

(γ)

Po-208 2.9 yr 590 α 5.1 < < 

Po-209 100 yr 17 α 4.9 < < 

Po-210 140 days 4,500 α 5.3 < < 

Ci = curie, g = gram, and MeV = million electron volts; a “<” means the 

radiation energy is less than 0.001 MeV. (See the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk Coefficients for an 

explanation of terms and interpretation of radiation energies.) Values are 

given to two significant figures. Polonium-210 is a decay product of 

radium-226 and is also shown on that fact sheet. The basic properties of 

polonium-208 and polonium-209 (which are not in the natural decay series) 

are also given here because they are included in the general discussion below.

 

What Is It? Polonium is a radioactive element that occurs naturally in very low 

concentrations in the earth’s crust (at about one part in 1015, or one millionth of a 

trillionth). Polonium was the first element discovered by Marie and Pierre Curie in 

1898, while seeking the cause of radioactivity of pitchblende ore containing uranium. 

Polonium in its pure form is a low-melting, fairly volatile metal. Over 25 isotopes of 

polonium are known, with atomic masses ranging from 192 to 218 (isotopes are 

different forms of an element that have the same number of protons in the nucleus but a 

different number of neutrons.) All polonium isotopes are radioactive, with only three 

having appreciable half-lives: polonium-208, polonium-209, and polonium-210. 

Polonium-210, historically called “radium F,” is the predominant naturally occurring isotope of polonium and the 

one most widely used. Polonium-210 

is a radioactive decay product in the 

natural uranium-238 decay series; 

along with lead-210 it is one of two 

relatively long-lived decay products 

of radon-222. Polonium-210 has a 

half-life of 138 days, and it decays to 

stable lead-206 by emitting an alpha 

particle. One-thousandth of a gram 

(1 mg) of polonium-210 emits as 

many alpha particles as 5 g of 

radium-226. The energy released by 

its decay is so large (140 watts/g) 

that a capsule containing about half a 

gram reaches a temperature above 

500oC. 

Where Does It Come From? Because it is produced during the decay of naturally ubiquitous uranium-238, 

polonium-210 is widely distributed in small amounts in the earth’s crust. Although it can be produced by the 

chemical processing of uranium ores or minerals, uranium ores contain less than 0.1 mg polonium-210 per ton. 

Originally, polonium-210 was obtained from the rich pitchblende ore found in Bohemia, but it can also be obtained 

from aged radium salts that contain about 0.2 mg per gram of radium. Although a number of other polonium 

isotopes are present in the natural decay series, their short half-lives preclude any appreciable concentrations. 

Due to its scarcity, polonium-210 is usually produced artificially in a nuclear reactor by bombarding bismuth-209 (a 

stable isotope) with neutrons. This forms radioactive bismuth-210, which has a half-life of 5 days. Bismuth-210 

decays to polonium-210 through beta decay. Milligram amounts of polonium-210 have been produced by this 

method. The longer-lived isotopes polonium-209 (half-life 103 years) and polonium-208 (half-life 2.9 years) are 

also produced in reactors or particle accelerators, but these are very expensive. 

How Is It Used? Polonium-210 is used mainly in static eliminators, which are devices designed to eliminate 

static electricity in machinery where it can be caused by processes such as paper rolling, manufacturing sheet 

plastics, and spinning synthetic fibers. The polonium-210 is generally electroplated onto a backing foil and inserted 

into a brush, tube, or other holder. Alpha particles from the polonium ionize adjacent air, and the air ions then 

neutralize static electricity on the surfaces in contact with the air. These devices generally need to be replaced every 

year because of the short half-life of this radioisotope. Polonium-210 is also used in brushes to remove dust from 

photographic films and camera lenses. Static eliminators typically contain from tens to hundreds of mCi 

(thousandth of a curie) of radioactivity. Polonium-210 can also be combined with beryllium to produce neutron 

sources, and in fact it was used as neutron-producing initiators of at least the first generation of atomic weapons. In 

addition, polonium-210 has been investigated as a heat source for thermoelectric power devices for space 

applications. 

Polonium 
What’s in the Environment? Polonium-210 is naturally present in all environmental media at very low 

concentrations. In soils, the concentration is similar to that of uranium, averaging about 1 pCi/g (or one trillionth 

curie per gram). Because polonium-210 is produced from the decay of radon-222 gas, it can be found in the 

atmosphere from which it is deposited on the earth’s surface. Average annual air concentrations 

range from 0.005 to 0.04 pCi/m3

. Polonium-210 is also emitted to the atmosphere during the 

calcining of phosphate rock as part of the production of elemental phosphorous. Although direct 

root uptake by plants is generally small, polonium-210 can be deposited on broad-leaved 

vegetables. Deposition from the atmosphere on tobacco leaves results in elevated concentrations of 

polonium-210 in tobacco smoke, resulting in greater intakes in smokers compared to non-smokers. 

It is estimated that the average Western diet includes from 1 to 10 pCi of polonium-210 per day. Polonium-210 can 

be significantly elevated in residents of northern lands who subsist on reindeer that consume lichens, which absorb 

trace elements from the atmosphere. 

What Happens to It in the Body? Polonium can be taken into the body by eating food, drinking water, or 

breathing air. Between 50% and 90% of the polonium taken in by ingestion will promptly leave the body in feces. 

The fraction remaining in the body enters the bloodstream. In general, the spleen and kidneys concentrate polonium 

more than other tissues except for temporary deposition in the lung after inhalation of an insoluble form. It is 

estimated that approximately 45% of ingested polonium will be deposited in the spleen, kidneys, and liver, with 

10% deposited in bone marrow and the remainder distributed throughout the body. The amount of polonium in the 

body will decrease with a half-time of 50 days. 

Studies of smokers have shown that inhaled polonium can be highly localized in the lungs, with about twice as 

much polonium found in the ribs of smokers compared to nonsmokers. It is estimated that the dose to the skeleton 

is elevated about 30% in smokers. Another source of polonium-210 in the body is its gradual ingrowth from the 

decay of radium-226 and lead-210 deposited in bone. The average amount of polonium-210 in the body is 

approximately 1 nCi (one billionth of a curie). 

What Are the Primary Health Effects? Polonium-210 is a health hazard only if it is taken into the body. 

External exposure is not a concern because polonium is an alpha emitter. The primary means of exposure are 

ingestion of food and water containing polonium-210 and inhalation of polonium-contaminated dust. Inhalation is 

of particular concern in the vicinity of a source of airborne dust, such as a phosphate plant, and in areas of high 

radon concentrations, or for cigarette smokers. 

Substantial radiation doses from polonium can be 

expected in many tissues of the body; it supplies a more 

nearly whole-body dose than almost all other alpha 

emitters. Effects are more common in the kidney than 

the spleen, despite a higher dose in the spleen. The 

lymph nodes and liver can also be affected. Polonium 

that is inhaled, either from radon in the air or cigarette 

smoke, can be deposited on the mucous lining of the 

respiratory tract. When alpha particles are then emitted 

within the lung, the cells lining the airways can be 

damaged, potentially leading to lung cancer over time. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including polonium-210 (see box at 

right). Risk coefficients for inhalation are about 6 times 

higher than for dietary ingestion. Similar to other 

radionuclides, the risk coefficients for ingestion of tap 

water containing polonium-210 are about 75% of those 

shown for dietary ingestion. Polonium-210 poses no 

external risk when outside the body.

Radiological Risk Coefficients 

This table provides risk coefficients for inhalation and 

absorption. Recommended default absorption types 

were used for inhalation, and dietary values were used 

for ingestion. Risks are for lifetime cancer mortality per 

unit intake (pCi), averaged over all ages and both 

genders (10-9 is a billionth). Other values, including for 

morbidity, are also available. No risk or dose conversion 

factors are available for Po-208 or Po-209. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Po-210 1.0 × 10-8 1.6 × 10-9

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: K(-40) 

Atomic Number: 19 

(protons in nucleus)

Atomic Weight: 39 

(naturally occurring)

Radioactive Properties of Potassium-40

Radiation Energy (MeV)

Isotope Half-Life 

(yr)

Natural 

Abundance 

(%)

Specific 

Activity 

(Ci/g)

Decay

Mode Alpha 

(α)

Beta

(β)

Gamma 

(γ)

K-40 1.3 billion 0.012 0.0000071 β, EC - 0.52 0.16 

EC = electron capture, Ci = curie, g = gram, and MeV = million electron volts; a dash means 

that the entry is not applicable. (See the companion fact sheet on Radioactive Properties, 

Internal Distribution, and Risk Coefficients for explanation of terms and interpretation of 

radiation energies.) Potassium-40 decays by both emitting a beta particle (89%) and electron 

capture (11%). Values are given to two significant figures. 

What Is It? Potassium is a soft, silver-white metal. An important 

constituent of soil, it is widely distributed in nature and is present in all 

plant and animal tissues. Potassium-40 is a naturally occurring 

radioactive isotope of potassium. (An isotope is a different form of an 

element that has the same number of protons in the nucleus but a 

different number of neutrons.) Two stable (nonradioactive) isotopes of 

potassium exist, potassium-39 and potassium-41. Potassium-39 

comprises most (about 93%) of naturally occurring potassium, and potassium-41 accounts for essentially 

all the rest. Radioactive postassium-40 comprises a very small fraction (about 0.012%) of naturally 

occurring potassium. 

Several radioactive isotopes of potassium exist in addition to potassium-40. These isotopes all have halflives of less 

than one day 

so they are not 

of concern for 

Department of 

Energy (DOE) 

environmental 

management 

sites such as 

Hanford. The 

half-life of 

potassium-40 

is 1.3 billion years, and it decays to calcium-40 by emitting a beta particle with no attendant gamma 

radiation (89% of the time) and to the gas argon-40 by electron capture with emission of an energetic 

gamma ray (11% of the time). Potassium-40 is an important radionuclide in terms of the dose associated 

with naturally occurring radionuclides. 

Where Does It Come From? Potassium-40 is present as a very small fraction of naturally occurring 

potassium, which is an element found in large amounts throughout nature. Potassium is the seventh most 

abundant element in the crust of the earth and the sixth most abundant element in solution in the oceans. 

It is present in mineral waters and brines, and in various minerals such as carnallite, feldspar, saltpeter, 

greensand, and sylvite. Potassium is an important constituent of fertile soil and is an essential nutrient for 

plant growth and in the human diet. 

How Is It Used? Potassium metal, which is so soft it can be cut with a knife, is used in photoelectric 

cells. Potassium is one of the most reactive metals in nature, and it forms a number of compounds that 

have many commercial uses. For example, the white solid potassium bromide is used in photography, 

engraving, and lithography. The red crystal potassium chromate and yellow crystal potassium bichromate 

are powerful oxidizing agents used in matches and fireworks, and they are also used to dye textiles and 

tan leather. The white crystal potassium iodide is very soluble in water and is used in photography to 

prepare gelatin emulsions. It is also used in medicine to treat rheumatism and overactive thyroid glands. 

Potassium nitrate is a white solid used in matches, explosives, and fireworks, and it is also used to pickle 

meat. The purple crystal potassium permanganate is used as a disinfectant and germicide and as an 

oxidizing agent in various chemical reactions. The white solid potassium carbonate is used to make glass 

and soft soap. The white solids potassium sulfate and potassium chloride are used to fertilize soil, 

because potassium (along with nitrogen and phosphorous) is an essential element for plant growth. 

Potassium is also an essential element for humans, as a key electrolyte for maintaining basic 

cardiovascular functions; many people take potassium supplements as capsules or tablets. There are no 

specific commercial or medical uses associated with the radioactive properties of potassium-40. 

Potassium-40 
 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Maximum values are given 

for inhalation since no default absorption types were 

provided, and dietary values were used for ingestion. 

Risks are for lifetime cancer mortality per unit intake 

(pCi), averaged over all ages and both genders (10P

-9P

is a billionth, and 10P

-1P2

P

 is a trillionth). Other values, 

including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCiP

-1P

) 

Ingestion 

(pCiP

-1P

) 

Potassium-40 2.1 × 10P

-10P 2.2 × 10P-11P

For more information, see the companion fact sheet 

on Radioactive Properties, Internal Distribution, and 

Risk Coefficients and the accompanying Table 1. 

What’s in the Environment? Potassium is present in the earth’s crust, oceans, and all organic 

material. Its concentration in the earth’s crust is about 15,000 milligrams per kilogram (mg/kg) or 1.5%, 

and its concentration in seawater is about 416 mg per liter (mg/L). Because potassium-40 represents 

0.012% of naturally occurring potassium, its concentration in the earth’s crust is about 

1.8 mg/kg, or 13 picocurie per gram (pCi/g). Potassium binds preferentially to soil, with the 

concentration associated with sandy soil particles estimated to be 15 times higher than in the 

interstitial water (in pore spaces between soil particles); it binds more tightly to loam and 

clay soil, so those concentration ratios are higher (above 50). Together with nitrogen and phosphorous, 

potassium is a major soil fertilizer, so levels of potassium-40 in soils are strongly influenced by fertilizer 

use; it is estimated that about 3,000 Ci of potassium-40 are added annually to U.S. soils. Potassium-40 

behaves in the environment the same as other potassium isotopes, being assimilated into the tissues of all 

plants and animals through normal biological processes. It is the predominant radioactive component in 

human tissues and in most food. For example, milk contains about 2,000 pCi/L of natural potassium-40. 

What Happens to It in the Body? Potassium-40 can be taken into the body by drinking water, eating 

food, or breathing air. Once taken in, potassium-40 behaves in the body in the same manner as other 

potassium isotopes. Humans require potassium to sustain biological processes, with most (including 

potassium-40) being almost completely absorbed upon ingestion, moving quickly from the 

gastrointestinal tract to the bloodstream. The potassium-40 that enters the bloodstream after ingestion or 

inhalation is quickly distributed to all organs and tissues. Potassium-40 is eliminated from the body with 

a biological half-life of 30 days. The potassium content of the body is under strict homeostatic control (in 

which the amount retained is actively regulated by the body to achieve the normal range required for 

system functions), and it is not influenced by variations in environmental levels. Hence, the potassium-40 

content in the body is constant, with an adult male having about 0.1 microcurie or 100,000 pCi. Each year 

this isotope delivers doses of about 18 millirem (mrem) to soft tissues of the body and 14 mrem to bone. 

What Is the Primary Health Effect? Potassium-40 can present both an external and an internal health 

hazard. The strong gamma radiation associated with the electron-capture decay process (which occurs 

11% of the time) makes external exposure to this isotope a concern. While in the body, potassium-40 

poses a health hazard from both the beta particles and gamma rays. Potassium-40 behaves the same as 

ordinary potassium, both in the environment and within the human body – it is an essential element for 

both. Hence, what is taken in is readily absorbed into the bloodstream and distributed throughout the 

body, with homeostatic controls regulating how much is retained or cleared. The health hazard of 

potassium-40 is associated with cell damage caused by the ionizing radiation that results from radioactive 

decay, with the general potential for subsequent cancer induction. 

What Is the Risk? Lifetime cancer mortality 

risk coefficients have been calculated for nearly all 

radionuclides, including potassium-40 (see box at 

right). While ingestion is generally the most 

common type of exposure, the risk coefficients for 

this route are lower than those for inhalation. As 

for other radionuclides, the risk coefficient for tap 

water is about 70% of that for dietary ingestion. In 

addition to risks from internal exposures, an 

external gamma exposure risk also exists for 

potassium-40. To estimate a lifetime cancer 

mortality risk, if it is assumed that 100,000 people 

were continuously exposed to a thick layer of soil 

with an initial average concentration of 1 pCi/g 

potassium-40, then 4 of these 100,000 people 

would be predicted to incur a fatal cancer over 

their lifetime. (This is in comparison to the 

20,000 people from the group predicted to die of cancer from all other causes per the U.S. average.) 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Symbol: Pa

Atomic Number: 91 

(protons in nucleus)

Atomic Weight: 231 

(naturally occurring)

Radioactive Properties of the Key Protactinium Isotope 

and Associated Radionuclides

Radiation Energy (MeV)

Isotope Half-Life 

Natural 

Abundance 

(%)

Specific 

Activity 

(Ci/g)

Decay

Mode Alpha 

(α)

Beta

(β)

Gamma 

(γ)

Pa-231 33,000 yr >99 .048 α 5.0 0.065 0.048 

Ac-227 22 yr 73 α, β 0.068 0.016 < 

Th-227 

(99%) 19 days 31,000 α 5.9 0.053 0.11 

Fr-223 

(1%) 22 min 39 million β - 0.40 0.059 

Ra-223 11 days 52,000 α 5.7 0.076 0.13 

Rn-219 4.0 sec 13 billion α 6.8 0.0063 0.056 

Po-215 0.0018 sec 30 trillion α 7.4 < < 

Pb-211 36 min 25 million β - 0.46 0.051 

Bi-211 2.1 min 420 million α, β 6.6 0.010 0.047 

Tl-207 4.8 min 190 million β - 0.49 0.0022 

Ci = curie, g = gram, and MeV = million electron volts; a “<” means the radiation energy 

is less than 0.001 MeV, and a dash means the entry is not applicable. (See the companion 

fact sheet on Radioactive Properties, Internal Distribution, and Risk Coefficients for an 

explanation of terms and interpretation of radiation energies.) Thorium-227 decays by 

both emitting an alpha particle (1%) and a beta particle (99%). Certain properties of 

additional radionuclides are included here because they accompany the protactinium 

decays. Values are given to two significant figures.

 

What Is It? Protactinium is a malleable, shiny, silver-gray radioactive metal that 

does not tarnish rapidly in air. It has a density greater than that of lead and occurs 

in nature in very low concentrations as a decay product of uranium. There are three 

naturally occurring isotopes, with protactinium-231 being the most abundant. 

(Isotopes are different forms of an element that have the same number of protons in 

the nucleus but a different number of neutrons.) The other two naturally occurring 

isotopes are protactinium-234 and protactinium-234m (the “m” meaning 

metastable), both of which have very short half-lives (6.7 hours and 1.2 minutes, respectively) and occur in 

extremely low concentrations. Protactinium was first identified in 1913 by Kasimir Fajans and O.H. Gohring 

(as the isotope protactinium-234m), and protactinium-231 was identified in 1917. The name comes from the 

Greek work protos (meaning first) and the element actinium, because protactinium is the precursor of actinium. 

Of the 20 known isotopes of protactinium, only protactinium-231 has a half-life greater than one year and is a 

concern for Department of Energy (DOE) environmental management sites. The half-lives of all other 

protactinium isotopes 

are less than a month. 

Protactinium-231 is a 

decay product of 

uranium-235 and is 

present at sites that 

processed uranium 

ores and associated 

wastes. This isotope 

decays by emitting an 

alpha particle with a 

half-life of 33,000 

years to actinium-227, 

which has a half-life 

of 22 years and 

decays by emitting an 

alpha or beta particle. 

Actinium-227 and its 

decay products are 

included with the list 

of radionuclides 

associated with 

protactinium-231 in 

the table to the right 

for completeness, as 

these radionuclides 

are typically present 

with protactinium-231. Much of the hazard associated with protactinium-231 is attributable to actinium-227. 

Where Does It Come From? Protactinium is widely distributed in very small amounts in the earth’s crust, 

and it is one of the rarest and most expensive naturally occurring elements. It is present in uranium ores at a 

concentration of about 1 part protactinium to 3 million parts uranium. Of the three naturally occurring isotopes, 

protactinium-231 is a decay product of uranium-235, and protactinium-234 and protactinium-234m are decay 

products of uranium-238. Essentially all (99.8%) of the decays of thorium-234, which is the immediate decay 

product of uranium-238, are to protactium-234m; only about 0.2% are to protactinium-234. (See the companion 

fact sheets on Uranium and Natural Decay Series for additional information.) 

How Is It Used? There are no industrial or commercial uses of protactinium due to its scarcity, expense, and 

radiotoxicity. Its only uses are associated with basic scientific research activities. 

Protactinium 
What’s in the Environment? Protactinium is naturally present in soil, rocks, surface water, groundwater, 

plants, and animals in very low concentrations – on the order of one part per trillion, or 

0.1 picocuries (pCi)/g. Higher levels are present in uranium ores and other geologic materials. 

Essentially all naturally occurring protactinium is present as protactinium-231. Protactinium 

preferentially adheres quite well to soil, and the concentration associated with sandy soil 

particles is typically 550 times higher than in interstitial water (water in the pore space between 

the soil particles); concentration ratios are even higher (about 2,000 and above) for loam and 

clay soils. Protactinium is generally not a major contaminant at DOE sites and is not a concern for groundwater. 

What Happens to It in the Body? Protactinium can be taken into the body by eating food, drinking water, or 

breathing air. When protactinium is inhaled, a significant fraction can move from the lungs through the blood to 

other organs, depending on the solubility of the compound. Gastrointestinal absorption from food or water is a 

likely source of internally deposited protactinium in the general population. Most of the protactinium taken in 

by ingestion will promptly leave the body in feces; only about 0.05% of the amount ingested is absorbed from 

the gastrointestinal tract into the bloodstream. After leaving the intestine or lung, about 40% of the protactinium 

that does enter the bloodstream deposits in the skeleton, about 15% deposits in the liver, about 2% deposits in 

the kidneys, and the rest is excreted. The biological half-life in the skeleton is about 50 years. Of the 

protactinium deposited in the liver, 70% is assumed to be retained with a biological half-life of 10 days, with the 

remaining 30% having a biological half-life of 60 days. Of the protactinium deposited in the kidneys, 20% is 

assumed to be retained with a biological half-life of 10 days, with the remaining 80% having a biological halflife of 60 days. (This information is per simplified models that do not reflect intermediate redistribution.) 

What Are the Primary Health Effects? Protactinium is generally a health hazard only if it is taken into the 

body, although there is a small external risk associated with the gamma rays emitted by protactinium-231 and a 

number of short-lived decay products of actinium-227. The main means of exposure are ingestion of food and 

water containing protactinium and inhalation of protactinium-contaminated dust. Ingestion is generally the 

exposure of concern unless there is a nearby source of contaminated airborne dust. Because protactinium is 

taken up in the body much more readily if inhaled rather than ingested, both exposure routes can be important. 

The major health concern is cancer resulting from the ionizing radiation emitted by protactinium deposited in 

the skeleton, liver, and kidneys. The health risks 

associated with protactinium-234m are included with those 

for uranium-238 (see the companion fact sheet on 

Uranium). Protactinium-234m decays by emitting an 

energetic beta particle so precautions against this radiation 

are needed when handling uranium; for example, heavy 

rubber gloves are worn to protect the hands and forearms. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including protactinium (see box at right). 

The inhalation risk factor for protactinium-231 represents 

one of the largest risk factors for any radionuclide. 

Actinium-227 and its decay products account for more 

than 80% of this inhalation risk. While the risk factor for 

ingestion is much lower than for inhalation, ingestion is 

generally the most common means of entry into the body. 

Similar to other radionuclides, the risk coefficient for tap 

water is about 75% of that shown for dietary ingestion. 

In addition to risks from internal exposures, there is a risk from external gamma exposure to protactinium-231. 

Using the external gamma risk coefficients to estimate lifetime cancer mortality risks, if it is assumed that 

100,000 people were continuously exposed to a thick layer of soil with an initial average concentration of 

1 pCi/g protactinium-231, then 8 of these 100,000 people would be predicted to incur a fatal cancer. (This is in 

comparison to the 20,000 people from this group predicted to die of cancer from all other causes per the 

U.S. average.) As for internal exposures, much of this risk is from actinium-227 and its decay products. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and absorption. Maximum values are 

given for inhalation as no default absorption types 

were provided, and dietary values were used for 

ingestion. These values include the contributions 

from the actinium-227 and its short-lived decay 

products. Risks are for lifetime cancer mortality per 

unit intake (pCi), averaged over all ages and both 

genders (10-9 is a billionth, and 10-12 is a trillionth). 

Other values, including for morbidity, are also 

available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Pa-231 2.5 × 10-7 6.0 × 10-10

For more information, see the companion fact sheet 

on Radioactive Properties, Internal Distribution, 

and Risk Coefficients and accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Symbol: Ra

Atomic Number: 88 

(protons in nucleus)

Atomic Weight: 226 

(naturally occurring)

Radioactive Properties of Key Radium Isotopes and Associated Radionuclides

Radiation Energy (MeV)

Isotope Half-Life 

Natural 

Abundance 

(%)

Specific 

Activity 

(Ci/g)

Decay

Mode Alpha 

(α)

Beta

(β)

Gamma 

(γ)

Ra-226 1,600 yr >99 1.0 α 4.8 0.0036 0.0067 

Rn-222 3.8 days 160,000 α 5.5 < < 

Po-218 3.1 min 290 million α 6.0 < < 

Pb-214 27 min 33 million β - 0.29 0.25 

Bi-214 20 min 45 million β - 0.66 1.5 

Po-214 0.00016 

sec 

330 trillion α 7.7 < < 

Pb-210 22 yr 77 β - 0.038 0.0048 

Bi-210 5.0 days 130,000 β - 0.39 - 

Po-210 140 days 4,500 α 5.3 < < 

Ra-228 5.8 yr <<1 280 β - 0.017 < 

Ac-228 6.1 hr 2.3 million β - 0.48 0.97 

Th-228 1.9 yr 830 α 5.4 0.021 0.0033 

Ra-224 3.7 days 160,000 α 5.7 0.0022 0.010 

Rn-220 56 sec 930 million α 6.3 < < 

Po-216 0.15 sec 350 billion α 6.8 < < 

Pb-212 11 hr 1.4 million β - 0.18 0.15 

Bi-212 61 min 15 million α, β 2.2 0.47 0.19 

Po-212 

 (64%) 

0.00000031 

sec 

180,000 

trillion α 8.8 - - 

Tl-208 

 (36%) 3.1 min 300 million β - 0.60 3.4 

Ci = curie, g = gram, and MeV = million electron volts; a “<” means the radiation energy 

is less than 0.001 MeV, and a dash means the entry is not applicable. (See the companion 

fact sheet on Radioactive Properties, Internal Distribution, and Risk Coefficients for an 

explanation of terms and interpretation of radiation energies.) Bismuth-212 decays by both 

emitting an alpha particle (36%) and a beta particle (64%). Certain properties of 

additional radionuclides are included here because they accompany the radium decays. 

Values are given to two significant figures. 

 

What Is It? Radium is a radioactive element that occurs naturally in very low concentrations 

(about one part per trillion) in the earth’s crust. Radium in its pure form is a silvery-white 

heavy metal that oxidizes immediately upon exposure to air. Radium has a density about onehalf that of lead and exists in nature mainly as radium-226, although several additional isotopes 

are present. (Isotopes are different forms of an element that have the same number of protons in 

the nucleus but a different number of neutrons.) Radium was first discovered in 1898 by Marie 

and Pierre Curie, and it served as the basis for identifying the activity of various radionuclides. 

One curie of activity equals the rate of radioactive decay of one gram (g) of radium-226. 

Of the 25 known isotopes of radium, only two – radium-226 and radium-228 – have half-lives greater than one year and are of 

concern for Department of 

Energy environmental 

management sites. 

Radium-226 is a radioactive 

decay product in the 

uranium-238 decay series 

and is the precursor of 

radon-222. Radium-228 is 

a radioactive decay product 

in the thorium-232 decay 

series. Both isotopes give 

rise to many additional 

short-lived radionuclides, 

resulting in a wide spectrum 

of alpha, beta and gamma 

radiations. Lead-210, 

which has a 22-year halflife, is included in the list of 

short-lived radionuclides 

associated with radium-226 

for completeness, as this 

isotope and its short-lived 

decay products are typically 

present with radium-226. 

Radium-226 decays slowly 

(half-life of 1,600 years) by 

emitting an alpha particle. 

Radium-228 has a much 

shorter half-life (5.8 years) 

and decays by emitting a 

beta particle. While 

radium-226 poses a hazard 

due to its long half-life, 

radium-228 poses a longterm hazard only if its 

parent (thorium-232) is 

present. 

Where Does It Come 

From? Radium is widely 

distributed in small amounts 

in the earth’s crust. It is present in all uranium and thorium minerals; its concentration in uranium ores is about one part 

radium to 3 million parts uranium. The chemical properties of radium are similar to those of barium, and the two substances 

are removed from uranium ore by precipitation and other chemical processes. Originally, radium was obtained from the rich 

pitchblende ore found in Bohemia. The carnotite sands of Colorado furnish some radium, but richer ores are found in the 

Republic of Zaire and the Great Lake Region of Canada. Radium is a major contaminant in mine and milling wastes, such as 

uranium mill tailings, and is present in various radioactive wastes associated with past uranium processing activities. 

How Is It Used? Radium-226 is the only radium isotope used commercially. Historically, the main use of radium has been 

as a component in luminous paint used on the dials of watches, clocks, and other instruments, although it is no longer used for

Radium
this purpose. Radium is currently used in brachytherapy to treat various types of cancer. (Brachytherapy is a method of 

radiation treatment in which sealed sources are used to deliver a radiation dose at a distance of up to a few centimeters by 

surface, intracavitary, or interstitial application.) 

What’s in the Environment? Essentially all naturally occurring radium is present as radium-226. Radium exists naturally 

in soil, rocks, surface water, groundwater, plants, and animals in generally low concentrations – on the order of one part per 

trillion, or 1 picocurie (pCi)/g. Higher levels are present in uranium ores and other geologic materials. 

Because of the separation process used to extract uranium from ores, radium-226 is a major contaminant in 

uranium mill tailings. The concentration of radium in plants is typically about 0.03 (or 3%) of that in soil. 

However, Brazil nuts in areas of high natural radium have much higher (orders of magnitude) concentration 

ratios. The average concentration of radium in food has been estimated at less than 0.01 to 0.03 pCi/g. 

Radium preferentially adheres well to soil particles, with concentrations in sandy soil generally on the order 

of 500 times higher than in interstitial water (water in the pore spaces between soil particles); it is even less mobile in clay 

soils, with concentration ratios over 9,000. The maximum contaminant level developed by the U.S. Environmental Protection 

Agency for radium (as radium-226 and radium-228, combined) in drinking water supplies is 5 pCi per liter (pCi/L). 

What Happens to It in the Body? Radium can be taken into the body by eating food, drinking water, or breathing air. 

Most of the radium taken in by ingestion (about 80%) will promptly leave the body in feces. The remaining 20% enters the 

bloodstream and is carried to all parts of the body. Inhaled radium can remain in the lungs for several months and will 

gradually enter the bloodstream and be carried throughout the body. The metabolic behavior of radium in the body is similar to

that of calcium. For this reason, an appreciable fraction is preferentially deposited in bone and teeth. The amount in bone 

decreases with time from the exposure, generally dropping below 10% in a few months to 1% and less in a few years. Release 

from the bone is slow, so a portion of inhaled and ingested radium will remain in the bones throughout a person’s lifetime. 

What Are the Primary Health Effects? Radium poses an external as well as an internal health hazard. The strong 

external gamma radiation associated with several short-lived decay products of radium-226 and radium-228 makes external 

exposure a concern, and shielding is often needed to handle waste and other materials containing large concentrations of these 

radionuclides. The majority of epidemiological data on the health effects of radium-226 and radium-228 in humans comes 

from studies of radium dial painters, radium chemists, and technicians exposed through medical procedures in the early 1900s. 

These studies, as well as studies on experimental animals, indicate that chronic exposure to radium can induce bone sarcomas. 

The minimum latency period is seven years after the first exposure, but tumors can continue to appear throughout a lifetime. 

The inhalation risk is associated primarily with radium decay 

products, i.e., radon and its short-lived daughters. Each of the 

two radium isotopes decays into a gaseous radon isotope. 

Radon-222 is a short-lived decay product of radium-226, and 

radon-220 is a short-lived decay product of radium-228. The 

primary hazard associated with radon arises from the inhalation 

of its short-lived decay products, which are charged ions that 

readily attach to dust particles. These particles can be inhaled 

into the lungs and deposited on the mucous lining of the 

respiratory tract. Unattached decay products tend to be inhaled 

deeper into the lungs where the residence time is longer. When 

alpha particles are then emitted within the lung, the cells lining 

the airways can be damaged, potentially leading to lung cancer 

over time. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all radionuclides, 

including radium (see box at right). The ingestion and 

inhalation coefficients for radium-226 and radium-228 are 

generally comparable. While ingestion is the most common 

means of radium entry into the body, risk coefficients for that 

exposure route are lower than for inhalation. Similar to other 

radionuclides, the risk coefficients for tap water are about 75% of those shown for dietary ingestion. 

In addition to risks from internal exposures, a risk from external gamma exposure is associated with these two isotopes. Using

the external gamma risk coefficients to estimate lifetime cancer mortality risks, if it is assumed that 100,000 persons were 

continuously exposed to a thick layer of soil with an initial average concentration of 1 pCi/g, then 40 of these 100,000 people 

would be predicted to incur a fatal cancer if the soil contained radium-226, and 7 if it contained radium-228. (This is in 

comparison to the 20,000 people from the group predicted to die of cancer from all other causes per the U.S. average.) These 

risks are associated with the gamma rays emitted by various decay products of these two radium isotopes. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and absorption. Recommended default 

absorption types were used for inhalation, and dietary 

values were used for ingestion. These values include the 

contributions from the short-lived radium decay 

products. (See text for information on the external 

exposure pathway.) Risks are for lifetime cancer 

mortality per unit intake (pCi), averaged over all ages 

and both genders (10-9 is a billionth). Other values, 

including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Ra-226 2.4 × 10-8 2.9 × 10-9

Ra-228 9.0 × 10-8 1.3 × 10-9

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Sm

Atomic Number: 62 

(protons in nucleus)

Atomic Weight: 150 

(naturally occurring)

Radioactive Properties of Key Samarium Isotopes 

Radiation Energy (MeV)

Isotope Half-Life 

(yr)

Specific

Activity

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Sm-146 100,000,000 0.000024 α 2.5 - - 

Sm-151 90 27 β - 0.020 < 

Ci = curie, g = gram, and MeV = million electron volts; a “<” means the 

radiation energy is less than 0.001 MeV, and a dash means the entry is not 

applicable. (See the companion fact sheet on Radioactive Properties, Internal 

Distribution, and Risk Coefficients for an explanation of terms and 

interpretation of radiation energies.) The three naturally occurring samarium 

isotopes are not included here. Values are given to two significant figures. 

What Is It? Samarium is a very hard, silvery-white metal. It is a member 

of the lanthanide series and ignites in air at high temperatures (150°C). In 

nature, samarium occurs as seven isotopes, three of which are radioactive 

with extremely long half-lives. (Isotopes are different forms of an element 

that have the same number of protons in the nucleus but a different number 

of neutrons.) The isotopes and their natural occurrences are samarium-144 

(3.1%), samarium-147 (15%), samarium-148 (11%), samarium-149 (14%), 

samarium-150 (7.4%), samarium-152 (27%), and samarium-154 (23%). The three naturally occurring 

radioactive isotopes and their half-lives are samarium-147 (110 billion years), samarium-148 (8,000 

trillion years), and samarium-149 (10,000 trillion years). Their extremely long half-lives make these 

three radioactive isotopes essentially indistinguishable from the stable (nonradioactive) isotopes. 

Nine major radioactive isotopes of samarium exist in addition to the three naturally occurring ones. Of 

these nine, only two have half-lives long enough to warrant concern at Department of Energy (DOE) 

environmental management sites such as Hanford. The half-lives of the other samarium isotopes are less 

than one year. Radioactive samarium isotopes are produced by the fissioning of uranium and plutonium, 

generally in nuclear 

reactors. The shorter-lived 

samarium isotopes (those 

with half-lives less than one 

year) produced more than 

20 years ago have long 

since decayed away. Of the 

two longer-lived isotopes, 

samarium-146 is not a 

major concern at DOE sites 

because of its low 

prevalence and very long 

half-life (and hence low 

specific activity) that limits 

its radioactive hazards. The isotope of most concern is samarium-151, a radionuclide with a half-life of 

90 years that decays by emitting a relatively low-energy beta particle. 

Where Does It Come From? Samarium is found in a variety of minerals such as monazite, bastnasite, 

cerite, gadolinite, and samarskite. These minerals contain different mixes of rare earth metals, which are 

the elements from lanthanum through lutetium in the periodic table. Samarium generally makes up less 

than a few percent of these minerals. China currently produces the vast majority of the rare earth metals, 

totaling about 70,000 metric tons (MT) annually. The United States comes in a distant second, producing 

around 5,000 MT annually, and samarium makes up only a small fraction of this amount. 

Samarium-151 is produced by nuclear fission. When an atom of uranium-235 (or another fissile nuclide) 

fissions, it generally splits asymmetrically into two large fragments – fission products with mass numbers 

in the range of about 90 and 140 – and two or three neutrons. (The mass number is the sum of the 

number of protons and neutrons in the nucleus of the atom.) The fission yield of samarium-151is about 

0.4%. That is, about four atoms of samarium-151 are produced per 1,000 fissions. Samarium-151 is a 

significant radionuclide in spent nuclear fuel, high-level radioactive wastes resulting from the processing 

of spent nuclear fuel, and radioactive wastes associated with the operation of nuclear reactors and fuel 

reprocessing plants. In addition, samarium-149 is a very good absorber of neutrons and is therefore often 

used in nuclear reactor control rods; capture of two neutrons by samarium-149 results in samarium-151. 

Samarium 
How Is It Used? Samarium has a variety of commercial uses. It is used as a catalyst for certain organic 

reactions and in pyrophoric alloys for cigarette lighter flints. Samarium oxide is also used in special 

infrared absorbing glass and in the cores of carbon arc-lamp electrodes. An alloy of samarium with cobalt 

is used to make a magnetic material with the highest resistance to demagnetization of any known 

material. Because one isotope (samarium-149) is a very good neutron absorber, this element has also 

been used as a component in nuclear reactor control rods 

What’s in the Environment? Samarium is present in the earth’s crust at a concentration of about 

6 milligram per kilogram (mg/kg), while its concentration in seawater is about 0.00045 micrograms per 

liter (μg/L). Trace amounts of samarium-151 are present in soil around the globe from radioactive 

fallout. It can also be present at certain nuclear facilities, such as reactors and spent fuel reprocessing 

plants. 

The transport of samarium in the environment is strongly influenced by its chemical 

form. It is generally one of the less mobile radioactive metals in soil, although certain 

forms can move downward with percolating water some distance to underlying layers of 

soil. Samarium preferentially adheres well to soil particles, with concentrations in sandy 

soil generally more than 200 times higher than in interstitial water (the water in the pore 

space between soil particles); it is even less mobile in clay soils, with concentration ratios over 1,000. 

Thus, samarium is generally not a major contaminant in groundwater at DOE sites. 

What Happens to It in the Body? Samarium can be taken into the body by eating food, drinking 

water, or breathing air. Gastrointestinal absorption from food or water is the principal source of internally 

deposited samarium in the general population. Samarium is not well absorbed into the body after intake. 

Only about 0.05% of the amount ingested being absorbed into the bloodstream through the digestive tract, 

with the rest being excreted. Of the samarium that reaches the blood, 45% is deposited in the liver and 

another 45% is deposited on the surface of the bone, where it can irradiate the bone-forming cells; the 

remaining 10% is excreted. The samarium that is internally deposited is retained in the body with a 

biological half-life of almost 10 years (3,500 days). (This information is per simplified models that do 

not reflect intermediate redistribution.) 

What Are the Primary Health Effects? 

Samarium-151 is a health hazard only if it is 

taken into the body. External gamma exposure is 

not a major concern because samarium-151 emits 

only a small amount of low-energy gamma 

radiation. While in the body, samarium-151 

poses a health hazard from both the beta particles 

and gamma rays it emits, and the main health 

concern is associated with the increased 

likelihood of inducing cancer in the liver and 

bone. 

What Is the Risk? Lifetime cancer mortality 

risk coefficients have been calculated for nearly 

all radionuclides, including samarium (see box at 

right). While the coefficients for ingestion are 

somewhat lower than for inhalation, ingestion is 

generally the most common means of samarium 

entry into the body. Similar to other 

radionuclides, the risk coefficients for tap water 

are about 70% of those for dietary ingestion. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Maximum values are given 

for inhalation since no default absorption types were 

provided, and dietary values were used for ingestion 

Risks are for lifetime cancer mortality per unit intake 

(picocurie, pCi), averaged over all ages and both 

genders (10-9 is a billionth, and 10-12 is a trillionth). 

Other values, including for morbidity, are also 

available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Samarium-146 1.2 × 10-8 4.0 × 10-11 

Samarium-151 8.6 × 10-12 4.6 × 10-13 

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Se

Atomic Number: 34

(protons in nucleus)

Atomic Weight: 79

(naturally occurring)

Radioactive Properties of the Key Selenium Isotope

Radiation Energy (MeV)

Isotope 

HalfLife 

(yr)

Specific

Activity

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Se-79 650,000 0.070 β - 0.056 - 

Ci = curie, g = gram, and MeV = million electron volts; a dash means the 

entry is not applicable. (See the companion fact sheet on Radioactive 

Properties, Internal Distribution, and Risk Coefficients for an explanation 

of terms and interpretation of radiation energies.) Values are given to two 

significant figures.

What Is It? Selenium is a non-metallic mineral that resembles sulfur and can 

exist as a gray crystal, red powder, or vitreous black form. It occurs in nature 

as six stable isotopes. (Isotopes are different forms of an element that have the 

same number of protons in the nucleus but a different number of neutrons.) 

Selenium-80 is the most prevalent, comprising about half of natural selenium. 

The other five stable isotopes and their relative abundances are selenium-74 

(0.9%), selenium-76 (9.4%), selenium-77 (7.6%), selenium-78 (24%), and 

selenium-82 (8.7%). 

Of the nine major radioactive 

selenium isotopes, only one – 

selenium-79 – has a half-life long 

enough to warrant concern at 

U.S. Department of Energy (DOE) 

environmental management sites 

such as Hanford. The half-life of 

selenium-75 is 120 days and the 

half-lives of all other isotopes are 

less than eight hours. Selenium-79 

decays by emitting a beta particle 

with a half-life of 650,000 years with 

no attendant gamma radiation. Selenium-79 is present in spent nuclear fuel and the wastes resulting from 

reprocessing this fuel. The low specific activity and relatively low energy of its beta particle limits the 

radioactive hazards of this isotope. 

Where Does It Come From? Selenium occurs naturally as a trace element in most soils, rocks and waters, 

and it accompanies sulfur in volcanic effluents. Higher concentrations are present in soils near volcanoes and in 

minerals such as clausthalite, naumannite, tiemannite, and senenosulfur, but selenium does not occur in 

concentrated deposits. Although it cannot be economically recovered directly from the earth, selenium is 

commonly generated as a byproduct of electrolytic copper refining. Selenium-79 is produced by nuclear fission. 

When an atom of uranium-235 (or other fissile nuclide) fissions, it generally splits asymmetrically into two large 

fragments – fission products with mass numbers in the range of about 90 and 140 – and two or three neutrons. 

(The mass number is the sum of the number of protons and neutrons in the nucleus of the atom.) Selenium-79 is 

one such fission product, and it is produced with a yield of about 0.04%. That is, four atoms of selenium-79 are 

produced per 10,000 fissions. Selenium-79 is a component of spent nuclear fuel, high-level radioactive wastes 

resulting from processing spent fuel, and radioactive wastes associated with the operation of nuclear reactors 

and fuel reprocessing plants. 

How Is It Used? Selenium has a number of industrial applications. It is used in photocells, solar cells, and 

exposure meters for photography and as a toner for photographic and copier uses. Selenium is also used to 

decolorize glass and to impart a scarlet red color to clear glass, glazes and enamels. Additional uses are in 

rectifiers and other electronic equipment, as a vulcanizing agent in processing of rubber, and as an additive in 

stainless steel and other alloys. Selenium compounds are used as an insecticide to combat insects that attack 

cultivated plants and also to treat various skin and scalp conditions (such as dandruff). 

What’s in the Environment? Selenium is naturally present in the earth’s crust at an average concentration of 

about 0.05 milligram per kilogram (mg/kg), and its concentration in seawater is about 

0.45 microgram per liter (μg/L). Trace amounts of selenium-79 are present in soil around the 

globe from radioactive fallout. It can also be present at certain nuclear facilities, such as reactors 

and facilities that process spent nuclear fuel. The highest concentrations of selenium-79 at the 

Hanford site are in areas that contain waste from processing irradiated fuel, such as in tanks in the 

central portion of the site. Selenium is generally one of the less mobile radioactive metals in soil 

as it preferentially adheres well to soil particles. Concentrations in sandy soil are estimated to be 150 times 

Selenium 
higher than in interstitial water (in pore spaces between the soil particles), and it is even less mobile in clay soils 

where concentration ratios exceed 700. The low fission yield of selenium-79 limits its presence at DOE sites, so 

it is generally not a major groundwater contaminant at these sites. Its concentration in plants is typically 0.025 

(or 2.5%) of that in soil, although levels are much higher in seleniferous plants. Certain foods are especially 

high in selenium, such as garlic. 

What Happens to It in the Body? Selenium can be taken into the body by eating food, drinking water, or 

breathing air. Children, and to a lesser extent adults, can also be exposed by ingesting soils. Gastrointestinal 

absorption is the principal source of internally deposited selenium in the general population. About 80% of 

selenium incorporated in food and soluble inorganic compounds are absorbed from the gastrointestinal tract into 

the bloodstream. However, elemental selenium and selenides are relatively inactive biologically, and only about 

5% of these forms are absorbed from the intestines. After reaching the blood, selenium selectively deposits in 

the liver (15%), kidneys (5%), spleen (1%) and pancreas (0.5%). The remainder is deposited uniformly 

throughout all other organs and tissues. Of the selenium deposited in any organ or tissue, 10% is retained with a 

biological half-life of 3 days, 40% is retained with a biological half-life of 30 days, and 50% is retained with a 

biological half-life of 150 days. (This information is per simplified models that do not reflect intermediate 

redistribution.) Selenium does not appear to be absorbed through the skin. 

What Are the Primary Health Effects? Selenium is a health hazard only if it is taken into the body. 

External gamma exposure is not a concern because selenium-79 decays by emitting a beta particle with no 

gamma radiation. While inside the body, this selenium poses a health hazard from the beta particles emitted 

during its radioactive decay, and the main concern is associated with the increased likelihood of inducing 

cancer. Selenium (nonradioactive) is an essential element for humans that is required for normal enzyme 

function and is necessary for normal growth and metabolism. The recommended daily intakes are 55 and 70 μg 

for women and men, respectively. It is also an important antioxidant that helps protect cell membranes and 

prevent free radical generation. Some studies suggest that selenium decreases the risk of certain forms of cancer 

and diseases of the heart and blood vessels. Chronic exposures to moderate levels can cause selenosis, which is 

characterized by hair loss, brittle nails, and neurological abnormalities. A common cause of selenium poisoning 

in humans is overdosing on vitamins. Exposure to selenium or selenium dioxide in air can result in respiratory 

irritation, bronchial spasms, and coughing. Selenium can cause “blind staggers” in animals that drink water or 

eat plants or animals in areas of high selenium. The U.S. Environmental Protection Agency (EPA) has 

classified one specific form of selenium, selenium sulfide, as a probable human carcinogen. Selenium sulfide is 

not present in foods and is biologically different from the types of selenium compounds found in food or in the 

environment. Selenium may be protective against the toxic effects of mercury exposures. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including selenium (see box at right).

Ingestion is the exposure route of most concern because 

this risk coefficient is about three times higher than that 

for inhalation, and ingestion is generally the most 

common means of radioactive selenium intake. Similar 

to other radionuclides, the risk coefficient for tap water 

is about 75% of that for dietary ingestion. The EPA 

toxicity value for estimating potential for non-cancer 

effects from oral exposure to chemicals is called a 

reference dose (RfD), which is an estimate of the highest 

dose that can be taken in every day without causing an 

adverse non-cancer effect. The RfD for selenium, shown 

at right, is based on studies of individuals with selenosis 

who live in areas with naturally high concentrations of 

selenium in soil. Those exposures were estimated to be 

at levels about 70 to 100 times higher than the 

recommended daily intake. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Recommended default 

absorption types were used for inhalation, and dietary 

values were used for ingestion. Risks are for lifetime 

cancer mortality per unit intake (picocurie, pCi), 

averaged over all ages and both genders (10-12 is a 

trillionth). Other values, including for morbidity, are 

also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Selenium-79 2.3 × 10-12 6.7 × 10-12

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 

Chemical Toxicity Value 

Form Non-Cancer Effect: Oral RfD

Selenium and 

compounds 

0.005 

mg/kg-day 
Argonne National Laboratory, EVS Human Health Fact Sheet, November 2006 

Symbol: Sr

Atomic Number: 38 

(protons in nucleus)

Atomic Weight: 88 

(naturally occurring)

Radioactive Properties of the Key Strontium Isotope

and an Associated Radionuclide 

Radiation Energy (MeV)

Isotope HalfLife 

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Sr-90 29 yr 140 β - 0.20 - 

Y-90 64 hr 550,000 β - 0.94 < 

Ci = curie, g = gram, and MeV = million electron volts; a dash means the entry is 

not applicable, and a “<” means the radiation energy is less than 0.001 MeV. 

(See the companion fact sheet on Radioactive Properties, Internal Distribution, 

and Risk Coefficients for an explanation of terms and interpretation of radiation 

energies.) Properties of yttrium-90 are included here because this radionuclide 

accompanies strontium decays. Values are given to two significant figures. 

What Is It? Strontium is a soft, silver-gray metal that occurs in nature as four 

stable isotopes. (Isotopes are different forms of an element that have the same 

number of protons in the nucleus but a different number of neutrons.) 

Strontium-88 is the most prevalent form, comprising about 83% of natural 

strontium. The other three stable isotopes and their relative abundance are 

strontium-84 (0.6%), strontium-86 (9.9%), and strontium-87 (7.0%). Strontium 

is present in nature chiefly as celestite (SrSO4) and strontianite (SrCO3), and it 

comprises about 0.025% of the earth’s crust. 

Sixteen major radioactive isotopes of strontium exist, but only strontium-90 has a half-life sufficiently long 

(29 years) to warrant 

concern for nuclear facilities 

such as the U.S. Department 

of Energy Hanford site. The 

half-lives of all other 

strontium radionuclides are 

less than 65 days. 

Strontium-90 decays to 

yttrium-90 by emitting a 

beta particle, and yttrium-90 

decays by emitting a more 

energetic beta particle with a 

half-life of 64 hours to 

zirconium-90. The main 

health concerns for 

strontium-90 are related to the energetic beta particle from yttrium-90. 

Where Does It Come From? While four stable isotopes of strontium occur naturally, strontium-90 is 

produced by nuclear fission. When an atom of uranium-235 (or other fissile nuclide) fissions, it generally 

splits asymmetrically into two large fragments – fission products with mass numbers in the range of about 

90 and 140 – and two or three neutrons. (The mass number is the sum of the number of protons and 

neutrons in the nucleus of the atom.) Strontium-90 is such a fission product, and it is produced with a yield 

of about 6%. That is, about six atoms of strontium-90 are produced per 100 fissions. Strontium-90 is a 

major radionuclide in spent nuclear fuel, high-level radioactive wastes resulting from processing spent 

nuclear fuel, and radioactive wastes associated with the operation of reactors and fuel reprocessing plants. 

How Is It Used? Strontium has a variety of commercial and research uses. It has been used in certain 

optical materials, and it produces the red flame color of pyrotechnic devices such as fireworks and signal 

flares. Strontium has also been used as an oxygen eliminator in electron tubes and to produce glass for 

color television tubes. In addition, strontium-90 has been used as an isotopic energy source in various 

governmental research applications, including in radiothermal generators to produce electricity for a variety 

of purposes including devices to power remote weather stations, navigational buoys, and satellites. 

What’s in the Environment? Beyond the four stable isotopes naturally present in soil, strontium-90 is 

also present in surface soil around the world as a result of fallout from past atmospheric 

nuclear weapons tests. Current strontium-90 levels in surface soil typically range from 0.01 to 

1 picocurie per gram (pCi/g), reflecting various rainfall and wind patterns, elevation, and 

terrain; most levels fall between 0.05 and 0.5 pCi/g, with 0.1 pCi/g as a general average. 

Strontium-90 is relatively mobile and can move down through soil with percolating water to groundwater. 

Environmental transport of strontium is strongly influenced by its chemical form. Strontium preferentially 

adheres to soil particles, and the amount in sandy soil is typically about 15 times higher than in interstitial 

Strontium
water (in the pore spaces between soil particles); concentration ratios are typically higher (110) in clay soil. 

As a note, many years ago the U.S. Environmental Protection Agency (EPA) established a maximum 

contaminant level for strontium-90 in public drinking water supplies. That value based on extant dosimetry 

models is 8 pCi per liter (pCi/L). The value using current, improved dosimetry models would be 36 pCi/L. 

What Happens to It in the Body? Strontium can be taken into the body by eating food, drinking water, 

or breathing air. Gastrointestinal absorption from food or water is the principal source of internally 

deposited strontium in the general population. On average, 30 to 40% of ingested strontium is absorbed 

into the bloodstream. The amount absorbed tends to decrease with age, and is higher (about 60%) in 

children in their first year of life. Adults on fasting and low-calcium diets can also increase intestinal 

absorption to these levels, as the body views strontium as a replacement for calcium. Strontium behaves 

similarly to calcium (although it is not homeostatically controlled, i.e., the body does not actively regulate 

levels within the cells), but living organisms generally use and retain it less effectively. For adults, about 

31% of the activity entering the blood (plasma) from the gastrointestinal tract is retained by bone surfaces; 

the remainder goes to soft tissues or is excreted in urine and feces. Much of the activity initially deposited 

on bone surfaces is returned to plasma within a few days based on an updated biokinetic model that 

accounts for redistribution in the body. About 8% of the ingested activity remains in the body after 

30 days, and this decreases to about 4% after 1 year. This activity is mainly in the skeleton. 

What Are the Primary Health Effects? Strontium is a health hazard only if it is taken into the body. 

External gamma exposure is not a major concern because strontium-90 emits no gamma radiation and its 

decay product yttrium-90 emits only a small amount. Strontium-90 concentrates in bone surfaces and bone 

marrow, and its relatively long radioactive half-life (29 years) make it one of the more hazardous products 

of radioactive fallout. The health effects associated with strontium-90 were studied concurrent with 

development of the atomic bomb during World War II by the Manhattan Engineer District. Bone tumors 

and tumors of the blood-cell forming organs are the main health concern. These tumors are associated with 

the beta particles emitted during the radioactive decay of strontium-90 and yttrium-90. 

What Is the Risk? Lifetime cancer mortality risk coefficients have been calculated for nearly all 

radionuclides, including strontium-90 (see box at 

right). Most of the risk is associated with the highenergy beta particle emitted by yttrium-90. 

Although the risk coefficient for ingestion is lower 

than for inhalation, ingestion is generally the most 

common way this radioisotope enters the body. 

Similar to other radionuclides, the risk coefficient 

for tap water is about 80% of that for dietary 

ingestion. In addition to potential radiogenic 

effects, strontium has been shown to inhibit 

calcification and cause bone deformities in 

animals, notably at high doses. The EPA toxicity 

value for estimating the potential for non-cancer 

effects from oral exposure is termed a reference 

dose (RfD), which is an estimate of the highest 

dose that can be taken in every day without causing 

an adverse non-cancer effect. The RfD for 

ingested strontium (see box at right) based on 

rachitic bone effects was developed by studying 

test animals given relatively high doses over their 

lifetimes, then adjusting and normalizing those 

results to a milligram per kilogram per day (mg/kgday) basis for humans. A noncancer toxicity value 

for inhalation exposure has not been developed. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. The recommended default 

absorption type was used for inhalation, and the dietary 

value was used for ingestion. These values include the 

contribution from the decay product yttrium-90. Risks 

are for lifetime cancer mortality per unit intake (pCi), 

averaged over all ages and both genders (10-9 is a 

billionth, and 10-12 is a trillionth). Other values, 

including for morbidity, are also available. 

 Lifetime Cancer Mortality Risk 

Isotope 

Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Strontium-90 1.0 × 10-10 7.5 × 10-11

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 

Chemical Toxicity Value 

Non-Cancer Effect: Oral RfD

0.6 mg/kg-day
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Tc

Atomic Number: 43 

(protons in nucleus)

Atomic Weight: - 

(not naturally occurring)

Radioactive Properties of Key Technetium Isotopes

Radiation Energy (MeV)

Isotope Half-Life 

(yr)

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Tc-97 2.6 million 0.0014 EC - 0.0056 0.011 

Tc-98 4.2 million 0.00088 β - 0.16 1.4 

Tc-99 210,000 0.017 β - 0.10 - 

EC = electron capture, Ci = curie, g = gram, and MeV = million electron 

volts; a dash means the entry is not applicable. (See the companion fact sheet 

on Radioactive Properties, Internal Distribution, and Risk Coefficients for an 

explanation of terms and interpretation of radiation energies.) Values are 

given to two significant figures.

What Is It? Technetium is a silvery gray metal that looks like platinum 

and tarnishes slowly in moist air. Essentially all the technetium on earth 

has been created by human activities involving fissionable materials. 

Taking its name from the Greek work technetos meaning artificial, it was 

first produced in 1937 by bombarding molybdenum with deuterons (a 

form of hydrogen with a neutron in the nucleus) in a cyclotron. 

There are no stable, i.e., nonradioactive, isotopes of technetium. (Isotopes are different forms of an 

element that have the same number of protons in the nucleus but a different number of neutrons.) Of the 

ten major radioactive isotopes, only three – technetium-97, technetium-98 and technetium-99 – have halflives sufficiently long to warrant concern over time. The half-lives of the other isotopes are less than 

90 days. Only one of the three long-lived isotopes, technetium-99, is produced in sufficient quantities to 

be of concern at Department of Energy (DOE) environmental management sites such as Hanford. This 

fission product decays by emitting a beta particle to produce the stable isotope ruthenium-99. The very 

long half-life (and thus low specific activity) of technetium-99 limits its radioactive hazards. 

Technetium-98 also decays 

by emitting a beta particle 

while technetium-97 decays 

by electron capture. These 

two radionuclides have very 

long half-lives (in excess of 

a million years). An 

additional radionuclide, 

technetium-99m (the “m” 

means metastable), is used 

in medical diagnostic 

procedures. This isotope 

has a half-life of about six 

hours and is a decay product 

of molybdenum-99, a radionuclide with a half-life of 66 hours that also decays by emitting a beta particle. 

Where Does It Come From? Technetium is produced as a result of nuclear transformations, typically 

in a nuclear reactor. When an atom of a fissile nuclide such as uranium-235 fissions, it generally splits 

asymmetrically into two large fragments – fission products with mass numbers in the range of about 90 

and 140 – and two or three neutrons. (The mass number is the sum of the number of protons and neutrons 

in the nucleus of the atom.) Technetium-99 and molybdenum-99 are two such fission products, with a 

relatively high yield of about 6%. That is, about six atoms of each isotope are produced per 100 fissions. 

Technetium-99m is a short-lived decay product of molybdenum-99. (An extremely small amount of 

technetium was created naturally in sustained underground nuclear reactions estimated to have occurred 

about 1.9 billion years ago in Gabon, Africa. This phenomenon occurred because much higher 

concentrations of uranium-235 were present at that time; the current uranium-235 concentration, about 

0.72%, will not sustain such natural reactions.) Technetium-99 is a key radionuclide in spent nuclear fuel, 

high-level radioactive wastes resulting from processing spent fuel, and radioactive wastes associated with 

operating nuclear reactors and fuel reprocessing plants. 

How Is It Used? Technetium is a very good corrosion inhibitor for steel, and protection can be 

achieved by adding only very small amounts during production. However, this use is limited by the 

Technetium
radioactive nature of technetium. Technetium-99m is commonly used in nuclear medicine as a 

radioactive tracer. In this application, the radionuclide is chemically attached to a drug chosen for its 

tendency to collect in specific organs of the body, and the solution is then injected into the patient. After 

a short time (its half-life is only 6 hours), an image is collected with a radiosensitive detector for analysis. 

This technique is very useful in identifying cancer metastases in locations distant from primary tumors. 

What’s in the Environment? Technetium is not a naturally occurring element. Technetium-99 is 

present in soil due to fallout from past atmospheric nuclear weapons tests. Estimated concentrations in 

surface soil are very low, on the order of 0.0001 picocuries per gram (pCi/g), due to its 

low specific activity. Technetium-99 is very mobile in the environment, especially under 

aerobic conditions (i.e., where oxygen is present). From the surface it can move rapidly 

downward with percolating water because most technetium compounds do not bind well 

to soil particles. The concentration associated with sandy soil particles is estimated at 0.1 

of that in interstitial water (in the pore spaces between the soil particles), although technetium binds more 

tightly to clay soils (with concentration ratios 10 times higher). For this reason, technetium-99 has been 

found in groundwater at several DOE sites. 

What Happens to It in the Body? Technetium pertechnetate (TcO4) is readily taken up from the 

intestines and lungs following ingestion or inhalation, with about 50 to 80% of the amount ingested being 

transferred to the bloodstream. After reaching the blood, about 4% of the technetium pertechnetate 

deposits in the thyroid where it is retained with a biological half-life of 0.5 days; the other two organs to 

which this isotope preferentially distributes are the stomach wall (10%) and liver (3%). The rest of what 

enters the blood is uniformly distributed throughout all other organs and tissues with a short residence 

time. Of the amount that reaches body tissues, half is excreted in urine and half is excreted in feces. For 

the technetium that is distributed to organs other than the thyroid, about 75% leaves the body with a 

biological half-life of 1.6 days, 20% clears with a half-life of 3.7 days, and 5% clears with a half-life of 

22 days. (This information is per simplified models that do not reflect intermediate redistribution.) 

What Are the Primary Health Effects? Technetium-99 is a health hazard only if it is taken into the 

body. It does not pose an external hazard because it decays by emitting a relatively low-energy beta 

particle with no gamma radiation. The main concern is cancer induction from the beta particles 

associated with its radioactive decay. 

Technetium can concentrate in several organs 

depending on its chemical form, so there is no 

primary organ of concern. This is one reason 

why the short-lived isotope technetium-99m has 

such wide usage in nuclear medicine as a 

diagnostic tool. The low energy of the beta 

particle, the lack of significant gamma or X-rays, 

and the rapid excretion of technetium-99 from the 

body limit the potential for health effects. 

What Is the Risk? Lifetime cancer mortality 

risk coefficients have been calculated for nearly 

all radionuclides, including technetium (see box 

at right). While the coefficients for ingestion are 

somewhat lower than for inhalation, ingestion is 

generally the most common means of entry into 

the body. Similar to other radionuclides, the risk 

coefficients for tap water are about 70% of those 

for dietary ingestion. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Recommended default 

absorption types were used for inhalation, and dietary 

values were used for ingestion. Risks are for lifetime 

cancer mortality per unit intake (pCi), averaged over 

all ages and both genders (10-12 is a trillionth). Other 

values, including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Technetium-97 7.6 × 10-13 2.3 × 10-13

Technetium-98 2.6 × 10-11 6.0 × 10-12

Technetium-99 1.3 × 10-11 2.3 × 10-12

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Symbol: Th

Atomic Number: 90 

(protons in nucleus)

Atomic Weight: 232 

(naturally occurring)

Radioactive Properties of Key Thorium Isotopes and Associated Radionuclides

Radiation Energy (MeV)

Isotope Half-Life 

Natural 

Abundance 

(%)

Specific 

Activity (Ci/g)

Decay 

Mode Alpha 

(α)

Beta

(β)

Gamma 

(γ)

Th-232 14 billion yr >99 0.00000011 α 4.0 0.012 0.0013 

Th-230 77,000 yr <<1 0.020 α 4.7 0.015 0.0016 

Th-229 7,300 yr <<1 0.22 α 4.9 0.12 0.096 

Ra-225 15 days 40,000 β - 0.11 0.014 

Ac-225 10 days 59,000 α 5.8 0.022 0.018 

Fr-221 4.8 min 180 million α 6.3 0.010 0.031 

At-217 0.032 sec 1.6 trillion α 7.1 < < 

Bi-213 46 min 20 million α, β 0.13 0.44 0.13 

Po-213 (98%) 0.0000042 sec 13,000 trillion α 8.4 - - 

Tl-209 (2%) 2.2 min 410 million β - 0.69 2.0 

Pb-209 3.3 hr 4.7 million β - 0.20 - 

Ci = curie, g = gram, and MeV = million electron volts; a ”<” means the radiation energy is 

less than 0.001 MeV, and a dash means the entry is not applicable. (See the companion fact 

sheet on Radioactive Properties, Internal Distribution, and Risk Coefficients for explanation of 

terms and interpretation of radiation energies.) Bismuth-213 decays by both emitting an alpha 

particle (2%) and a beta particle (98%). Certain properties of additional radionuclides are 

included here because they accompany the thorium decays. Values are to two significant figures. 

What Is It? Thorium is a radioactive element that occurs naturally in low concentrations 

(about 10 parts per million) in the earth’s crust. It is about three times as abundant as 

uranium and about as abundant as lead or molybdenum. Thorium in its pure form is a 

silvery-white heavy metal that is about as dense as lead. In nature, almost all thorium is 

thorium-232, although several additional isotopes can be present in small amounts. 

(Isotopes are different forms of an element that have the same number of protons in the 

nucleus but a different number of neutrons.) Thorium is a soft, ductile metal that is 

pyrophoric in powdered form. When heated in air, thorium turnings ignite and burn 

brilliantly with a white light. 

Of the 26 known isotopes of thorium, only 12 have half-lives greater than one second, and of these only 3 have half-lives 

sufficiently long to 

warrant a concern. 

These key isotopes 

decay very slowly 

by emitting an 

alpha particle. 

The half-lives of 

thorium-232 and 

thorium-230, the 

isotopes of most 

concern, are very 

long. Their low 

specific activity 

means these two 

isotopes are not 

highly radioactive. 

Both thorium-232 

and thorium-230 

are present in soil 

and ores in secular 

equilibrium with 

radium-228 and 

radium-226, 

respectively. The 

health risks for 

these two radium 

isotopes (shown in 

the Radium fact 

sheet) must be added to those shown here to estimate the total risk. Thorium-229 is not generally associated with nuclear 

fuel cycle activities previously performed by the Department of Energy (DOE), and hence is not a radionuclide of concern 

at DOE environmental management sites. The health risks associated with thorium-228, which has a half-life of 1.9 years, 

are commonly included with those for radium-228 because thorium-228 cannot persist for an extended period of time in the 

absence of radium-228. (See the companion fact sheet for Radium.)

Where Does It Come From? Thorium is widely distributed in small amounts in the earth’s crust. The chief commercial 

source is monazite sands in the United States (in North Carolina, South Carolina, Idaho, Colorado, Montana, and Florida) 

as well as in Brazil, India, Australia, and South Africa. The concentration of thorium oxide in monazite sands is about 3 to 

10%. Thorium is also found in the minerals thorite (thorium silicate) and thorianite (mixed thorium and uranium oxides). 

The isotope thorium-230, a decay product of uranium-238, is found in uranium deposits as well as in uranium mill tailings. 

How Is It Used? The principal use of thorium has been in the preparation of the Welsbach mantle for portable gas 

lanterns. These mantles contain thorium oxide with about 1% cerium oxide and other ingredients, and they glow with a 

dazzling light when heated in a gas flame. Thorium is an important alloying element in magnesium and is used to coat 

tungsten wire for components of electronic equipment. Thorium can also be added to ceramic items such as crucibles to 

make them more heat resistant, as well as to refractive glass to allow for smaller and more accurate camera lenses. In 

addition, thorium is used in welding rods and electric bulb filaments to improve product performance. 

Thorium
Thorium can also be used as a fuel in nuclear reactors. While thorium-232 itself is not fissile, it transforms into the fissile 

isotope uranium-233 upon absorption of a neutron. Although use of the thorium-232/uranium-233 fuel cycle has been 

demonstrated in pilot-scale studies, it has not been proven to be economically or technically viable for use in commercial 

nuclear power plants. 

What’s in the Environment? Thorium is naturally present in soil, rocks, surface water, groundwater, plants, and 

animals at low concentrations, on the order of ten parts per million. Higher levels are present in certain geological 

materials such as monzanite sands. Essentially all naturally occurring thorium is present as thorium-232. Thorium-230 is a 

radioactive decay product of uranium-238 and is found in low concentrations in uranium deposits and mill 

tailings. In its natural state, thorium occurs as an oxide (ThO2), phosphate (ThPO4), and silicate (ThSiO4). 

Thorium preferentially adheres very tightly to soil particles, with concentrations in sandy soil generally 

more than 3,000 times higher than in interstitial water (water in the pore spaces between soil particles); it is 

even less mobile in clay soils, with concentration ratios over 5,000. The concentration of thorium in plants 

is typically about 0.0042 (or 0.42%) of that in soil. Data have not indicated that it biomagnifies in terrestrial or aquatic 

food chains. Because of its low solubility, thorium is not generally a major groundwater contaminant at DOE sites. 

What Happens to It in the Body? Thorium can be taken into the body by eating food, drinking water, or breathing air. 

Most thorium that is inhaled or ingested in food and water is excreted within a few days, with only a small fraction being 

absorbed into the bloodstream. Gastrointestinal absorption from food or water is the principal source of internally 

deposited thorium in the general population. About 0.02 to 0.05% of the amount ingested is absorbed into the bloodstream 

through the intestines. Of the amount entering the blood, about 70% deposits in bone where it is retained with a biological 

half-life of about 22 years, 4% deposits in the liver where it is retained with a biological half-life of 700 days, and 16% is 

uniformly distributed to all other organs and tissues of the body where it is cleared with a biological half-life of 700 days. 

(per simplified models that do not reflect intermediate redistribution). Most of the remaining 10% is directly excreted. 

Thorium is predominantly deposited on the endosteal surfaces of mineral bone and only slowly redistributes throughout the 

bone volume. 

What Are the Primary Health Effects? Thorium is generally a health hazard only if it is taken into the body. External 

gamma exposure is not a major concern because thorium emits only a small amount of gamma radiation. Although 

thorium-229 has a much higher gamma component than either thorium-232 or thorium-230, thorium-229 comprises a very 

small fraction of natural thorium. (Note that if significant concentrations of radium occur along with thorium, which is 

common, the external gamma dose associated with the radium must also be addressed.) The major means of exposure to 

thorium are ingestion of food and water containing thorium and inhalation of thorium-contaminated dust. Ingestion is 

generally the main exposure concern, unless there is a nearby source of airborne dust containing thorium such as uranium 

mill tailings. Thorium is taken up in the body much more readily if inhaled rather than ingested (see table below), so both 

exposure routes can be important. The main health concern for environmental exposures is generally bone cancer. 

Most of the human data for thorium exposure comes from 

diagnostic studies. Colloidal thorium-232 dioxide 

(Thorotrast) was injected into patients as a radiographic 

contrast medium between 1928 and 1955. The 

epidemiological data from these studies show that the 

primary health effects of high doses of injected Thorotrast 

are blood disorders and liver tumors. Some evidence of 

increased incidence of lung, pancreatic, and hematopoietic 

cancers was found in workers occupationally exposed to 

thorium via inhalation. However, these workers were also 

exposed to several other toxic agents, so direct causation 

cannot be inferred. Few data are available regarding the 

health effects associated with low (e.g., environmental) 

levels of exposure from either inhalation or ingestion. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including thorium (see box at right). The risk 

coefficients for the isotopes of most concern (thorium-230 

and thorium-232) are similar. While the coefficients for 

thorium-229 are about five to eight times higher, this 

isotope is generally not of concern at DOE sites. The risk 

coefficients for ingestion, the most common type of 

exposure, are much lower than those for inhalation for all three isotopes. Similar to other radionuclides, the risk 

coefficients for tap water are nearly 80% of those for dietary ingestion. 

Radiological Risk Coefficients

This table provides selected risk coefficients for 

inhalation and ingestion. Recommended default 

absorption types were used for inhalation, and dietary 

values were used for ingestion. These values include the 

contributions from the short-lived thorium decay 

products. Risks are for lifetime cancer mortality per unit 

intake (picocurie, pCi), averaged over all ages and both 

genders (10-9 is a billionth, and 10-12 is a trillionth). 

Other values, including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Th-229 2.2 × 10-7 4.7 × 10-10 

Th-230 2.7 × 10-8 8.0 × 10-11

Th-232 4.1 × 10-8 9.1 × 10-11

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Radioactive Properties of Key Tin Isotopes 

and Associated Radionuclides 

Radiation Energy (MeV)

Isotope Half-Life 

 

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma

(γ) 

Sn-121m 55 yr 54 β, IT - 0.035 0.0049 

Sn-121 

 (78%) 27 hr 970,000 β - 0.11 - 

Sn-126 250,000 yr 0.029 β - 0.17 0.057 

Sb-126 12 d 85,000 β - 0.28 2.8 

IT = isomeric transition, Ci = curie, g = gram, and MeV = million electron 

volts; a dash means the entry is not applicable. (See the companion fact sheet 

on Radioactive Properties, Internal Distribution, and Risk Coefficients for an 

explanation of terms and interpretation of radiation energies.) Tin-121m 

decays by two means: emitting a beta particle (22%) and isomeric transition 

(78%). Certain properties of tin-121 and antimony-126 are included here 

because these radionuclides accompany the tin decays. Values are given to 

two significant figures. 

Symbol: Sn 

Atomic Number: 50

(protons in nucleus)

Atomic Weight: 119

(naturally occurring)

What Is It? Tin is a soft, silvery white metal with a highly crystalline structure 

that is malleable and ductile. When a bar of tin is bent, a crackling sound is 

emitted caused by the friction of the tin crystals. Tin exists in nature as ten stable 

isotopes. (Isotopes are different forms of an element that have the same number 

of protons in the nucleus, but a different number of neutrons.) These ten isotopes 

and their approximate abundances are tin-112 (1.0%), tin-114 (0.7%), tin-115 

(0.3%), tin-116 (15%), tin-117 (7.7%), tin-118 (24%), tin-119 (8.6%), tin-120 

(33%), tin-122 (4.6%), and tin-124 (5.8%). 

Thirteen major radioactive isotopes of tin exist of which only two – tin-121m and tin-126 – have half-lives long 

enough to warrant concern at 

Department of Energy (DOE) 

environmental management sites 

such as Hanford. Tin-121m has a 

half-life of 55 years, and tin-126 

has a half-life of 250,000 years; 

the half-lives of the other isotopes 

are less than one year. Both 

isotopes are fission products, with 

the fission yield of tin-126 being 

significantly larger than the yield 

of tin-121m. Thus, tin-126 is the 

more prevalent isotope at DOE 

sites. While tin-126 has a low 

specific activity that tends to limit 

its radioactive hazards, its shortlived daughter antimony-126 has a 

high gamma component, making 

external exposure to tin-126 a 

potential concern. 

Where Does It Come From? The principal ore of tin is the mineral cassiterite (tin oxide, SnO2) found either 

in rocks within often irregular veins or lodes, or in debris that has built up from the gradual wearing down of tinbearing rocks to form alluvial deposits found in river beds and valleys or on the ocean floor close to shore. Most 

of the world’s tin is produced by Malaysia, Brazil, Indonesia, Thailand, Bolivia, and Australia. The United 

States imports more than one-fifth of the average annual world production of tin. 

Tin-121m and tin-126 are produced by nuclear fission. When an atom of uranium-235 (or other fissile nuclide) 

fissions, it generally splits asymmetrically into two large fragments – fission products with mass numbers in the 

range of about 90 and 140 – and two or three neutrons. (The mass number is the sum of the number of protons 

and neutrons in the nucleus of the atom.) Tin-121m and tin-126 are two such fission products; the fission yield 

of tin-121m is very low (about 0.00003%), while the fission yield of tin-126 is about 0.06%. That is, much less 

than one atom of tin-121m and six atoms of tin-126 are produced per 10,000 fissions. These two tin isotopes are 

components of spent nuclear fuel, high-level radioactive wastes resulting from processing spent fuel, and 

radioactive wastes associated with the operation of nuclear reactors and fuel reprocessing plants. 

How Is It Used? Tin is used in a number of industrial processes throughout the world. Tinplate (sheet steel 

coated with a thin layer of tin) is the primary material used for food cans, and tin is also commonly used in 

bakeware and food storage containers. Tin is used to produce common alloys such as bronze (tin and copper) 

and solder (tin and lead). It is also used as an alloy with titanium in the aerospace industry and as an ingredient 

Tin
in some insecticides. Stannic sulfide, also known as mosaic gold, is used in powdered form for bronzing articles 

made of plaster of paris or wood. Dental amalgam contains about 13% tin together with silver and mercury, and 

mercury-free alternative dental filling materials with about twice the amount of tin are under development. Tin 

and tin-alloy coatings are widely used in the manufacture of bearings and in many kinds of machinery and 

fabricated parts, for both their anti-corrosion and lubricant properties. 

 

What’s in the Environment? Tin is present in the earth’s crust at a concentration of about 2.2 milligram per 

kilogram (mg/kg), and its concentration in seawater is about 0.01 micrograms (μg) per liter. 

Trace amounts of tin-121m and tin-126 are present in soil around the globe from radioactive 

fallout. These isotopes can also be present at certain nuclear facilities, such as reactors and 

spent fuel reprocessing plants. Tin is generally one of the less mobile radioactive metals in 

soil, although certain forms can move downward with percolating water some distance to 

underlying layers of soil depending on site conditions. Tin preferentially adheres quite well 

to soil, and the concentration associated with sandy soil particles is estimated to be about 

130 times higher than in interstitial water (the water in the pore spaces between the soil particles), with even 

higher concentration ratios in loam and clay. Thus, tin is generally not a major contaminant in groundwater at 

DOE sites. 

What Happens to It in the Body? Tin can be taken into the body by eating food, drinking water, or breathing 

air. Gastrointestinal absorption from food or water is the principal source of internally deposited tin in the 

general population. Gastrointestinal absorption is generally quite low, with only about 2% of the amount 

ingested being transferred to the bloodstream. Thirty-five percent of tin that reaches the blood is deposited in 

mineral bone, 15% is distributed throughout all other organs and tissues of the body, and the remaining 50% is 

excreted. Of the tin deposited in any organ or tissue, 20% is retained with a biological half-life of 4 days, 20% 

is retained with a biological half-life of 25 days, and 60% is retained with a biological half-life of 400 days. 

(This information is per simplified models that do not reflect intermediate redistribution.) 

 

What Are the Primary Health Effects? Tin poses an external as well as an internal health hazard. The 

strong gamma radiation associated with tin-126 makes external exposure to this isotope a concern. The main 

means of internal exposure are ingestion of food and water containing tin isotopes. While in the body, tin poses 

a health hazard from both the beta particles and gamma rays, and the main health concern is associated with the 

increased likelihood of inducing cancer in bone and other organs and tissues in which it may be deposited. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including tin (see box at right). While the 

coefficients for ingestion are lower than for inhalation, 

ingestion is generally the most common means of entry 

into the body. Similar to other radionuclides, the risk 

coefficients for tap water are about 70% of those for 

dietary ingestion. 

In addition to risks from internal exposures, there is an 

external gamma exposure risk associated with tin-126. To 

estimate a lifetime cancer mortality risk, if it is assumed 

that 10,000 people were continuously exposed to a thick 

layer of soil with an initial average concentration of 

1 pCi/g tin-126, then 6 of these 10,000 people would be 

predicted to incur a fatal cancer. (This is in comparison to 

the 2,000 people from this group predicted to die of cancer 

from all other causes per the U.S. average.) This risk is 

largely associated with the gamma ray emitted by its shortlived decay product antimony-126. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Maximum values are 

given for inhalation since no default absorption 

types were provided, and dietary values were used 

for ingestion. These values include the 

contributions from the short-lived tin decay 

products. Risks are for lifetime cancer mortality per 

unit intake (picocurie, pCi), averaged over all ages 

and both genders (10-9 is a billionth, and 10-12 is a 

trillionth). Other values, including for morbidity, 

are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Tin-121m 4.1 × 10-11 2.9 × 10-12

Tin-126 3.9 × 10-10 3.0 × 10-11

For more information, see the companion fact sheet 

on Radioactive Properties, Internal Distribution, 

and Risk Coefficients and accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Radioactive Properties of Tritium

Radiation Energy (MeV)

Isotope 

HalfLife 

(yr)

Natural 

Abundance 

(%)

Specific 

Activity 

(Ci/g)

Decay

Mode Alpha 

(α)

Beta

(β)

Gamma 

(γ)

H-3 12 <<1 9,800 β - 0.0057 - 

Ci = curie, g = gram, and MeV = million electron volts; a dash means the entry is not 

applicable. (See the companion fact sheet on Radioactive Properties, Internal 

Distribution, and Risk Coefficients for an explanation of terms and interpretation of 

radiation energies.) Values are given to two significant figures. 

Symbol: H (H-3)

Atomic Number: 1 

(protons in nucleus)

Atomic Weight: 1 

(naturally occurring H)

What Is It? Tritium is the only radioactive isotope of hydrogen. (An 

isotope is a different form of an element that has the same number of 

protons in the nucleus but a different number of neutrons.) The nucleus 

of a tritium atom consists of a proton and two neutrons. This contrasts 

with the nucleus of an ordinary hydrogen atom (which consists solely of a 

proton) and a deuterium atom (which consists of one proton and one 

neutron). Ordinary hydrogen comprises over 99.9% of all naturally 

occurring hydrogen. Deuterium comprises about 0.02%, and tritium comprises about a billionth of a 

billionth (10-16 percent) of natural hydrogen. 

The most common 

forms of tritium 

are tritium gas 

(HT) and tritium 

oxide, also called 

“tritiated water.” 

In tritiated water, 

a tritium atom 

replaces one of the 

hydrogen atoms so 

the chemical form 

is HTO rather than H2O. The chemical properties of tritium are essentially the same as those of ordinary 

hydrogen. It decays with a half-life of 12 years by emitting a beta particle to produce helium-3. Tritium 

has a relatively high specific activity and is generated by both natural and artificial processes. It is of 

concern at Department of Energy (DOE) sites that operated tritium production facilities, such as Hanford. 

Where Does It Come From? Tritium is naturally present as a very small percentage of ordinary 

hydrogen in water, both liquid and vapor. This tritium is produced as a result of the interaction of cosmic 

radiation with gases in the upper atmosphere, and the natural steady-state global inventory is about 

7.3 kilograms (kg). (About five times this amount remains from past atmospheric nuclear weapons tests.) 

After being produced in the atmosphere, it is readily incorporated into water and falls to earth as rain, thus 

entering the natural hydrological cycle. Tritium is also produced as a fission product in nuclear weapons 

tests and in nuclear power reactors, with a yield of about 0.01%. That is, about one atom of tritium is 

produced per 10,000 fissions. Each year a large commercial nuclear power reactor produces about 

20,000 curies (2 grams) of tritium, which is generally incorporated in the nuclear fuel and cladding. 

Because little tritium is naturally present, it must be produced artificially for use on a practical scale. 

Tritium can be made in production nuclear reactors, i.e., reactors designed to optimize the generation of 

tritium and special nuclear materials such as plutonium-239. Tritium is produced by neutron absorption 

of a lithium-6 atom. The lithium-6 atom, with three protons and three neutrons, and the absorbed neutron 

combine to form a lithium-7 atom with three protons and four neutrons, which instantaneously splits to 

form an atom of tritium (one proton and two neutrons) and an atom of helium-4 (two protons and two 

neutrons). The United States has recovered an estimated 225 kg of tritium, of which 150 kg has decayed 

into helium-3, leaving a current inventory of approximately 75 kg. While tritium can also be produced in 

accelerators by bombarding helium-3 with neutrons, this approach has not been proven on a large scale. 

How Is It Used? Tritium is used as a component in nuclear weapons to boost the yield of both fission 

and thermonuclear (or fusion) warheads. Tritium is also used as a tracer in biological and environmental 

Tritium (Hydrogen-3) 
Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and dietary ingestion of tritiated water. 

Risks are for lifetime cancer mortality per unit intake 

(pCi), averaged over all ages and both genders (10-12

is a trillionth). Other values, including for morbidity, 

are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Tritium (H-3) 3.9 × 10-14 4.4 × 10-14

For more information, see the companion fact sheet 

on Radioactive Properties, Internal Distribution, and 

Risk Coefficients and the accompanying Table 1. 

studies, and as an agent in luminous paints such as those used to make building exit signs, airport runway 

lights, and watch dials.

What’s in the Environment? Tritium is present in water (liquid and vapor) as a result of natural 

processes in the atmosphere, as well as from fallout from past atmospheric nuclear weapons tests and the 

operation of nuclear reactors and fuel reprocessing plants. The form of most concern, 

tritium oxide (HTO), is generally indistinguishable from normal water and can move 

rapidly through the environment in the same manner as water. Tritium is naturally 

present in surface waters at about 10 to 30 picocuries per liter (pCi/L). The maximum 

contaminant level developed by the Environmental Protection Agency for tritium in 

drinking water supplies is 20,000 pCi/L or 0.02 microcuries per liter (a picocurie is a 

millionth of a microcurie). Higher concentrations can be present in water at facilities 

that produce and utilize tritium, including certain DOE sites. 

What Happens to It in the Body? Tritium can be taken into the body by drinking water, eating food, 

or breathing air. It can also be taken in through the skin. Nearly all (up to 99%) inhaled tritium oxide can 

be taken into the body from the lungs, and circulating blood then distributes it to all tissues. Ingested 

tritium oxide is also almost completely absorbed, moving quickly from the gastrointestinal tract to the 

bloodstream. Within minutes it is found in varying concentrations in body fluids, organs, and other 

tissues. Skin absorption of airborne tritium oxide can also be a significant route of uptake, especially for 

exposure to high concentrations of tritiated water vapor, as could occur under conditions of high humidity 

during hot weather, because of the normal movement of water through the skin. For someone immersed 

in a cloud of airborne tritium oxide (HTO), the uptake by absorption through the skin would be about half 

that associated with inhalation. No matter how it is taken into the body, tritium is uniformly distributed 

through all biological fluids within one to two hours. Tritium is eliminated from the body with a 

biological half-life of 10 days, the same as for water. During the time it is in the body, a small fraction of 

the tritium is incorporated into easily exchanged hydrogen sites in organic molecules. 

What Is the Primary Health Effect? Tritium poses a health hazard only if it is taken into the body, 

because tritium decays by emitting a low-energy beta particle with no gamma radiation. This beta 

particle cannot penetrate deeply into tissue or travel far in air. The most likely form of uptake is as 

tritium oxide (or tritiated water), as the uptake of tritium gas is typically very low (less than 1%). 

Tritiated water behaves the same as ordinary water, both in the environment and in the human body. 

Hence, a significant fraction of the inhaled and 

ingested tritium is directly absorbed into the 

bloodstream. The health hazard of tritium is 

associated with cell damage caused by the 

ionizing radiation that results from radioactive 

decay, with the potential for subsequent cancer 

induction. 

What Is the Risk? Lifetime cancer mortality 

risk coefficients have been calculated for nearly 

all radionuclides, including tritium (see box at 

right). The values given here are for tritiated 

water; additional values are available, including 

for inhalation and ingestion of organically bound 

tritium and inhalation of tritium on particulates. 

As for other nuclides, the risk coefficient for tap 

water is about 80% of that for dietary ingestion. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: U

Atomic Number: 92 

(protons in nucleus)

Atomic Weight: 238

(naturally occurring)

Radioactive Properties of Key Uranium Isotopes and Associated Radionuclides

Radiation Energy (MeV)

Isotope Half-Life 

Natural 

Abundance 

(%)

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

U-232 72 yr 0 22 α 5.3 0.017 0.0022 

U-233 160,000 yr 0 0.0098 α 4.8 0.0061 0.0013 

U-234 240,000 yr 0.0055 0.0063 α 4.8 0.013 0.0017 

U-235 700 million yr 0.72 0.0000022 α 4.4 0.049 0.16 

Th-231 26 hr 540,000 β - 0.17 0.026 

U-236 23 million yr 0 0.000065 α 4.5 0.011 0.0016 

U-238 4.5 billion yr >99 0.00000034 α 4.2 0.010 0.0014 

Th-234 24 days 23,000 β - 0.060 0.0093 

Pa-234m 1.2 min 690 million β - 0.82 0.012 

Ci = curie, g = gram, and MeV = million electron volts; a dash means the entry is not applicable. (See 

the companion fact sheet on Radioactive Properties, Internal Distribution, and Risk Coefficients for an 

explanation of terms and interpretation of radiation energies.) Properties of thorium-231, thorium-234, 

and protactinium-234m are included here because these radionuclides accompany the uranium decays. 

Values are given to two significant figures. 

What Is It? Uranium is a radioactive element that occurs naturally in low concentrations (a 

few parts per million, ppm) in soil, rock, surface water, and groundwater. It is the heaviest 

naturally occurring element, with an atomic number of 92. Uranium in its pure form is a 

silver-colored heavy metal that is nearly twice as dense as lead. In nature, uranium exists as 

several isotopes: primarily uranium-238, uranium-235, and a very small amount of 

uranium-234. (Isotopes are different forms of an element that have the same number of 

protons in the nucleus but a different number of neutrons.) In a typical sample of natural 

uranium, almost all the mass (99.27%) consists of atoms of uranium-238. Less than 1% 

(about 0.72%) of the mass consists of atoms of uranium-235, and a very small amount (0.0055% by mass) is uranium-234. 

Uranium decays 

very slowly by 

emitting an alpha 

particle. The halflife of uranium238 is 4.5 billion 

years, which 

means it is not 

very radioactive 

as indicated by its 

low specific activity. The very 

long half-lives of 

these isotopes are 

the reason why 

uranium still 

exists on earth. 

Three additional 

isotopes 

(uranium-232, 

uranium-233, and 

uranium-236) are 

not naturally 

present but can be produced by nuclear transformations. These three isotopes also decay by emitting an alpha particle. 

Where Does It Come From? While small amounts of natural uranium are found almost everywhere in soil, rock, and 

water, uranium ores are found in just a few places – usually in hard rock or sandstone, in deposits normally covered with 

earth and vegetation. Uranium has been mined in the southwest United States, Canada, Australia, parts of Europe, the 

former Soviet Union, Namibia, South Africa, Niger, and elsewhere. It is a contaminant at many U.S. Department of Energy 

sites (including Hanford) and other facilities that used natural uranium, including mining, milling, and production facilities. 

How Is It Used? For many years, uranium was used to color ceramic glazes, producing colors that ranged from orangered to lemon yellow. It was also used for tinting in early photography. The radioactive properties of uranium were not 

recognized until 1896, and its potential for use as an energy source was not realized until the middle of the 20th century. In 

nuclear reactors, uranium serves as both a source of neutrons (via the fission process) and a target material for producing 

plutonium. (Plutonium-239 is produced when uranium-238 absorbs a neutron.) Today, its primary use is as fuel in nuclear 

power reactors to generate electricity. Uranium is also used in small nuclear reactors to produce isotopes for medical and 

industrial purposes around the world. Natural uranium must be enriched in the isotope uranium-235 for use as a nuclear 

fuel in light-water reactors, and this enrichment has generally been achieved by gaseous diffusion techniques. Highly 

enriched uranium is a primary component of certain nuclear weapons. A byproduct of the enrichment process is depleted 

uranium, i.e., uranium depleted in the isotope 235. (See the companion fact sheet for Depleted Uranium.) 

 

What’s in the Environment? Uranium is naturally present in all environmental media at very low concentrations (a few 

parts per million). Higher levels are present in certain areas, including those with natural uranium ores such as in the 

southwestern United States. In its natural state, uranium occurs as an oxide ore, U3O8. Additional compounds that may be 

present include other oxides (UO2, UO3) as well as fluorides, carbides or carbonates, silicates, vanadates, and phosphates. 

In addition to the three naturally occurring isotopes, uranium-232, uranium-233, and uranium-236 are present at Hanford. 

At that site, uranium-233 was produced in targets and disposed of in the 300 Area; uranium-236 measurements in 

Uranium
groundwater there have been used to distinguish the presence of natural uranium from uranium associated with reprocessed 

nuclear fuel. The environmental transport of uranium is strongly influenced by its chemical form. It is 

generally one of the more mobile radioactive metals and can move down through soil with percolating 

water to underlying groundwater. Uranium preferentially adheres to soil particles, with a soil 

concentration typically about 35 times higher than that in the interstitial water (the water between the soil 

particles); concentration ratios are usually much higher for clay soils (e.g., 1,600). Uranium can 

bioconcentrate in certain food crops and in terrestrial and aquatic organisms. However, data do not 

indicate that it biomagnifies in terrestrial or aquatic food chains. The U.S. Environmental Protection 

Agency (EPA) established a maximum contaminant level (MCL) for uranium in drinking water of 0.030 milligram per liter 

(mg/L). This equates to about 27 picocuries (pCi) per liter considering the ratio of isotopes typically present in drinking 

water sources. 

What Happens to It in the Body? Uranium can be taken into the body by eating food, drinking water, or breathing air. 

Gastrointestinal absorption from food or water is the main source of internally deposited uranium in the general population. 

After ingestion, most uranium is excreted within a few days and never enters the bloodstream. The small fraction (0.2 to 

5%) that is absorbed into the bloodstream is deposited preferentially in bone (about 22%) and kidneys (about 12%), with 

the rest being distributed throughout the body (12%) and excreted. Most of what goes to the kidneys leaves within a few 

days (in urine), while that deposited in bone can remain for many years. After inhalation, generally only a small fraction 

penetrates to the lung’s alveolar region, where it can remain for years and from which it can also enter the bloodstream. 

What Are the Primary Health Effects? Uranium is a health hazard only if it is taken into the body. External exposure 

is generally not a major concern because uranium emits only a small amount of low-energy gamma radiation. While 

uranium-235 has a much higher gamma component than either uranium-234 or uranium-238, uranium-235 only comprises 

about 2% of the total activity of natural uranium. The primary means of exposure are ingestion of food and water 

containing uranium isotopes and inhalation of uranium-contaminated dust. Ingestion is usually the exposure of concern 

unless there is a nearby source of airborne dust, such as a uranium mine or mill. Because uranium is absorbed much more 

readily if inhaled rather than ingested, both exposure routes can be important. The major health concern is kidney damage 

caused by the chemical toxicity of soluble uranium compounds. That effect can be reversible depending on the level of 

exposure. (Uranium has also been implicated in reproductive effects in laboratory animals and developmental effects in 

young animals, but it is not known if these problems exist for humans.) A second concern is for uranium deposited in bone, 

which can lead to bone cancer as a result of the ionizing radiation associated with its radioactive decay products. 

What Is the Risk? Lifetime cancer mortality risk 

coefficients have been calculated for nearly all 

radionuclides, including uranium (see box at right). 

Although ingestion is generally the common means of 

entry, these risk coefficients are much lower than those 

for inhalation so both exposure routes need to be 

considered. Similar to other radionuclides, the risk 

coefficients for tap water are about 75% of those for 

dietary ingestion. On an activity (curie) basis, the risk 

coefficients are essentially the same for all uranium 

isotopes (although the factor for ingesting uranium-232 is 

somewhat higher), so the risk is essentially independent of 

the ratio of various isotopes in a compound. For this 

reason, the risk from exposure to depleted uranium is 

essentially the same as for enriched uranium on an 

activity basis. Uranium-235 also poses an external gamma 

exposure risk. To estimate a lifetime cancer mortality risk, 

if it is assumed that 100,000 people were continuously 

exposed to a thick layer of soil with an initial 

concentration of 1 pCi/g uranium-235, then 3 of those 

100,000 people would be predicted to incur a fatal cancer. 

(This is in comparison to about 20,000 people from the 

group predicted to die of cancer from all other causes per 

the U.S. average.) Uranium can also kidney damage due 

to its chemical toxicity. The toxicity value used to 

estimate the potential for non-cancer effects following 

ingestion is a reference dose (RfD), which is an estimate 

of the highest dose that can be taken in every day over a 

lifetime without causing an adverse health effect. In addition to the on-line RfD shown above, EPA more recently derived 

a value of 0.0006 mg/kg-day to support the drinking water MCL. These values were developed by analyzing the biological 

effects of test animals given relatively large amounts of uranium, then adjusting and normalizing the results to a mg/kg-day 

basis for humans. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for inhalation 

and ingestion. Recommended default absorption types were 

used for inhalation, and dietary values were used for 

ingestion. These values include contributions from shortlived uranium decay products. Risks are for lifetime cancer 

mortality per unit intake (pCi), averaged over all ages and 

both genders (10-9 is a billionth, and 10-12 is a trillionth). 

Other values, including for morbidity, are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Uranium-232 1.8 × 10-8 2.7 × 10-10

Uranium-233 1.1 × 10-8 6.3 × 10-11

Uranium-234 1.1 × 10-8 6.1 × 10-11

Uranium-235 9.5 × 10-9 6.2 × 10-11

Uranium-236 9.9 × 10-9 5.8 × 10-11

Uranium-238 8.8 × 10-9 7.5 × 10-11

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 

Chemical Toxicity Value 

Non-Cancer Effect: Oral RfD (soluble salts)

0.003 mg/kg-day
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Specific Activity of Enriched and Depleted Uranium 

(For additional information, see the Uranium fact sheet.) 

The specific activity (SA) of a uranium compound depends on its 

isotopic composition. The SA of natural uranium (containing 

0.72% uranium-235) is 6.77×10-7 curies per gram (Ci/g). The SA 

for other mixtures of uranium-238, uranium-235, and uranium-234 

can be estimated using the approach developed by the U.S. Nuclear 

Regulatory Commission (see 10 CFR 20): 

For depleted uranium (uranium-235 containing less than 0.72%): 

SA = 3.6 × 10-7 Ci/g 

For enriched uranium (uranium-235 containing more than 0.72%): 

SA = {0.4 + 0.38(enrichment) + 0.0034(enrichment)2}×10-6 Ci/g, 

where enrichment is the percent uranium-235. 

Thus, the SA of depleted uranium is approximately half that of 

natural uranium. (To express SA in standard international units, 

multiply the value in Ci/g by 3.7 × 1010 becquerels [Bq]/Ci.)

What Is Depleted Uranium? Depleted uranium is created as a byproduct during the processing of natural 

uranium to make it suitable for use as fuel in nuclear power plants or as a component in nuclear weapons. In 

natural uranium, 99.27% of the mass consists of atoms of uranium-238. About 0.72% of the mass consists of 

atoms of uranium-235, and a very small amount (0.0055% by mass) is uranium-234. Although uranium-235 

is the rarer of the two major uranium isotopes, it is the one that most readily undergoes nuclear fission and is 

thus the most useful for common nuclear applications. For most of these applications, the proportion of the 

uranium-235 isotope found in natural uranium must be increased through a process called enrichment. 

The uranium enrichment process results 

in the production of “enriched” uranium 

(containing >0.72% uranium-235) and 

“depleted” uranium (containing <0.72% 

uranium-235). The uranium enrichment 

process also removes much of the 

uranium-234 from the depleted uranium. 

Most depleted uranium in the United 

States contains between 0.2 and 0.4% 

uranium-235, with the remainder being 

uranium-238 and a slight amount of 

uranium-234. Depleted uranium is less 

radioactive than natural uranium because 

some of the uranium-235 and most of the 

uranium-234 have been removed. 

Depleted uranium will not undergo a 

nuclear chain reaction. 

How Is It Produced? In the United States, uranium is enriched by the gaseous diffusion process in which 

the compound uranium hexafluoride (UF6) is heated and converted from a solid to a gas. The UF6 gas is 

forced through a long series of compressors and converters with porous barriers. Because uranium-235 has a 

slightly lighter isotopic mass than uranium-238, UF6 molecules made with uranium-235 diffuse through the 

barriers at a slightly higher rate than do the molecules containing uranium-238 (the uranium-234 molecules 

diffuse through the barriers at a rate even higher than that of the uranium-235 molecules). 

At the end of the process there are two UF6 streams, with one having a higher concentration of uranium-235 

than the other. The stream with the higher uranium-235 concentration is referred to as enriched UF6, while 

the stream that is reduced in its concentration of uranium-235 is referred to as depleted UF6. The depleted 

UF6 can be converted to other chemical forms, such as depleted uranium oxide or depleted uranium metal. 

Other methods can be used to enrich uranium, and they too produce depleted uranium as a byproduct. The 

most common enrichment process used outside of the United States is gas centrifuge enrichment. Laserbased enrichment processes have also been investigated but not commercially developed. 

How Is It Used? Although the vast majority of depleted uranium is stored as a byproduct of the enrichment 

process at the gaseous diffusion plant sites, several current and potential uses exist. Because of its high 

density, depleted uranium is currently used for radiation shielding. Depleted uranium metal was previously 

used on large commercial aircraft as counterweights in the wings. Military applications of depleted uranium 

include use as tank armor, armor-piercing projectiles (antitank weapons), and counterweights in missiles and 

high-performance aircraft. In a potential future use, depleted uranium could be mixed with highly enriched 

uranium from retired nuclear weapons to produce nuclear reactor fuel. This process is called blending, and 

to date only natural or slightly enriched uranium has been considered for this application. Limited amounts 

of depleted uranium can also be used for the fabrication of mixed uranium and plutonium oxide (MOX) fuel, 

which is currently being evaluated for use in commercial nuclear power plants. 

Depleted Uranium 

(This complements the general fact sheet for uranium.)
What Are Depleted Uranium Munitions? Research, tests, and evaluations were conducted during the 

1970s to develop improved armor-penetrating munitions capable of defeating a heavily armored target. 

High-density materials such as tungsten and depleted uranium were considered candidates for this type of 

munitions; depleted uranium was ultimately selected because of its very high density, availability, 

noncompetitive uses, and pyrophoricity (spontaneous combustion upon exposure to air). 

The two main antitank weapons made of depleted uranium are the 

120-millimeter (mm) cannon shell used by the Army's M1 Abrams 

tank, and the 30-mm munitions used by A-10 antitank aircraft. 

Depleted uranium weapons are also used in Britain's Challenger 

tank, and the U.S. and British Navy Phalanx gun systems. The M1 

Abrams tank fires a 120-mm shell that contains a kinetic energy 

penetrator, a solid rod with a pointed tip. The penetrator is made 

from about 5 kilograms (kg) (11 pounds [lb]) of depleted uranium 

metal alloyed with less than 1% titanium. The shell is a sabot 

round, meaning that the penetrator is contained within a carrier as 

it is fired from the barrel. The carrier separates from the penetrator 

after exiting the gun barrel. In addition, many M1 Abrams tanks 

incorporate depleted uranium metal as part of the external armor. 

Depleted uranium rounds are also fired by a 30-mm, seven-barrel 

gatling gun mounted in the nose of the A-10 Thunderbolt aircraft, 

the only U.S. military plane that employs depleted uranium rounds. Depleted uranium is the primary 

munition for the A-10 Thunderbolt for combat. Each 30-mm depleted uranium projectile contains 

approximately 0.3 kg (0.66 lb) of extruded depleted uranium metal alloyed with 0.75% titanium. The 

projectile is encased in a 0.8-mm-thick aluminum shell as the final depleted uranium round. 

Depleted uranium penetrators have a “sharpening effect” upon impact that allows greater penetration through 

armor. Weapon testing shows that when a depleted uranium round penetrates an armored vehicle, it may 

pass completely through the vehicle or ricochet around and fragment inside the vehicle. Metal fragments 

from the penetrator and the vehicle’s hull can scatter inside the vehicle, killing and injuring personnel, 

destroying equipment, and causing secondary explosions and fires. As much as 70% of a depleted uranium 

penetrator can be aerosolized when it strikes a tank. Aerosols containing uranium oxides may contaminate 

the area downwind. Uranium metal and oxide fragments may also contaminate the soil around the struck 

vehicle. Tests of depleted uranium penetrators striking depleted uranium armored vehicles have shown that 

most of the contamination will occur within 5 to 7 meters (16 to 23 feet) of the vehicle. 

For Whom Is Exposure a Concern? Soldiers can be exposed to depleted uranium when intact rounds 

are in storage or uploaded in armored vehicles, or on the battlefield following the use of depleted uranium 

munitions. Studies have shown that exposure to intact depleted uranium weapons systems, both munitions 

and armor, pose very little risk from external radiation. Uranium and its decay products are primarily alpha 

emitters, emitting only very low levels of gamma radiation. However, individuals involved in handling or 

processing depleted uranium generally wear heavy gloves to protect against the high-energy beta particle 

from protactinium-234m, which is only a concern for skin exposure. 

On the battlefield, exposure to depleted uranium can occur several ways. After impact with an armored 

vehicle, depleted uranium in the form of oxide and metal fragments will contaminate the struck vehicle and 

the surrounding area, especially if an explosion or fire occurs. Depleted uranium oxide can be aerosolized, 

presenting an inhalation hazard to soldiers in the vicinity or downwind of the struck vehicle. In addition, 

depleted uranium dust can be resuspended in the air from soil contamination around the vehicle. Soil 

contamination also can be ingested inadvertently from hand to mouth contact. In another example of 

battlefield exposure, at least 36 soldiers in Operation Desert Storm were reported to have wounds involving 

imbedded depleted uranium fragments. (Many of these fragments were not removed because the risk of 

120-mm antitank kinetic energy 

rounds fired by the M1 Abrams tank 
surgery was judged to be too great.) Following combat, maintenance and recovery personnel can be exposed 

to equipment contaminated with depleted uranium. Dust containing depleted uranium can be resuspended 

during recovery operations, posing an inhalation hazard. Spent depleted uranium rounds and penetrators that 

were fired but missed their targets may be scattered on the battlefield and can thus potentially lead to 

additional exposures. 

What Happens to It in the Body? After it is ingested, most uranium is excreted from the body within a 

few days and never enters the bloodstream. The small fraction that is absorbed into the bloodstream (0.2 to 

5%) is deposited preferentially in bone and kidneys. Most of what goes to the kidneys leaves within a few 

days (in urine), while that deposited in bone can remain for many years. After inhalation, generally only a 

small fraction penetrates to the lung’s alveolar region, where it can remain for years and can also enter the 

bloodstream. 

What Are the Primary Health Effects? Uranium is a health hazard only if it is taken into the body. 

External gamma exposure is generally not a major concern because uranium emits only a small amount of 

low-energy gamma radiation, and beta exposure is only of concern for direct handling operations (due to 

potential skin effects). The main means of exposure are ingestion of food and water containing uranium 

isotopes and inhalation of uranium-contaminated dust. Ingestion is usually the exposure route of concern 

unless there is a nearby source of airborne dust. Because uranium is taken up in the body much more readily 

if inhaled than if ingested, both exposure routes can be important. The major health concern is kidney 

damage caused by the chemical toxicity of soluble uranium compounds; these effects can be reversible 

depending on the level of exposure. Uranium is not considered a chemical carcinogen. A second concern is 

for uranium deposited in bone, which can lead to bone cancer as a result of the ionizing radiation associated 

with the radioactive decay products. Uranium has caused reproductive problems in laboratory animals and 

developmental problems in young animals, but it is not known if these problems exist for humans. 

What Are the Risks? Lifetime cancer mortality risk coefficients have been calculated for nearly all 

radionuclides, including uranium (see the companion fact sheet for Uranium). On an activity (curie) basis, 

these risk coefficients are essentially the same for all three natural uranium isotopes. Hence, the risk is 

essentially independent of the ratio of the various isotopes in a uranium compound. For this reason, the risk 

of a fatal cancer from exposure to depleted uranium is essentially the same as for enriched uranium on an 

activity basis. However, because of the difference in specific activity, equal activities of enriched and 

depleted uranium have significantly different masses. 

The decay of uranium isotopes in depleted uranium results in the generation of radioactive decay products. 

For example, thorium-234 and protactinium-234m (the “m” meaning metastable) are produced from the 

decay of uranium-238. The recycling of irradiated uranium may also result in some radioactive impurities in 

depleted uranium. Some of these decay products and impurities could potentially also exhibit chemical 

toxicity. The decay products and impurities contribute to the radiation emitted by depleted uranium; 

however, the concentrations are generally too small to be of significance when assessing the chemical health 

effects from exposures to depleted uranium. 

The extent of chemical damage from exposure to a depleted uranium compound depends on its solubility and 

the route of exposure. In most assessments, only inhalation and ingestion are considered because although 

dermal absorption of some soluble compounds (e.g., uranyl nitrate) is possible, these exposures generally are 

not significant in association with industrial emissions or environmental exposures. When soluble or 

moderately soluble compounds such as uranyl fluoride (UO2F2) or uranium tetrafluoride (UF4) are inhaled or 

ingested, some of the uranium enters the bloodstream and reaches the kidney and other internal organs; thus, 

chemical toxicity is of primary importance. 

When insoluble compounds such as UO2 and U3O8 are inhaled, the uranium is generally deposited in the 

lungs and can remain there for long periods of time (months or years). The main concern from exposure to 

these insoluble compounds is increased cancer risk from the internal exposure to radioactivity. Ingested 

insoluble compounds are poorly absorbed from the gastrointestinal tract, and so generally have low toxicity. 
Available health effect criteria for acute and chronic uranium exposures are provided in Tables 1 and 2, 

respectively, derived from animal studies. Only limited data exist for human exposures.

TABLE 1. Reference Levels for Chemical Effects of Acute Uranium Exposures

Health Endpoint Estimated Intake Source Air Concentration 

of Various Compounds a

Potential irreversible 

adverse effects 

30 mg soluble uranium NRC (1994) 

Fisher et al. (1994) 

UO2F2; UF4: 26 mg/m3

U3O8: 590 mg/m3

UO2: 570 mg/m3

Potential adverse 

effects 

10 mg soluble uranium McGuire (1991) UO2F2; UF4: 9 mg/m3

U3O8: 200 mg/m3

UO2: 190 mg/m3

a Assumes 1-hour exposure at an inhalation rate of 1.5 m3

 per hour; assuming a higher inhalation rate for a shorter 

exposure duration may be appropriate (ICRP 1994) and will result in lower air concentrations. A factor of 25 

decrease in systemic absorption of insoluble compounds (i.e., U3O8 and UO2) was also assumed. These 

concentrations may not be protective against the adverse radiological effects of insoluble uranium compounds.

TABLE 2. Reference Levels for Chemical Effects of Chronic Uranium Exposures

Route Reference Levela Source Notes 

Ingestion Soluble compounds

(1) 0.003 mg/kg/d 

(2) 0.002 mg/kg/d 

Insoluble compounds

(1) 0.003 mg/kg/d 

(2) Compound-specific based 

on absorption data 

EPA (2004); Maynard 

and Hodge (1949) 

ATSDR, 1999; 

Gilman et al., 1998 

EPA (2004); Maynard 

and Hodge (1949) 

ATSDR, 1999; 

Gilman et al., 1998 

Based on a lowest observed adverse 

effect level (LOAEL) of 3 mg/kg-day 

and an uncertainty factor of 1,000. 

Based on an intermediate-duration 

(91 days) oral exposure of rabbits to a 

soluble uranium compound. 

Same as for soluble compounds; no 

correction for decreased absorption. 

Inhalation Soluble compounds 

(1) 0.011 mg/m3

 

(2) 0.0012 mg/m3

(3) 0.0004 or 

 0.0003 mg/m3

Insoluble compounds

(1) 0.011 mg/m3

 

(2) 0.28 mg/m3

(3) 0.008 mg/m3

(4) 0.006 mg/m3

Derived 

Derived 

ATSDR (1999) 

Derived 

Derived 

ATSDR (1999) 

Derived 

Based on route-to-route extrapolation of 

ingestion value (EPA 1996). 

Based on adjustments to occupational 

exposure limit (EPA 1992). 

Based on intermediate duration 

(5 weeks) and chronic (1 year) 

inhalation exposures of dogs to soluble 

uranium compounds. 

Same as for soluble compounds; no 

correction for decreased absorption. 

Assume absorption is decreased by a 

factor of 25 (ICRP 1979). 

Based on intermediate duration 

(5 weeks) inhalation exposures of dogs 

to insoluble uranium dioxide. 

Based on adjustments made to 

occupational exposure limit (EPA 1992). 

a Soluble compounds include uranyl fluoride (UO2F2), uranyl nitrate hexahydrate (UO2[NO3]2@ 6H2O), and uranium 

tetrachloride (UCl4). Some insoluble compounds are uranium metal, uranium dioxide (UO2), uranium trioxide 

(UO3), and triuranium octaoxide (U3O8). Uranium hexafluoride (UF6) decomposes to UO2F2 and hydrogen fluoride 

(HF) in the presence of moisture. Uranium tetrafluoride (UF4) is very slightly soluble. Source: ATSDR (1999). 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Zr

Atomic Number: 40

(protons in nucleus)

Atomic Weight: 91

(naturally occurring)

Radioactive Properties of the Key Zirconium Isotope

and an Associated Radionuclide 

Radiation Energy (MeV)

Isotope Half-Life 

(yr)

Specific

Activity

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ) 

Zr-93 1.5 million 0.0025 β - 0.020 - 

Nb-93m 14 290 IT - 0.028 0.0019 

IT =isomeric transition, Ci = curie, g = gram, and MeV = million electron 

volts; a dash means the entry is not applicable. (See the companion fact sheet 

on Radioactive Properties, Internal Distribution, and Risk Coefficients for an 

explanation of terms and interpretation of radiation energies.) Certain 

properties of niobium-93m are included here because this radionuclide can 

accompany zirconium-93 decays. Values are given to two significant figures. 

What Is It? Zirconium is a lustrous, grayish-white, corrosion-resistant 

metal. When it exists in a finely divided form, zirconium can 

spontaneously ignite in air, especially at high temperatures; the solid metal 

is much more difficult to ignite. Zirconium occurs in nature as five stable 

isotopes. (Isotopes are different forms of an element that have the same 

number of protons in the nucleus but a different number of neutrons.) 

Zirconium-90 is the most prevalent form, comprising slightly more than 

half of natural zirconium. The other four stable isotopes and their relative abundances are zirconium-91 

(11%), zirconium-92 (17%), zirconium-94 (17%), and zirconium-96 (2.8%). Zirconium occurs widely in 

the earth’s crust, but not in concentrated deposits. 

Of the six major radioactive zirconium isotopes, only one – zirconium-93 – has a half-life long enough to 

warrant potential concern at Department of Energy (DOE) environmental management sites such as 

Hanford. The half-lives of all other isotopes are less than three months. Zirconium-93 decays by 

emitting a beta particle with 

a half-life of 1.5 million 

years to niobium-93m (the 

“m” means metastable), 

which in turn decays by 

isomeric transition with a 

half-life of 14 years. 

Zirconium-93 is present in 

spent nuclear fuel and the 

wastes resulting from 

reprocessing this fuel. The 

low-specific activity and 

low energy of its radiations 

limit the radioactive 

hazards of this isotope. 

Where Does It Come From? Zirconium is naturally present in a number of minerals, but it does not 

generally occur in concentrated deposits. The mineral zircon (zirconium orthosilicate), which is found in 

alluvial deposits in streambeds, ocean beaches, or old lake beds, is the only commercial source of 

zirconium. Commercial grade zirconium contains from 1 to 3% hafnium. The leading producers of 

zirconium are Australia and South Africa. Other producers include Ukraine, the United States, India, 

Brazil, and China. Annual worldwide production of zirconium is currently about 900,000 metric tons 

(MT) of which about 60 MT is produced domestically. 

Zirconium-93 is produced by neutron activation of components in nuclear reactors and by nuclear fission. 

When a fissile nuclide such as an atom of uranium-235 fissions, it generally splits asymmetrically into 

two large fragments – fission products with mass numbers in the range of about 90 and 140 – and two or 

three neutrons. (The mass number is the sum of the number of protons and neutrons in the nucleus of the 

atom.) These neutrons can cause additional fissions (producing a chain reaction), escape from the reactor, 

or irradiate nearby materials. Zirconium is used in the cladding of nuclear fuel elements, and neutron 

irradiation of the cladding can produce zirconium-93. Zirconium-93 is also a fission product with a 

relatively high yield of about 6%. That is, about six atoms of zirconium-93 are produced per 100 fissions. 

Zirconium-93 is present in spent nuclear fuel, high-level radioactive wastes resulting from processing 

spent nuclear fuel, and radioactive wastes associated with nuclear reactors and fuel reprocessing plants. 

How Is It Used? Zirconium is very resistant to corrosion that occurs in the presence of many common 

acids and alkalis as well as seawater. For this reason it is used extensively by the chemical industry 

Zirconium 
where corrosive agents are employed. Zirconium is used in vacuum tubes, as an alloying agent in steel, in 

surgical appliances, photoflash bulbs, explosive primers, and lamp filaments. It is also used in poison ivy 

lotions in the form of the carbonate. The mineral zircon has a high index of refraction and is used as a 

gem material. The impure oxide zirconia is used for laboratory crucibles that can withstand heat shock, 

for the lining of metallurgical furnaces, and by the glass and ceramic industries as a refractory material. 

Zircaloy, an alloy of zirconium containing small amounts of tin, iron, chromium, and nickel, is used as a 

cladding material for nuclear fuel elements. Zirconium has a low absorption cross-section for neutrons, 

which makes it an ideal material for use in nuclear reactor applications. Its use in commercial nuclear 

power generation now accounts for as much as 90% of the zirconium metal that is produced. 

What’s in the Environment? Zirconium is present in the earth’s crust at a concentration of about 

130 milligrams per kilogram (mg/kg), and its concentration in seawater is about 0.026 micrograms 

(μg)/liter. Trace amounts of zirconium-93 are present in soil around the globe from radioactive fallout. It 

can also be present at certain nuclear facilities, such as reactors and spent fuel 

reprocessing plants. Zirconium is generally one of the less mobile radioactive metals in 

soil, although certain forms can move downward some distance to underlying layers 

with percolating water. Zirconium preferentially adheres quite well to soil, and the 

concentration associated with sandy soil particles is typically about 600 times higher 

than in interstitial water (water in the pore space between soil particles), with higher 

concentration ratios (over 2,000) in loam and clay soils. Thus, zirconium is generally not a major 

contaminant in groundwater at DOE sites. 

What Happens to It in the Body? Zirconium can be taken into the body by eating food, drinking 

water, or breathing air. Gastrointestinal absorption from food or water is the principal source of internally 

deposited zirconium in the general population. Zirconium is not well absorbed into the body, with only 

about 0.2% of the amount ingested being absorbed into the bloodstream through the intestines. Of the 

zirconium that reaches the blood, half deposits in the skeleton with a biological half-life of 8,000 days and 

the other half deposits in all other organs and tissues of the body where it is retained with a biological 

half-life of 7 days (per simplified models that do not reflect intermediate redistribution). Since zirconium 

is not a major constituent of mineral bone, the amount deposited in the skeleton is assumed to remain on 

the bone surfaces and not be absorbed into the volume of bone. 

 

What Are the Primary Health Effects? Zirconium is a health hazard only if it is taken into the body. 

External gamma exposure is not a concern because zirconium-93 decays by emitting a beta particle with 

no gamma radiation, and niobium-93m decays by 

isomeric transition, in which only low-energy 

gamma radiation is emitted. While inside the 

body, zirconium presents a health hazard from the 

beta particles and gamma radiation, and the main 

concern is associated with the increased 

likelihood of inducing cancer. 

What Is the Risk? Lifetime cancer mortality 

risk coefficients have been calculated for nearly 

all radionuclides, including zirconium (see box at 

right). While the coefficients for ingestion are 

somewhat lower than for inhalation, ingestion is 

generally the most common means of entry into 

the body. The risks shown here include the 

contribution from niobium-93m, even though it 

may not be in radioactive equilibrium with 

zirconium-93 (about ten niobium half-lives or 140 years are needed to establish equilibrium). Similar to 

other radionuclides, the risk coefficients for tap water are about 75% of those for dietary ingestion. 

Radiological Risk Coefficients 

This table provides selected risk coefficients for 

inhalation and ingestion. Recommended default 

absorption types were used for inhalation, and dietary 

values were used for ingestion. These values include 

the contribution from niobium-93m. Risks are for 

lifetime cancer mortality per unit intake (picocurie, 

pCi), averaged over all ages and both genders (10-12 is 

a trillionth). Other values, including for morbidity, 

are also available. 

Lifetime Cancer Mortality Risk 

Isotope Inhalation 

(pCi-1) 

Ingestion 

(pCi-1) 

Zirconium-93 8.4 × 10-12 1.7 × 10-12

For more information, see the companion fact sheet on 

Radioactive Properties, Internal Distribution, and Risk 

Coefficients and the accompanying Table 1. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

 

What Is It? Ionizing radiation is energy that is carried by any of several types of particles and rays (electromagnetic 

radiation) given off by radioactive material, X-ray machines, and nuclear reactions. This energy can knock electrons out of 

molecules with which they interact, thus creating ions. Non-ionizing radiation, such as that emitted by a laser, is different 

because it does not create ions when it interacts with matter but dissipates energy generally in the form of heat. The three main 

types of ionizing radiation are alpha particles, beta particles, and gamma rays. 

An alpha particle consists of two protons and two neutrons and is identical to the 

nucleus of a helium atom. Because of its relatively large mass and charge, an alpha 

particle produces ions in a very localized area. An alpha particle loses some of its 

energy each time it produces an ion (its positive charge pulls electrons away from 

atoms in its path), finally acquiring two electrons from an atom at the end of its 

path to become a complete helium atom. An alpha particle has a short range 

(several centimeters) in air and cannot penetrate the outer layer of skin. 

Beta particles can be either negative (negatron) or positive (positron). Negatrons 

are identical to electrons and originate in the nucleus of an atom that undergoes 

radioactive decay by changing a neutron into a proton. The only difference 

between a negative beta particle (negatron) and an electron is the ancestry. A beta particle originates in the nucleus whereas an 

electron is external to the nucleus. Unless otherwise specified, the term “beta particle” generally refers to a negatron. A 

positron is emitted from an atom that decays by changing a proton into a neutron. Beta particles are smaller and more 

penetrating than alpha particles, but their range in tissue is still quite limited. When its energy is spent, a negatron attaches 

itself to an atom and becomes an ordinary electron, while a positron collides with an ambient electron and 

the two particles annihilate each other, producing two gamma rays. When a negatron passes close to the 

nucleus of an atom, the strong attractive Coulomb force causes the beta particle to deviate sharply and lose 

energy at a rate proportional to the square of the acceleration. This energy manifests itself as photons termed 

Bremsstrahlung. The amount of beta energy converted into photons is directly proportional to the energy of 

the beta particle. This effect is only significant for high-energy beta particles generally passing through very 

dense materials such as lead, i.e., those with higher atomic numbers and so more protons in the nucleus. 

Gamma rays are electromagnetic radiation given off by an atom as a means of releasing excess energy. They are bundles 

(quanta) of energy that have no charge or mass and can travel long distances through air (up to several hundred meters), body 

tissue, and other materials. A gamma ray can pass through a body without hitting anything, or it may hit an atom and give that

atom all or part of its energy. This normally knocks an electron out of the atom, ionizing it. This electron then uses the energy 

it receives from the gamma ray to create additional ions by knocking electrons out of other atoms. Because a gamma ray is 

pure energy, it no longer exists once it loses all its energy. The capability of a gamma ray to do damage is a function of its

energy, where the distance between ionizing events is large on the scale of the nucleus of a cell. 

Additional forms of ionizing radiation beyond the three types shown in the figure above include neutrons, protons, neutrinos, 

muons, pions, heavy charged particles, X-rays and others. Essentially all radioactive materials at the Hanford Site originated

from neutron interactions with uranium fuel to produce plutonium. Byproducts of this process include fission products (most 

of which are in the high-level waste currently in on-site storage), activation products in the containment and reactor coolant 

materials, and various radioactive wastes. However, the radioactive hazards that remain at the Hanford Site are largely those 

associated with the three general types of radiation shown above, so the discussion here is limited to these three. 

Where Does It Come From? All organisms are being exposed to ionizing radiation from natural sources all the time. 

Radiation doses are typically given in units of rem – an acronym for Roentgen equivalent man – or millirem (mrem), which is 

one one-thousandth of a rem. This unit was developed to allow for the consistent reporting of hazards associated with the 

various types and energies of radiation on the human body. The rem is the product of the absorbed dose in rads (i.e., the 

amount of energy imparted to tissue by the radiation, where 1 rad equals 0.01 joules/kg) and factors for the relative biological 

effectiveness (RBE) of the radiation. The RBE is directly related to the linear energy transfer (LET) or distance over which the 

radiation energy is imparted to the absorbing medium and is accounted for by a quality factor. For example, alpha particles are 

20 times more hazardous than beta particles for the same energy deposition and hence have a quality factor of 20, whereas the 

quality factor for beta particles is one. The International Commission on Radiological Protection (ICRP) has developed a 

methodology for reporting the effective dose equivalent. This is the product of the dose (in rem or mrem) to individual tissues 

and the tissue-specific weighting factors (fractional values less than one) that indicate the relative risk of cancer induction or 

hereditary defects from irradiation of that tissue, summed over all relevant tissues. By use of the effective dose equivalent, it is 

possible to compare the relative radiation hazards from various types of radiation that impact different organs of the body. The 

doses discussed below are effective dose equivalents. 

Paper

Ionizing Radiation 

Glass 

Lead 

Alpha 

Beta 

Gamma 
 

More than 80% of the total dose received by individuals in the United States comes from natural sources; 55% comes from 

radon gas decay products, 8% from cosmic radiation, 8% from terrestrial radiation, and 11% from internal sources such as 

potassium-40. The rest comes from man-made sources that include medical X-rays (11%), nuclear medicine (4%), consumer 

products (3%), radioactive fallout, and nuclear power plants (<1%). These percentages are representative, and not all people 

are exposed to the same sources to the same degree. The National Council on Radiation Protection and Measurements (NCRP) 

estimates an average radiation dose of 360 mrem/yr from natural and man-made sources, with about 300 mrem/yr from natural 

sources and 60 mrem/yr from man-made sources (mostly for medical procedures). For reference, a person receives a dose of 

about 5 mrem on a flight from New York City to Los Angeles, and a typical chest X-ray produces a dose of about 10 mrem. 

How Is It Used? Radioactive materials and other sources of ionizing radiation are widely used to diagnose and treat diseases 

in human and veterinary medicine. Medical and dental X-rays are used to detect problems such as broken bones and dental 

cavities. Sealed radiation sources are used to deliver very high, localized radiation doses to treat certain types of cancers. 

Ionizing radiation also has a number of industrial and commercial uses. Radioactive sources are used in consumer products 

such as smoke detectors and to sterilize food products. Ionizing radiation is also used to test materials, inspect welds, generate 

heat and electricity for space travel, determine soil moisture content, and track the movement of various elements in the 

environment and human body (through use of radioactive tracers). Additional uses continue to be identified. 

What’s in the Environment? Exposure to background radiation and naturally occurring radioactive materials results in an 

annual dose of about 300 mrem/yr. Of this total, about one-third is due to external ionizing radiation, the main contributors 

being cosmic rays and terrestrial gamma rays (29 mrem/yr each), and radionuclides within the body (40 mrem/yr). Cosmic 

rays are produced when subatomic particles originating outside the solar system interact with particles in the 

upper atmosphere to produce gamma rays, neutrons and leptons that can reach and penetrate the earth’s surface. 

It is these secondary particles and rays that produce the dose from cosmic radiation. Naturally occurring 

radioactive elements such as uranium, thorium and radium are present in soil, rock, water, and all other 

environmental media, and certain of these radionuclides (and their radioactive decay products) give off gamma 

rays as they undergo radioactive decay. The principal contributor to the dose from radionuclides in the body is potassium-40, 

which decays by emitting an energetic beta particle and gamma rays. About two-thirds of the background radiation dose (of 

300 mrem/yr) is due to intake of radionuclides into the body. The largest contributor is inhalation of radon-220 and radon-222

gases and their short-lived radioactive decay products, which are charged particles that readily attach to airborne dust particles. 

In fact, inhalation of radon gas contributes nearly all of the 200 mrem/yr attributable to radionuclide intake, and most is due to 

the short-lived radon-222 decay products polonium-218 and polonium-214. Ingestion of food and water containing naturally 

occurring radionuclides accounts for only a few mrem/yr. 

 

What Are the Primary Health Effects? High doses of ionizing radiation can lead to effects such as skin burns, hair loss, 

birth defects, illness, cancer, and death, depending on the dose and the period of time over which it is received. Acute doses

(such as from a serious accident involving nuclear materials) can result in damage to the blood-forming organs, gastrointestinal 

tract, and central nervous system. Very high doses, e.g., on the order of 500 rem (or 500,000 mrem) or more, can cause death 

in many (but not all) individuals, depending on the degree of medical intervention. The main health concern associated with 

radiation exposure is the induction of various cancers. Additional effects may include genetic mutations (although none have 

been observed in humans) and teratogenic effects such as mental retardation. The U.S. Environmental Protection Agency 

(EPA) has indicated that for radioactively contaminated Superfund sites, the risk from cancer is limiting and should be used as

the sole basis for assessing radiation-related human health risks. The EPA noted that on average, about 50% of all cancers that 

can be induced by radiation (they can also be caused by other agents) are fatal, ranging from about 10% for thyroid cancer to 

100% for liver cancer. Other EPA estimates for specific isotopes indicate the average may be higher, e.g., 60 to 70% or more. 

What Happens to It in the Body? Radioactive materials can enter the body by inhalation, ingestion, or dermal absorption. 

In addition, gamma radiation external to the body can penetrate the skin and produce a dose in various tissues. Inhalation is the 

primary exposure mode for gaseous radionuclides (such as radon), and particulates near the source of an airborne release. A 

fractional amount of inhaled radionuclides is transferred from the lungs to the blood, where it distributes to other organs. The 

extent of absorption is strongly dependent on the radionuclide and its chemical form. Ingestion is the primary uptake mode for

radionuclides in soil, water, and food, including those naturally occurring (such as radium and uranium in soil and 

groundwater) and man-made (such as plutonium from radioactive fallout). A fractional amount of ingested radionuclides is 

absorbed from the gastrointestinal tract into the blood while the rest clears the body through normal biological processes via 

urine and feces. As with inhalation, the extent of uptake depends on the radionuclide and its chemical form. The skin is 

generally an effective barrier against absorption of radionuclides, so dermal absorption is a very minor route of exposure. An

exception to this is dermal absorption of tritiated water, i.e., water containing some amount of tritium (hydrogen-3) in place of 

a normal hydrogen atom in the water molecule, which is absorbed through the skin in the same manner as ordinary water. 

What Is the Risk? While the EPA has developed lifetime cancer mortality risk coefficients for nearly all radionuclides, the 

agency has not developed a risk coefficient for ionizing radiation as a general category. A nominal mortality value of 5 × 10-7

incremental cancer risk per mrem has been identified for low-LET radiation delivered at a low dose and dose rate.
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Natural Decay Series: 

Uranium, Radium, and Thorium

Uranium, radium, and thorium occur in three natural decay series, headed by uranium-238, thorium-232, 

and uranium-235, respectively. In nature, the radionuclides in these three series are approximately in a 

state of secular equilibrium, in which the activities of all radionuclides within each series are nearly equal. 

Two conditions are necessary for secular equilibrium. First, the parent radionuclide must have a half-life 

much longer than that of any other radionuclide in the series. Second, a sufficiently long period of time 

must have elapsed, for example ten half-lives of the decay product having the longest half-life, to allow 

for ingrowth of the decay products (see the companion fact sheet on Ionizing Radiation). Under secular

equilibrium, the activity of the parent radionuclide undergoes no appreciable changes during many halflives of its decay products. 

The radionuclides of the uranium-238, thorium-232, and uranium-235 decay series are shown in Figures

N.1, N.2, and N.3, along with the major mode of radioactive decay for each. Radioactive decay occurs 

when an unstable (radioactive) isotope transforms to a more stable isotope, generally by emitting a 

subatomic particle such as an alpha or beta particle. Radionuclides that give rise to alpha and beta 

particles are shown in these figures, as are those that emit significant gamma radiation. 

Gamma radiation is not a mode of radioactive decay (such as alpha and beta decay). Rather, it is a

mechanism by which excess energy is emitted from certain radionuclides, i.e., as highly energetic 

electromagnetic radiation emitted from the nucleus of the atom. For simplicity, only significant gamma 

emissions associated with the major decay modes are shown in Figures N.1 through N.3; that is, 

radionuclides listed are those for which the radiation dose associated with gamma rays may pose a health 

concern. The gamma component is not shown for those radionuclides whose gamma emissions do not 

generally represent a concern. 

Of the two conditions noted above for secular equilibrium, the first is generally met for the uranium-238, 

thorium-232 and uranium-235 decay series in naturally occurring ores. While the second condition may

not be met for all ores or other deposits of uranium and thorium (given the extremely long half-lives for 

the radionuclides involved and the geological changes that occur over similar time scales), it is reasonable 

to assume secular equilibrium for naturally occurring ores to estimate the concentrations of the various 

daughter radionuclides that accompany the parent. The state of secular equilibrium in natural uranium 

and thorium ores is significantly altered when they are processed to extract specific radionuclides. 

After processing, radionuclides with half-lives less than one year will reestablish equilibrium conditions

with their longer-lived parent radionuclides within several years. For this reason, at processing sites what

was once a single, long decay series (for example the series for uranium-238) may be present as several

smaller decay series headed by the longer-lived decay products of the original series (that is, headed by 

uranium-238, uranium-234, thorium-230, radium-226, and lead-210 in the case of uranium-238). Each of 

these sub-series can be considered to represent a new, separate decay series. Understanding the physical

and chemical processes associated with materials containing uranium, thorium, and radium is important 

when addressing associated radiological risks. 

In the fact sheets developed for uranium, radium, and thorium, the contributions of radionuclides having 

half-lives less than one year were included in the risk coefficients. (Each fact sheet identifies which 

radionuclides are included in these coefficients.) In some situations, it may be necessary to add the 

radiological risk identified for a given radionuclide to that of its parent radionuclide to properly represent 

the total risk. For example, the radiological risk for thorium-232 is comprised of the risk for thorium-232 

plus the risk for radium-228. Decay series information should be used together with the information in 

these fact sheets to ensure that the radiological risks associated with uranium, radium, and thorium are 

properly estimated and represented. 
α 77,000 years

Protactinium-234m

FIGURE N.1 Natural Decay Series: Uranium-238

β

Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

NOTES:

The symbols α and β indicate alpha and beta 

decay, and the times shown are half-lives. 

An asterisk indicates that the isotope is also 

a significant gamma emitter. 

Uranium-238 also decays by spontaneous 

fission.

Uranium-234

Thorium-230

α 240,000 years

Radium-226

Radon-222

α 1,600 years

α 3.8 days

Uranium-238

Thorium-234

α 4.5 billion years

24 days

β

1.2 minutes

Bismuth-210

Polonium-210

α 140 days

22 years

β

5.0 days

Lead-206 (stable)

Bismuth-214*

Polonium-214

α 160 microseconds

27 minutes

β

20 minutes

Lead-210

Polonium-218

α 3.1 minutes

Lead-214*

β β
FIGURE N.2 Natural Decay Series: Uranium-235

Uranium-235*

Protactinium-231*

Actinium-227

Thorium-231

α

α

700 million years

33,000 years

26 hours

β

Thorium-227*

Radium-223*

α 19 days

22 years

(99%)

β

Lead-211*

Bismuth-211*

36 minutes

β

Polonium-215

α 1.8 milliseconds

NOTES:

The symbols α and β indicate alpha and beta 

decay, and the times shown are half-lives. 

An asterisk indicates that the isotope is also 

a significant gamma emitter. 

Radon-219*

α 4.0 seconds

Francium-223*

α

22 minutes

β α 11 days

22 years

(1%)

Thallium-207

Lead-207 (stable)

4.8 minutes

β

α 2.1 minutes

Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005
α 3.7 days

Actinium-228*

β

Thorium-228

Radium-224

α 1.9 years

Radon-220

α 56 seconds

Thorium-232

Radium-228

α 14 billion years

5.8 years

β

6.1 hours

Polonium-212

α 310 nanoseconds

Lead-208 (stable)

Bismuth-212*

11 hours

β

61 minutes

(64%)

Polonium-216

α 0.15 seconds

Lead-212*

β 61 minutes 

(36%)

α

Thallium-208*

β

3.1 minutes

FIGURE N.3 Natural Decay Series: Thorium-232

NOTES:

The symbols α and β indicate alpha and beta 

decay, and the times shown are half-lives. 

An asterisk indicates that the isotope is also 

a significant gamma emitter. 

Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

What Are They? Taking their name from being trans- or beyond uranium, transuranic radionuclides 

have atomic numbers greater than that of uranium, which is 92. All transuranic isotopes are radioactive. 

(Isotopes are different forms of an element that have the same number of protons in the nucleus but a 

different number of neutrons.) Those associated with americium, californium, curium, neptunium, and 

plutonium are generally of concern for U.S. Department of Energy (DOE) environmental management 

sites such as Hanford. More information on these five radionuclides and their decay products is presented 

in the separate radionuclide-specific fact sheets. An additional transuranic radionuclide that can be found 

at DOE sites is berkelium-247; summary radioactive properties for this transuranic isotope is given in the 

green box on the back of this page. 

Where Do They Come From? Most transuranic radionuclides at DOE sites were produced in nuclear 

reactors by neutron capture, with the remainder being produced in particle accelerators. Because multiple 

neutron capture reactions are usually needed to produce higher-numbered isotopes for a given element, 

the relative number of individual isotopes decreases with increasing atomic number. For example, after 

nuclear fuel is discharged from a plutonium production reactor, more plutonium-238 and plutonium-239 

would generally be present than the higher-numbered isotopes such as plutonium-242. Transuranic 

radionuclides typically do not occur naturally in the environment, although very minute amounts can be 

present with some uranium ores. The underground, natural sustained nuclear reactions estimated to have 

occurred about 1.9 billion years ago in Gabon, Africa, also produced some transuranic radionuclides, 

including neptunium and plutonium. 

Like certain other radioactive isotopes such as those of naturally occurring uranium, radium, and thorium, 

the transuranic radionuclides undergo radioactive decay to create typically long chains of decay products. 

Radioactive decay occurs when an unstable atom transforms to a more stable isotope by emitting an alpha 

or beta particle, although some isotopes also decay by spontaneous fission. During spontaneous fission, 

the atom self-disintegrates into two smaller atoms accompanied by a release of energy. The radioactive 

decay series for several key transuranic radionuclides are shown in Figures T.1 through T.3 at the end of 

this fact sheet. These figures illustrate the relationships among isotopes in terms of radioactive decay and 

ingrowth. They also identify the major decay modes and show which isotopes emit significant gamma 

radiation. A number of these transuranic decay series eventually transition into the decay series for the 

naturally occurring radionuclides, which are shown separately in Figures N.1 through N.3 within the 

companion fact sheet: Natural Decay Series: Uranium, Radium, and Thorium. Many concepts relevant 

to decay series, such as secular equilibrium, are discussed in that companion fact sheet and are therefore 

not repeated here. Interested readers are encouraged to refer to that sheet for more information. 

What Are Transuranic Wastes? Radioactive wastes containing more than 100 nanocuries per gram 

(nCi/g) of alpha-emitting transuranic radionuclides with half-lives greater than 20 years are termed 

“transuranic wastes.” (100 nCi/g is one tenth of a microcurie or one ten-millionth of a curie per gram.) 

This term technically applies to radioactive wastes generated after 1970. Prior to 1970, no specific 

considerations were identified for radioactive wastes containing transuranic radionuclides. At that time, 

these wastes were managed in the same manner as other low-level radioactive wastes regardless of the 

concentration of transuranic radionuclides, and all these wastes were generally disposed of by shallow 

land burial. Today, transuranic wastes are segregated from other low-level radioactive wastes and 

managed separately because the hazards associated with the higher concentrations of long-lived alphaemitting radionuclides are believed to warrant more stringent handling and disposal considerations. 

Radioactive wastes that would meet today’s definition of transuranic wastes but were disposed of before 

1970 by shallow land burial and other similar techniques at DOE sites are often referred to as “buried 

transuranic wastes.” The DOE is currently evaluating the most appropriate manner to manage these 

wastes over the long term on a site-specific basis. Transuranic (or TRU) wastes generated by more recent 

defense-related activities, including a substantial inventory in controlled storage, are currently being sent 

to the Waste Isolation Pilot Plant located near Carlsbad, New Mexico, for deep underground disposal. 

This facility began accepting wastes in 1999 and has a permitted capacity of about 175,600 m3

. 

Transuranic Radionuclides 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

The following table summarizes key properties of selected transuranic radionuclides. (For americium, 

californium, curium, neptunium, and plutonium, this information is also given in the radionuclide-specific 

fact sheets and the supporting Table 2.) The table below includes radionuclides beyond those that define 

transuranic wastes, i.e., those with half-lives <20 years or that decay by means other than alpha emission. 

(These are indicated by regular green shading, while those that define TRU wastes are in yellow-green.) 

What Are the Health Risks? Transuranic wastes pose a concern for environmental management 

because of the relatively long half-lives of the transuranic radionuclides, combined with their emission of 

alpha particles (and to a lesser degree beta particles). Because the alpha particles cannot penetrate the 

outer layer of skin, these radionuclides pose a hazard primarily if they are taken into the body. Some 

isotopes also emit gamma radiation, which results in an additional hazard from external exposure. The 

main concern is cancer induction due to energy deposited in tissues during radioactive decay. Lifetime 

cancer mortality risk coefficients have been calculated for nearly all radionuclides, including transuranic 

Radioactive Properties of Selected Transuranic Radionuclides 

Radiation Energy (MeV) Isotope Half-Life Specific 

Activity (Ci/g)

Decay

Mode Alpha (α) Beta (β) Gamma (γ)

Americium-241 430 yr 3.5 α 5.5 0.052 0.033 

Americium-242m 150 yr 9.8 IT 0.025 0.044 0.0051 

Americium-242 16 hr 820,000 β, EC - 0.18 0.018 

Americium-243 7,400 yr 0.20 α 5.3 0.022 0.056 

Neptunium-239 2.4 days 230,000 β - 0.26 0.17 

Berkelium-247 1,400 yr 1.1 α 5.6 0.061 0.11 

Californium-248 330 days 1,600 α 6.3 0.0060 0.0013 

Californium-249 350 yr 4.1 α 5.8 0.043 0.33 

Californium-250 13 yr 110 α 6.0 0.0057 0.0012 

Californium-251 900 yr 1.6 α 5.8 0.020 0.13 

Californium-252 2.6 yr 540 α 5.9 0.0056 0.0012 

Curium-242 160 days 3,400 α 6.1 0.010 0.0018 

Curium-243 29 yr 52 α 5.8 0.14 0.13 

Curium-244 18 yr 82 α 5.8 0.086 0.0017 

Curium-245 8,500 yr 0.17 α 5.4 0.065 0.096 

Curium-246 4,700 yr 0.31 α 5.4 0.0080 0.0015 

Curium-247 16 million yr 0.000094 α 4.9 0.021 0.32 

Curium-248 340,000 yr 0.0043 α 4.7 0.0060 0.0012 

Curium-250 6,900 yr 0.21 α, β 1.3 0.0016 - 

Plutonium-246 (25%) 11 days 49,000 β - 0.13 0.14 

Berkelium-250 (14%) 3.2 hr 3.9 million β - 0.29 0.89 

Americium-246 (25%) 39 min 20 million β - 0.66 0.70 

Neptunium-235 1.1 yr 1,400 EC <0.001 0.010 0.0071 

Neptunium-236 120,000 yr 0.013 β, EC - 0.21 0.14 

Plutonium-236 (9%) 2.9 yr 540 α 5.8 0.013 0.0021 

Neptunium-237 2.1 million yr 0.00071 α 4.8 0.070 0.035 

Protactinium-233 27 days 21,000 β - 0.20 0.20 

Plutonium-238 88 yr 17 α 5.5 0.011 0.0018 

Plutonium-239 24,000 yr 0.063 α 5.1 0.0067 <0.001 

Plutonium-240 6,500 yr 0.23 α 5.2 0.011 0.0017 

Plutonium-241 14 yr 100 β <0.001 0.0052 <0.001 

Plutonium-242 380,000 yr 0.0040 α 4.9 0.0087 0.0014 

Plutonium-244 83 million yr 0.000018 α 4.6 0.0071 0.0012 

Decay products are in italics. Values are given to two significant figures. EC = electron capture, IT = isomeric 

transition, Ci = curie, g = gram, and MeV = million electron volts. A dash indicates that the entry is not applicable. 

(See radionuclide-specific fact sheets and the companion fact sheet on Radioactive Properties, Internal Distribution, 

and Risk Coefficients for further information, including on terminology and interpretation of radiation energies.) 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

isotopes. Selected radiological risk coefficients are summarized in the yellow box below. This summary 

includes information presented on the radionuclide-specific fact sheets for americium, californium, 

curium, neptunium, and plutonium and the supporting Table 1. Information for berkelium-247 is taken 

from the same source reference used for that table. Shown below are selected risk coefficients of various 

transuranic radionuclides for inhalation and dietary ingestion, and for external gamma irradiation where 

that entry is appropriate. The mortality risk represents the lifetime risk of incurring a fatal cancer, and the 

morbidity risk represents the risk of incurring both fatal and non-fatal cancers. 

Mortality and Morbidity Risk Coefficients for Selected Transuranic Radionuclides 

Lifetime Cancer Risk 

Isotope Mortality Morbidity 

Inhalation Ingestion External Inhalation Ingestion External 

Americium-241 2.4 × 10-8 9.5 × 10-11 1.9 × 10-8 2.8 × 10-8 1.3 × 10-10 2.8 × 10-8

Americium-242m 1.3 × 10-8 6.8 × 10-11 - 1.6 × 10-8 9.0 × 10-11 - 

Americium-243 2.3 × 10-8 9.8 × 10-11 4.3 × 10-7 2.7 × 10-8 1.4 × 10-10 6.4 × 10-7

Berkelium-247 4.0 × 10-8 1.2 × 10-10 2.1 × 10-7 4.8 × 10-8 1.6 × 10-10 3.1 × 10-7

Californium-248 2.4 × 10-8 3.8 × 10-11 - 2.6 × 10-8 6.2 × 10-11 - 

Californium-249 4.0 × 10-8 1.2 × 10-10 9.3 × 10-7 4.8 × 10-8 1.6 × 10-10 1.4 × 10-6

Californium-250 3.5 × 10-8 8.0 × 10-11 - 3.7 × 10-8 1.1 × 10-10 - 

Californium-251 4.1 × 10-8 1.3 × 10-10 2.6 × 10-7 4.9 × 10-8 1.7 × 10-10 3.8 × 10-7

Californium-252 2.1 × 10-8 4.1 × 10-11 - 2.2 × 10-8 5.7 × 10-11 - 

Curium-242 1.4 × 10-8 3.2 × 10-11 - 1.5 × 10-8 5.5 × 10-11 - 

Curium-243 2.4 × 10-8 8.5 × 10-11 2.9 × 10-7 2.7 × 10-8 1.2 × 10-10 4.2 × 10-7

Curium-244 2.3 × 10-8 7.5 × 10-11 - 2.5 × 10-8 1.1 × 10-10 - 

Curium-245 2.4 × 10-8 9.5 × 10-11 1.6 × 10-7 2.8 × 10-8 1.3 × 10-10 2.4 × 10-7

Curium-246 2.4 × 10-8 9.3 × 10-11 - 2.8 × 10-8 1.3 × 10-10 - 

Curium-247 2.2 × 10-8 9.0 × 10-11 8.9 × 10-7 2.5 × 10-8 1.3 × 10-11 1.3 × 10-6

Curium-248 8.8 × 10-8 3.4 × 10-10 - 1.0 × 10-7 4.8 × 10-10 - 

Curium-250 5.0 × 10-7 2.0 × 10-9 9.7 × 10-7 5.8 × 10-7 2.8 × 10-9 1.4 × 10-6

Neptunium-235 1.0 × 10-12 2.8 × 10-13 - 1.2 × 10-12 5.1 × 10-13 - 

Neptunium-236 2.6 × 10-9 1.5 × 10-11 2.2 × 10-7 3.0 × 10-9 2.3 × 10-11 3.2 × 10-7

Neptunium-237 1.5 × 10-8 5.8 × 10-11 5.4 × 10-7 1.8 × 10-8 9.1 × 10-11 8.0 × 10-7

Plutonium-238 3.0 × 10-8 1.3 × 10-10 - 3.4 × 10-8 1.7 × 10-10 - 

Plutonium-239 2.9 × 10-8 1.3 × 10-10 - 3.3 × 10-8 1.7 × 10-10 - 

Plutonium-240 2.9 × 10-8 1.3 × 10-10 - 3.3 × 10-8 1.7 × 10-10 - 

Plutonium-241 2.8 × 10-10 1.9 × 10-12 - 3.3 × 10-10 2.3 × 10-12 - 

Plutonium-242 2.8 × 10-8 1.3 × 10-10 - 3.1 × 10-8 1.7 × 10-10 - 

Plutonium-244 2.7 × 10-8 1.3 × 10-10 - 3.1 × 10-8 1.6 × 10-10 - 

Source: Cancer Risk Coefficients for Environmental Exposure to Radionuclides, Federal Guidance Report 13, 

U.S. Environmental Protection Agency, EPA 402-R-99-001, September 1999. Values are averaged over all ages 

and both genders and include the contributions from short-lived decay products, i.e., those with half-lives less than 

one year. (For context, 10-9 is a billionth, 10-12 is a trillionth, and a pCi is a picocurie, or a trillionth of a curie.) To 

convert to standard international units, multiply by 27 pCi per becquerel (Bq). For ingestion and inhalation, units 

are risk per pCi. For inhalation, the values corresponding to the recommended default absorption type for 

particulates are shown; the maximum value is given if no absorption type was recommended. For ingestion, the 

dietary values shown are the highest for ingestion exposures; values for tap water ingestion are typically 70 to 80% 

of those for diet. For external exposure, risk coefficients are given for those radionuclides with gamma-ray energies 

>0.03 MeV per decay, accounting for the fraction of time that the radioactive decay results in the emission of 

gamma rays. A dash indicates the radionuclide or its decay products does not emit significant gamma radiation (see 

the fact sheet for Radioactive Properties, Internal Distribution, and Risk Coefficients). Units for external gamma 

risk coefficients shown in the table are risk per pCi/g soil for one year of exposure. Risk coefficients do not exist for 

curium-248, curium-250, or plutonium-244, so values shown have been derived (see the Curium and Plutonium fact 

sheets for derivation approach); the external value for curium-250 is attributable to its short-lived decay products. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Two parent radionuclides are shown here.

Curium-242

Americium-242m

See Figure N.1 

in the companion 

fact sheet for the 

natural decay series.

Plutonium-238

Uranium-234

Neptunium-238*

α

α

α

150 years

(0.5%)

160 days

2.1 days

β 88 years

Berkelium-250*

Curium-250

Americium-246*

Plutonium-246*

α

11 days

β

6,900 years

6,900 years

β

3.2 hours

β

Californium-250

See Figure N.1 

in the companion 

fact sheet for the 

natural decay series.

Curium-246

Uranium-238

α

α

13 years

4,700 years

39 minutes

β

25%

14%

Plutonium-242

α 380,000 years

NOTES:

The symbols α and β indicate alpha and beta

decay, and the times shown are half-lives. 

An asterisk indicates that the isotope is also

a significant gamma emitter. 

Plutonium-238, plutonium-242, curium-242,

curium-246, curium-250, and californium-250

also decay by spontaneous fission.

FIGURE T.1 Transuranic Decay Series: Americium-242m and Curium-242, and Curium-250
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Three parent radionuclides are shown here.

FIGURE T.2 Transuranic Decay Series: Californium-251, Berkelium-247, and Curium-243; 

Californium-252 and Californium-248

Berkelium-247*

Curium-247*

Americium-243*

Neptunium-239*

Plutonium-243

α

α

α

16 million years

5.0 hours

β

1,400 years

7,400 years

Californium-251*

α 900 years

Curium-243*

Plutonium-239

Uranium-235*

α

α

2.4 days

β

29 years

24,000 years

See Figure N.2 in the companion 

fact sheet for the natural decay series.

Plutonium-244

Neptunium-240*

Uranium-240

α 83 million years

14 hours

β

Curium-248

α 340,000 years

Plutonium-240

See Figure N.3 

in the companion 

fact sheet for the 

natural decay series.

Uranium-236

Thorium-232

α

α

6,500 years 65 minutes

β

23 million years

Curium-244

α 18 years

Two parent radionuclides

are shown here.

Californium-248

α 330 days

Californium-252

α 2.6 years

NOTES:

The symbols α and β indicate alpha and beta

decay, and the times shown are half-lives. 

An asterisk indicates that the isotope is also

a significant gamma emitter. 

Plutonium-240, plutonium-244, curium-244, 

curium-248, californium-248, and californium-252 

also decay by spontaneous fission.
Protactinium-233*

Curium-245*

Americium-241*

Neptunium-237*

Plutonium-241

α

α

α

8,500 years

2.1 million years

430 years

14 years

β

Californium-249*

α 350 years

Uranium-233

Thorium-229*

Radium-225

α

α

160,000 years

7,300 years

27 days

β

Actinium-225

Francium-221*

Astatine-217

α

α

10 days

4.8 minutes

Bismuth-213*

Thallium-209*

46 minutes

β

Polonium-213

Lead-209

α 4.2 microseconds

2.2 minutes

β

Bismuth-209 (stable)

α 0.032 seconds

α 46 minutes 3.3 hours

β

15 years

β

NOTES:

The symbols α and β indicate alpha and beta

decay, and the times shown are half-lives. 

An asterisk indicates that the isotope is also

a significant gamma emitter. 

Californium-249 also decays by spontaneous 

fission.

98%

2%

Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

FIGURE T.3 Transuranic Decay Series: Californium-249
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

 

The fact sheets for the individual radionuclides contain summary-level information on the radioactive properties of the major 

isotopes of concern at Department of Energy (DOE) environmental management sites such as Hanford, as well as information on 

the potential health risks associated with exposure to these radionuclides. This information was developed using standard 

references and publications as identified in Table 3. The fact sheets are intentionally brief and it is not possible to include all 

relevant information associated with the highlights summarized in these sheets. This companion fact sheet has been prepared to

provide additional context to assist the reader in understanding the basis of the information presented in the individual fact sheets 

and to allow for proper interpretation of the radionuclide-specific data. 

Radioactive Properties. Each radionuclide fact sheet contains a table that summarizes the key properties of the various 

radioactive isotopes of that element. The information provided in that table includes the radioactive half-life, specific activity, 

radioactive decay modes, and the average energy of the emitted radiations. To simplify the presentation, all values have been 

rounded to two significant figures. Much of the information was obtained from Appendix G of Federal Guidance Report (FGR) 

Number 13 of the Environmental Protection Agency (EPA), Cancer Risk Coefficients for Environmental Exposure to 

Radionuclides, 402-R-99-001 (September 1999). These data were also checked with the more detailed decay information in 

International Commission on Radiological Protection (ICRP) Report 38, Radionuclide Transformations, Energy and Intensity of 

Emissions (1983). The ICRP report is a major source of the information given in Appendix G of FGR Number 13. 

Half-Life. The radioactive half-life is the length of time for a given amount of radioactive material to decrease to one half its 

initial amount by radioactive decay. Half-lives are constant for each radionuclide and can range from less than a second to 

billions of years. Only those radionuclides with half-lives longer than about one year are of concern for DOE environmental 

management sites, as shorter-lived radionuclides will have already decayed away to innocuous levels because production 

activities involving radioactive materials at major sites such as Hanford ceased more than ten years ago. 

Specific Activity. The specific activity is the activity per mass and is given in units of curies (Ci) per gram in the individual 

fact sheets. For reference, the specific activity of radium-226 is about 1 Ci per gram, and for context 1 gram of material is about 

0.035 ounce. The specific activities in the fact sheets (in units of curies per gram) were calculated using the following equation 

given in the Health Physics and Radiological Health Handbook (1992, p. 264). Specific Activity = ARa-226 × TRa-226 / Ai × Ti 

where ARa-22a = 226, the atomic weight of radium-226; TRa-226 = 1,600, the value used for the half-life of radium-226, in years (to 

two significant figures); Ai = the atomic weight of the isotope; and Ti = half-life of the isotope in years. The specific activity 

can be expressed in international units by multiplying the value in the fact sheet by 3.7 × 1010 becquerels (Bq) per Ci. 

Decay Mode. The radioactive decay modes identified in the fact sheets include beta-particle emission, alpha-particle emission, 

isomeric transition (IT), electron capture (EC), and spontaneous fission (SF). The companion fact sheet on Ionizing Radiation

contains additional information on the first two decay modes. The IT decay mode is a process whereby a nucleus in an elevated 

energy state (typically a metastable isotope designated by the letter “m”) releases excess energy by emitting a gamma ray. The

product of the decay is not a new isotope, but rather the same isotope in a reduced (more stable) energy configuration. The EC

decay mode is a process in which an inner-shell electron orbiting the nucleus of an atom is “captured” by the nucleus where it 

combines with a proton to become a neutron, and excess energy is given off in the form of gamma rays. An outer-shell electron 

fills the “hole” left in the inner shell, and the excess energy associated with the movement of an outer-shell electron to an inner 

shell is given off as X-rays. The SF decay mode is a process in which an unstable nucleus splits (fissions) into two smaller 

products without needing additional neutrons to initiate the process, i.e., it is spontaneous. For simplicity, only the major decay 

modes are shown in the fact sheets; decay modes that occur less than 1% of the time are not included. Report 38 of the ICRP 

includes a very detailed accounting of all decay modes for each radionuclide and can be consulted for additional information. 

In addition to the radioactive decay modes discussed above, there are additional mechanisms by which unstable atoms release 

energy. Internal conversion is a process in which the excess energy of a nucleus in an excited state is transferred to an electron 

orbiting the nucleus, which results in the electron being emitted from the atom. This process competes with gamma-ray emission

as a mechanism for releasing excess energy from the nucleus, such as occurs during IT decay. Both internal conversion and EC 

result in a “hole” in the inner shell of orbital electrons, which is filled by an outer-shell electron with excess energy given off in 

the form of X-rays. These X-rays can interact with other orbital electrons, transferring sufficient energy to them to result in the 

emission of additional electrons. Such emitted electrons generated by interactions with X-rays are termed Auger electrons and 

have very little kinetic energy. 

Decay Energy. The average energy reported for individual isotopes in the fact sheets represents the energy of the indicated 

radiation multiplied by the fractional yield for the given decay mode. That is, the energy reported for the various types of 

radiation represents the average energy per decay of the radionuclide. The energy of the radiation is given in units of million 

electron volts (MeV). One MeV is equal to 0.16 trillionth of a joule. The following two examples are provided to illustrate how 

this information should be interpreted. 

Radioactive Properties, Internal Distribution, and Risk Coefficients 
 

Consider a radionuclide that decays by emitting an alpha particle with two different energies: half the time the energy is 5 MeV 

and the other half of the time the energy is 6 MeV. The energy reported for the alpha particle decay in this case would be 

5.5 MeV. Consider a second example in which a radionuclide decays half the time by emitting a beta particle with an average 

energy of 0.5 MeV, and the other half of the time it decays by emitting an alpha particle with an energy of 6 MeV. The energy 

of the beta particle would be reported as 0.25 MeV (half of 0.5 MeV), and the energy of the alpha particle would be given as 

3 MeV (half of 6 MeV). As a note, the average energy of a beta particle is typically about one-third the maximum energy 

(which is often the energy reported in radionuclide charts), or about 30% of the maximum for negatrons and 40% for positrons. 

(See the companion fact sheet on Ionizing Radiation for a discussion of negatrons and positrons.) The summary-level tables in 

the individual radionuclide fact sheets include the contributions of all primary (alpha and beta particles and gamma rays) and 

secondary (X-rays and Auger electrons) radiations. 

The average energy reported for gamma rays in the fact sheets includes the contributions of X-rays and has been adjusted to 

account for the fractional yield. The only difference between these two types of electromagnetic radiation is their origin, and 

hence energy. Gamma rays originate in the nucleus as a means of releasing excess energy from the atom, while X-rays are 

emitted when electrons outside the nucleus move from higher to lower energy states. Radionuclides having gamma-ray energies 

less than 0.03 MeV per decay, considering the fractional yield (as described above), generally do not present a health concern 

from external gamma exposure. The average energy per decay reported for beta particles includes the contributions of all 

electrons and positrons regardless of their origin (internal or external to the nucleus). The average energy per decay reported for 

alpha particles does not include the contribution of the recoiling atom (which is typically quite small, e.g., a few percent of the 

total energy associated with the alpha-decay process). 

Some radionuclides decay into short-lived daughters that always accompany the parent. (The term parent is used to describe the 

original isotope, and daughter is used for the decay product.) For example, strontium-90 decays to yttrium-90 by emitting a beta 

particle with a 29-year half-life. The daughter yttrium-90 quickly decays by emitting a beta particle, with a half-life of 64 hours. 

So for all practical purposes, each decay of strontium-90 can be considered to yield two beta particles, one for strontium-90 and 

one for yttrium-90. Short-lived decay products need to be considered when estimating the potential health effects of exposures

to radionuclides. To facilitate this consideration, the radioactive properties of both the parent and its short-lived daughter(s) are 

presented in the individual radionuclide fact sheets. 

Internal Distribution. To estimate the human health risks associated with radionuclides, it is necessary to follow the 

movement of the isotopes from intake through excretion. These isotopes constantly emit radiation at a rate proportional to their 

specific activity as they pass through the body irradiating various organs. Some radionuclides very quickly deposit in one or two 

organs; others deposit more slowly throughout the entire body. Various models and computer codes have been developed by the 

ICRP, EPA, and other national and international organizations to estimate internal radiation doses and risks from intake of 

radionuclides. These models are based on extensive animal and human data and can be quite complex. A number of codes and 

models were considered by the EPA in developing the risk coefficients presented in FGR Number 13, as illustrated by the 

references in that document. The risk coefficients were calculated used the DCAL (Dose and Risk Calculation) software 

developed by Oak Ridge National Laboratory for the EPA. The DCAL is a comprehensive system for calculating radiation dose 

and risk coefficients using age-dependent models that incorporate information developed by the ICRP and other organizations on 

the distribution and retention of radionuclides by various organs in the body. The primary distribution of selected radionuclides 

in the body is shown in Figure 1 of the accompanying distribution fact sheet. 

Risk Coefficients. The EPA has developed mortality risk coefficients for nearly all radionuclides to estimate the lifetime risk 

of incurring a fatal cancer from environmental exposures using the DCAL software as described in FGR Number 13. These 

coefficients have been calculated by state-of-the-art methods and computer models that take into account age and gender 

dependence of intake, metabolism, dosimetry, and radiogenic risk, as well as competing causes of death, to estimate health risks 

from internal and external exposures. The values are given per unit uptake (picocurie, pCi) averaged over all ages and both 

genders. (For context, 10-9 is a billionth, 10-12 is a trillionth, and a pCi is a trillionth of a Ci.) To convert to standard 

international units, the given values should be multiplied by 27 pCi/Bq. 

Each radionuclide fact sheet contains a table with selected mortality risk coefficients for inhalation and ingestion, which are also 

summarized in Table 1. These values include the contributions from short-lived decay products, as identified on the radioactive 

properties summary table described above (e.g., the value for strontium-90 includes the contribution from yttrium-90). For 

inhalation, the values correspond to the recommended default absorption type for particulates, except as otherwise noted (e.g.,

the tritium values are for tritiated water). For ingestion, the dietary values shown are the highest for ingestion exposures; the 

values for tap water ingestion vary by radionuclide and are typically 70 to 80% of those for dietary intake. Coefficients are also 

available to estimate the risk of incurring all types of cancer (morbidity risk coefficients), and these values also vary by 

radionuclide. For most radionuclides, the ingestion mortality coefficients are on the order of 60 to 80% of the morbidity values, 

with iodine an exception at about 10%. For inhalation the percentages are a bit higher, ranging from 70% (for cesium and 

tritium) to nearly 100% (for uranium-234 and uranium-238), with the mortality coefficient for iodine again much lower at about 

10%. For major gamma-emitting radionuclides, risk coefficients are shown in Table 1 and risk text is included in the fact sheet. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Radiological Dispersal Device: 

Any method used to deliberately disperse 

radioactive material to create terror or 

harm. A dirty bomb is an example of an 

RDD. It is made by packaging explosives 

(like dynamite) with radioactive material 

to be dispersed when the bomb goes off. 

What Is an RDD? A radiological dispersal device (RDD) is an 

unconventional weapon that a terrorist might use to destabilize a 

community, as described at right. Although often used to 

represent a dirty bomb, the radioactivity in an RDD could also be 

distributed passively (nonexplosively), such as through spraying 

or spreading by hand. Alternately, a radiological exposure device 

(RED) might be used, which would simply involve placing a 

radioactive source in a public area to expose people passing by.

Where Would the Radioactive Material Come From? Radionuclides are used in a variety of 

industry, medicine, and scientific research applications, as illustrated by the examples below. Many of 

these are in sealed sources, used in civil engineering (in flow gauges and to test soil moisture and material 

thickness/integrity for construction), in petroleum engineering (in well logging for oil exploration), in the 

airline industry (in fuel gauges and to check welds and structural integrity), in medicine (cancer treatment, 

pacemakers, and diagnostics), in homes (smoke detectors), and to make electricity (in radiothermal 

generators or RTGs, that generate power in remote areas ranging from lighthouses to outer space). 

Radioactive sources can be portable or fixed, and most are quite small, ranging from tiny brachytherapy 

needles or seeds (implanted for localized cancer treatment) to thimble-sized plugs sealed within secure 

capsules for industrial gauges. Even the larger sources are fairly small; for example, the radioactive 

component of an RTG can range from the size of a roll of duct tape to a small wastebasket, although the 

outer housing can more than double the overall size. Most sources are encapsulated or sealed in housings 

of stainless steel, titanium, platinum, or other metal, and gamma emitters are encased in dense shielding 

(such as lead) to attenuate external gamma irradiation. 

Only some of the materials identified above are considered likely RDD candidates, based on portability 

coupled with relatively high levels of radioactivity. Not of concern are those with minute amounts of 

radioactivity, e.g., smoke detectors, camping lanterns, or brachytherapy needles; these would not 

constitute a dispersal or exposure issue even if thousands were collected to extract their radioactive 

material. (Key radionuclides of concern for RDDs are described in the next section.) Radioactive waste 

from the nuclear power industry or legacy weapons facilities is also considered a possible source, with its 

attractiveness depending on the specific radionuclides in that waste, their physical and chemical forms, 

and levels of radioactivity. High-activity wastes (e.g., from nuclear energy reactors) are well controlled, 

and the largest volumes of radioactive waste typically contain relatively low concentrations, so these 

materials are generally considered a secondary concern for RDDs. 

Examples of Radionuclides in Common Use 

Medicine Industry/Commerce Science 

Diagnosis Treatment Energy, Defense Testing, Production Food, Agriculture Home Research 

Tracer, flow 

(Tc-99m, 

I-131)

Gamma knife, 

blood/tissue 

sterilization 

(Cs-137, Co-60)

Commercial 

electricity 

(U, Pu)

Nondestructive test of 

structural integrity, 

radiographic imaging

(Co-60, Ir-192)

Food product 

sterilization (Co-60)

Smoke 

detector 

(Am-241)

High-energy 

physics 

(Cf-252, U-235) 

Tissue scan 

for clot, mass 

(Ga-67)

Needle, seed 

implants (Cs-137, 

Ir-192, Ra-226)

Remote power 

(Sr-90)

Density, moisture 

gauges (Am-241, 

Cs-137) 

Pest (fruit fly) 

sterilization 

(Cs-137, Co-60) 

Luminescent 

watch/clock 

dial (H-3)

Biokinetics 

(Pu, Sr-90, others)

X-ray 

(Cs-137, 

 Co-60)

Pacemaker 

(Pu-238)

Defense/weapons 

(Pu, H-3, U and 

depleted U)

Material thickness, 

flow, conveyor, level 

gauges (Am-241, 

Cs-137, Co-60, Kr-85)

Seed, spice 

sterilization 

(Cs-137, Co-60)

Gas camping 

lantern 

(Th-232) 

Biological tracer, 

protein/synthesis 

(C-14, H-3, 

N-15 P-32, S-35)

Am-americium, C-carbon, Cf-californium, Co-cobalt, Cs-cesium, Ga-gallium, H-3-tritium, I-iodine, Ir-iridium, K-potassium, 

Kr-krypton, N-nitrogen, P-phosphorous, Pu-plutonium, Ra-radium, S- sulfur, Sr-strontium, Tc- technetium, Th- thorium, U-uranium.

Radiological Dispersal Device (RDD) 
Which Radionuclides Are of Most Concern? Although dozens of radionuclides are used across 

various sealed sources 

(selected devices and 

associated sources are shown 

at right), only a small 

number are in concentrated 

amounts or are widely 

available. Nine isotopes of 

interest for RDDs are: 

• Americium-241 (Am-241)

• Californium-252 (Cf-252)

• Cesium-137 (Cs-137)

• Cobalt-60 (Co-60)

• Iridium-192 (Ir-192)

• Plutonium-238 (Pu-238)

• Polonium-210 (Po-210)

• Radium-226 (Ra-226)

• Strontium-90 (Sr-90)

Basic radiological properties 

for these nine isotopes are 

summarized below. (Note: 

radium-226 exists in nature, 

as does a small amount of polonium-210; the rest are man made.) 

 The specific activity of a radionuclide is inversely proportional to its half-life, as curies per gram (Ci/g). 

Unique to each isotope, it provides an indication of the rate at which that given radionuclide decays. Note 

that although iridium-192 has the highest specific activity among the nine, it decays to a stable isotope 

much more quickly than the others because its half-life is only 2.5 months. As a general rule of thumb, 

7 to 10 half-lives can indicate how long an isotope could be expected to remain radioactive. (Less than 

1% of the 

original amount 

remains after 

seven half-lives.) 

For example, if 

iridium-192 were 

dispersed, it 

would probably 

be gone within 

2 years, while 

levels of 

cobalt-60 would 

drop to 1 percent 

within 40 years, 

but it would take 

cesium-137 more 

than 200 years to 

be reduced to the 

same level. 

Basic Radiological Properties of Nine Key Radionuclides for RDDs 

Radiation Energy (MeV)

Isotope Half-Life

(years) 

Specific 

Activity 

(Ci/g)

Decay Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ)

Americium-241 430 3.5 α 5.5 0.052 0.033 

Californium-252 2.6 540 α (SF, EC) 5.9 0.0056 0.0012 

Cesium-137 30 88 β, IT - 0.19, 0.065 0.60 

Cobalt-60 5.3 1,100 β - 0.097 2.5 

Iridium-192 0.2 (74 d) 9,200 β, EC - 0.22 0.82 

Plutonium-238 88 17 α 5.5 0.011 0.0018 

Polonium-210 0.4 (140 d) 4,500 α 5.3 - -

Radium-226 1,600 1.0 α 4.8 0.0036 0.0067 

Strontium-90 29 140 β - 0.20, 0.94 - 

SF = spontaneous fission; IT = isomeric transition; EC = electron capture. A hyphen means not 

applicable. The radiation energies for cesium-137 include the contributions of barium-137 metastable 

(Ba-137m), and those for strontium-90 include the contributions of yttrium-90. 

Source for medical 

teletherapy machine 

(1-in.diameter, within 

stainless steel housing) 

(Cs-137, Co-69) 

Teletherapy 

machine 

Source for industrial 

radiography device 

(like for medical unit) 

(Co-60, Ir-192)

Industrial 

radiography 

device 

Source inside capsule at 

end of flexible cable 

(Co-60, Ir-192; 

well logs also use 

Am-241, Cs-137) 

Well logging 

devices with 

source at tip 

Russian radiothermal 

generators (RTGs) 

(Sr-90)
This concept of half life is illustrated below (from the Uranium Information Center, Melbourne, 

Australia, www.uic.com.au/ral.htm): 

Three of the nine isotopes considered candidates for an RDD are strong gamma-ray emitters: Cs-137 

(from Ba-137m), Co-60, and Ir-192. These three could pose an external hazard to individuals who handle 

them (e.g., potential terrorists) if their protective shielding was removed or not used. In fact, it is precisely 

their gamma radiation that makes these three isotopes valuable for commercial and medical applications. 

Gamma emitters are used to sterilize food and equipment, irradiate tumors, nondestructively evaluate 

high-integrity welds and castings (industrial radiography), and in industrial gauges. 

A fourth, Sr-90, emits beta particles and has limited but notable commercial uses. Like alpha emitters, 

beta emitters primarily represent an internal health hazard if ingested or inhaled. The major use of Sr-90 is 

in RTGs, and many of these were produced by the former Soviet Union to generate electricity in remote 

locations for applications such as lighthouses. 

The poor accountability for Soviet-era RTGs has 

been widely publicized, and they pose a 

considerable threat because they may contain 

tens of thousands of curies of Sr-90. 

Radioactive decay has substantially reduced 

initial levels of older RTGs, but levels in newer 

units could be much higher than 10,000 Ci. 

Summary information for these four isotopes in 

typical sealed sources is highlighted at left. 

The remaining isotopes are primarily alpha emitters: Am-241, Cf-252, Po-210, Pu-238, and Ra-226. 

Alpha particles are easily shielded with only minimal amounts of material, so they do not pose a 

significant external health hazard. Rather, their significance relates to health concerns if ingested or 

inhaled. In addition, Am-241 is commonly mixed with beryllium to produce a neutron-emitting source. 

Similarly, Cf-252 emits neutrons through spontaneous fission. Neutron emitters represent both an external 

and internal health hazard. Among other applications, alpha or neutron emitters have been used in soil 

moisture/density gauges, medical pacemakers, and well logging gauges used in the petroleum industry. 

How Dispersible Would these Radionuclides Be? Dispersibility will depend on the physical and 

chemical properties of the radioactive material used in an RDD. Metallic forms would be difficult to 

disperse while a powder could be dispersed fairly readily. Common forms of radionuclides in sealed 

sources are shown on the next page. Cobalt, iridium, and polonium generally exist as solid metals and 

would not be readily dispersible. Several of the others, including americium, californium, and plutonium, 

are typically oxides that could exist as a powder. Cesium is typically found as cesium chloride, which is 

also a powder and is quite soluble in water. Radium and strontium are used in various forms; strontium 

fluoride in certain sealed sources is sintered such that it is essentially insoluble and nondispersible. Even 

considering the forms in current sources, the specific physical and chemical characteristics of radioactive 

materials that could be in an RDD is uncertain because the original material could be chemically or 

physically altered (weaponized) to enhance dispersal. If the dispersal method is explosion via a dirty 

bomb, that would also likely physically and chemically alter the materials to produce a mixture that could 

include oxides as well as nitrates (from the explosives) over a range of particle sizes. 

Selected Highlights for Typical Commercial Sources

Initial Activity (Ci) Radionuclide Typical Form Medical Industrial

Cesium-137 Cesium chloride 0.01-10 10-1,000 

Cobalt-60 Metallic cobalt or alloy 0.01-10 10-1,000 

Iridium-192 Metallic iridium 0.01-10 10-1,000 

Strontium-90 Strontium chloride, 

fluoride, titanate 1-10 1-10,000 
What Would the Response to an RDD Involve? The response to an RDD event would consist of 

several phases. The first phase would involve immediate life-saving measures, such as treating blast 

victims and evacuating areas as indicated (e.g., based on radioactivity levels). The second phase would 

involve evaluating the extent of contamination and taking measures to control further contamination and 

minimize human exposures. The last phase would involve recovery and cleanup efforts, including 

decontamination and remediation of contaminated property. 

As background on potential health effects, evidence linking radiation exposure to observable biological 

effects has only been found at relatively high doses, i.e., acute doses exceeding 25 rads (see below). (For 

context, natural background radiation translates to an average annual dose of about 0.3 rem, which is far 

below the threshold for acute effects and corresponds to a 

lifetime risk of about 1 in 100.) On average, about half of all 

cancers that can be induced by radiation are fatal; this ranges 

from about 10% for thyroid cancer to essentially 100% for 

liver cancer. An RDD would most likely result in relatively 

small radiation exposures, which overall might not 

substantially differ from an annual background dose. But in 

the unlikely event someone was highly exposed, chelation 

therapy (to enhance excretion) and other medical 

interventions could be pursued, including to limit internal 

deposition. 

The degree of decontamination for a given area would depend on conditions specific to that setting. 

Federal, state and local officials are developing emergency response plans, but a common set of 

numerical standards for RDD cleanup has not been established. Although various radiation regulations 

would be used as guides, none are directly applicable to RDD scenarios. More than 10 years ago, in its 

manual of protective actions guides (PAGs), EPA identified nonbinding recommendations for responses 

in the early and intermediate stages of a radiological emergency. However, these PAGs were developed 

for accidents at nuclear power plants and not for incidents of radiological terrorism. Specifically the 

PAGs do not contain guidance for the long-term phase, final cleanup. 

The Department of Homeland Security has formed an interagency working group that includes 

representatives from eight Federal departments and agencies to develop new guidance for cleaning up 

after any RDD attack. Recommendations are expected to include use of the EPA PAG values for 

radiation exposure in the early and intermediate stages of a radiological terrorist attack. For the cleanup 

of areas contaminated by an RDD, rather than using a single numeric guideline that would not be able to 

account for all settings, it is expected that local stakeholders and decision makers would follow a process 

to develop cleanup plans tailored to the specific characteristics of the given situation, considering 

Chemical Forms of Radioactive Materials Often Found in Sealed Sources 

Radionuclide Form 

Americium-241 Americium oxide; americium-beryllium (AmBe) neutron sources are typically compressed powders 

Californium-252 Californium oxide 

Cesium-137 Cesium chloride 

Cobalt-60 Metallic cobalt, or cobalt-nickel alloy 

Iridium-192 Metallic iridium 

Plutonium-238 Plutonium dioxide, generally pressed into a ceramic-like material 

Polonium-210 Metallic foil 

Radium-226 Radium bromide or radium chloride 

Strontium-90 Metallic strontium, strontium chloride, strontium-fluoride, strontium-titanate 

Threshold Doses for Prodromal Effects 

Dose (rads) Indicator Effects 

50 Blood count changes 

100 Nausea, vomiting, appetite loss, 

malaise, and fatigue 

200 Diarrhea or bloody diarrhea 

300 Epilation (hair loss) 

500 Erythema (skin reddening) 
optimization approaches. Thus, residual radionuclide concentrations would be expected to vary from case 

to case, and there will not likely be a uniform, generic cleanup level. 

Where Can I Find More Information about RDDs? In the last several years, a number of reports, 

studies, articles, and books have been published that discuss issues related to RDDs. An 

introduction to selected resources follows; this list is not intended to be comprehensive. 

CDC (Centers for Disease Control and Prevention), 2003, Dirty Bombs, fact sheet, U.S. Department of 

Health and Human Services, July. Available at http://www.bt.cdc.gov/radiation/pdf/dirtybombs.pdf. 

Accessed Feb. 2005. 

DOE/NRC (U.S. Department of Energy/U.S. Nuclear Regulatory Commission), 2003, Radiological 

Dispersal Devices: An Initial Study to Identify Radioactive Materials of Greatest Concern and 

Approaches to Their Tracking, Tagging, and Disposition, report prepared by the DOE/NRC Interagency 

Working Group on Radiological Dispersal Devices for the NRC and Secretary of Energy, May 7. 

Available at http://www.nti. org/e_research/official_docs/doe/DOE052003.pdf. Accessed Feb. 2005. 

Illustrative Case Study: 1987 Radiological Accident in Goiania, Brazil 

In September 1987, a hospital in Goiania, Brazil, moved to a new location and left its radiation cancer 

therapy unit behind. Found by scrap metal hunters, it was dismantled and the cesium chloride source 

containing 1,400 Ci of cesium-137 was removed. Pieces were distributed to family and friends, and 

several who were intrigued by the glow spread it across their skin. Eleven days later, alert hospital staff 

recognized symptoms of acute radiation syndrome in a number of victims. 

The ensuing panic caused more than 112,000 people – 10% of the population – to request radiation 

surveys to determine whether they had been exposed. At a makeshift facility in the city’s Olympic 

Stadium, 250 people were found to be contaminated. 28 had sustained radiation-induced skin injuries 

(burns), while 50 had ingested cesium, so for them the internal deposition translated to an increased risk 

of cancer over their lifetime. Tragically, 2 men, 1 woman, and 1 child died from acute radiation 

exposure to the very high levels of gamma radiation from the breached source. 

In addition to the human toll, contamination had been tracked over roughly 40 city blocks. Of the 

85 homes found to be significantly contaminated, 41 were evacuated and 7 were demolished. It was 

also discovered that through routine travels, within that short time people had cross-contaminated 

houses nearly 100 miles away. Cleanup generated 3,500 m3

 radioactive waste at a cost of $20 million. 

The impacts of this incident continued beyond the health and physical damage to profound 

psychological effects including fear and depression for a large fraction of the city’s inhabitants. 

Further, frightened by the specter of radioactive contamination, neighboring provinces isolated Goiania 

and boycotted its products. The price of their manufactured goods dropped 40% and stayed low for 

more than a month. Tourism, a primary industry, collapsed and recent population gains were reversed 

by business regression. Total economic losses were estimated at hundreds of millions of dollars. A key 

lesson learned from this incident is the importance of enhancing the broader understanding of radiation. 

This fact sheet is intended to help support that objective. 

(For additional information see: International Atomic Energy Agency (IAEA), 1988, The Radiological 

Accident in Goiania, Vienna, Austria.) 
Eraker, E., 2004, “Cleanup after a Radiological Attack, U.S. Prepares Guidance,” The Nonproliferation 

Review, pp. 167−185, fall−winter. Available at http://cns.miis.edu/pubs/npr/vol11/113/113eraker.pdf. 

Accessed Feb. 2005. 

Ferguson, C.D., and W.C. Potter, 2004, The Four Faces of Nuclear Terrorism, Center for 

Nonproliferation Studies, Monterey Institute of International Studies, Monterey, CA. Information 

available at http://cns.miis.edu/pubs/books/4faces.htm. Accessed Feb. 2005. 

Ferguson, C.D., et al., 2003, Commercial Radioactive Sources: Surveying the Security Risks, Occasional 

Paper No. 11, Center for Nonproliferation Studies, Monterey Institute of International Studies, Monterey, 

CA, Jan. Available at http://cns.miis.edu/pubs/opapers/op11/op11.pdf. Accessed Feb. 2005. 

GAO (U.S. Government Accountability Office), 2003a, DOE Action Needed to Ensure Continued 

Recovery of Unwanted Sealed Radioactive Sources, GAO-03-483, report to the Ranking Minority 

Member, Subcommittee on Financial Management, the Budget, and International Security, Committee on 

Governmental Affairs, U.S. Senate, Apr. Available at http://www.gao.gov/highlights/d03483high.pdf. 

Accessed Feb. 2005. 

GAO, 2003b, U.S. and International Assistance Efforts to Control Sealed Radioactive Sources Need 

Strengthening, GAO-03-638, report to the Ranking Minority Member, Subcommittee on Financial 

Management, the Budget, and International Security, Committee on Governmental Affairs, U.S. Senate, 

May. Available at http://www.gao.gov/highlights/ d03638high.pdf. Accessed Feb. 2005. 

Kelly, H., 2002, Testimony of Dr. Henry Kelly, President, Federation of American Scientists, before the 

Senate Committee on Foreign Relations, March 6. Available at http://www.fas.org/ssp/docs/030602-

kellytestimony.htm. Accessed Feb. 2005. 

Lee, R., 2003, “Nuclear Smuggling: Patterns and Responses,” Parameters, US Army War College 

Quarterly XXXIII(1):95−111, spring. Available at http://carlisle-www.army-mil/usawc/Parameters/ 

03spring/ contents.htm. Accessed Feb. 2005. 

Medalia, J., 2004, Terrorist “Dirty Bombs”: A Brief Primer, CRS Report for Congress, Order Code 

RS21528, Congressional Research Service, Foreign Affairs, Defense, and Trade Division, updated Apr. 1. 

Available at http://www.fas.org/spp/starwars/crs/RS21528.pdf. Accessed Feb. 2005. 

NCRP (National Council on Radiation Protection and Measurements), 2001, Management of Terrorist 

Events Involving Radioactive Material, NCRP Report No. 138, Bethesda, MD, Oct. 24. Available at 

http://www.ncrponline.org/rpt138.html. Accessed Feb. 2005. 

NRC (U.S. Nuclear Regulatory Commission), 2003, Dirty Bombs, fact sheet, Office of Public Affairs, 

Washington, DC, March. Available at http://www.nrc.gov/reading-rm/doc-collections/fact-sheets/dirtybombs.html. Accessed Feb. 2005. 

Shea, D.A., 2004, Radiological Dispersal Devices: Select Issues in Consequence Management, CRS 

Report for Congress, Order Code RS21766, Congressional Research Service, Resources, Science, and 

Industry Division, updated Dec. 7. Available at http://www.fas.org/spp/starwars/crs/RS21766.pdf. 

Accessed Feb. 2005. 

Zimmerman, P.D., and C. Loeb, 2004, “Dirty Bombs: The Threat Revisited,” Defense Horizons 38:1−11, 

Jan. Available at http://hps.org/documents/RDD_report.pdf. Accessed Feb. 2005. 
Argonne National Laboratory, EVS Human Health Fact Sheet, March 2007

To address radionuclide releases to water (notably drinking water) or air, health-based concentrations can 

be developed to support environmental responses by using standard dose conversion factors and risk 

estimators from the U.S. Environmental Protection Agency (EPA), as presented in Federal Guidance 

Reports 11 and 13 (EPA 1988 and EPA 1999). Dose-based concentrations for 15 radionuclides that are of 

interest to various environmental programs are tabulated below. 

Drinking Water and Air Concentrations for Selected Radionuclidesa

Drinking Water Maximum 

Contaminant Level (MCL)-

Equivalents 

National Emissions Standard for 

Hazardous Air Pollutants 

(NESHAPs) Equivalents Radionuclide 

Specific 

Activity 

(Ci/g) (pCi/L) (μg/L) (pCi/m3

) (μg/m3) 

Americium-241 3.5 1.5 4.3 × 10-7 0.0031 8.8 × 10-10

Californium-252 540 5.1 9.4 × 10-9 0.0087 1.6 × 10-11

Cesium-137 88 110 1.2 × 10-6 43 4.9 × 10-7

Cobalt-60 1,100 200 1.9 × 10-7 6.3 5.7 × 10-9

Europium-154 270 570 2.1 × 10-6 4.8 1.8 × 10-8

Hydrogen-3 (tritium) 9,800 86,000 8.7 × 10-6 21,000 2.2 × 10-6

Iodine-129 0.00018 20 0.11 7.9 0.045 

Iridium-192 9,200 950 1.0 × 10-7 49 5.3 × 10-9

Plutonium-238 17 1.7 1.0 × 10-7 0.0035 2.1 × 10-10

Radium-226 1.0 5 5.0 × 10-6 0.16 1.6 × 10-7

Ruthenium-103 32,000 1,800 5.6 × 10-8 150 4.8 × 10-9

Ruthenium-106 3,400 200 5.9 × 10-8 2.9 8.4 × 10-10

Strontium-90 140 8 5.7 × 10-8 1.0 7.5 × 10-9

Technetium-99 0.017 3,700 0.22 1,400 0.079 

Uranium-238 0.00000034 10 30 0.12 0.35 

a

 Ci = curie; g = gram ; L = liter; m3 = cubic meters ; pCi = picocurie (10-12 curie); μg = microgram (10-6 gram); 

calculated values are given to two significant figures. Shading indicates the five isotopes with short-lived decay 

products: cesium-137 (barium-137m, 3-min half life); radium-226 (radon-222, 3.8-d half life, and its daughters); 

ruthenium-106 (rhodium-106, 30-s half life); strontium-90 (yttrium-90, 64-h half life); and uranium-238 

(thorium-234, 24-d half life; to protactinium-234m, 1.2-min half life). The strontium-90 MCL was established 

years ago using old dosimetry models; updated dosimetry gives 36 pCi/L. 

Because these concentrations are derived from EPA dose limits for drinking water and air, they are 

calculated using the current dose conversion factors from Federal Guidance Report 11 (EPA 1988) with 

the standard EPA intake assumptions of 2 L drinking water per day (2 L/d) for ingestion and 20 m3

 air/d 

for inhalation. In this table, the MCL (for strontium) and the MCL and NESHAP equivalents include the 

contributions of short-lived decay products. For the five radionuclides with short-lived daughters (shaded 

above), these daughters have half-lives of less than a month. This approach considers that beyond 

internal ingrowth after intake, the short-lived decay products would also be present at the same 

concentrations as the parents (i.e., in secular equilibrium) in the air or water to which someone is exposed. 

Longer-lived decay products are not included in the values above; notably the contributions of lead-210 

(which has a 22-year half-life) and its two short-lived decay products are not included in the value for 

radium-226 because it takes many years for them to occur by radionuclide ingrowth. By this approach, 

information in Table 1 can be used to assess introduced contaminants to which exposures would be finite 

Health-Based Radionuclide Concentrations in Drinking Water and Air 
rather than continuous, given that interest in discrete releases has increased within various environmental 

protection and preparedness programs. To also support analyses of long-term scenarios, such as chronic 

exposures to continuing airborne releases or to groundwater contaminated by past industrial releases, the 

contributions of lead-210 and its short-lived decay products to the radium-226 values for drinking water 

and air are provided in the following discussion. 

For radium-226 in air, incorporating the contributions of lead-210 and its two decay products lowers the 

value from 0.16 to 0.043 pCi/m3

. For drinking water, EPA has established MCLs for radium-226, 

strontium-90, and uranium. The MCL for radium-226 and radium-228 combined is 5 pCi/L, which is 

shown here as the level for radium-226. Because radium-228 originates from a different decay chain and 

has a much shorter half-life of 5.8 years, it would not be expected to coexist with introduced radium-226, 

so no adjustment is warranted for the combined value. Accounting for all decay products over time (to 

stable lead-206), the concentration for radium-226 that gives an annual dose of 4 mrem (as the 50-yr 

committed effective dose equivalent, or CEDE) would be 0.64 pCi/L; excluding the ingrowth of lead-210 

and its two short-lived decay products (which occurs over the longer term), the value would be 4.1 pCi/L. 

The MCL for strontium-90 was established many years ago based on extant dosimetry models. The value 

using updated dosimetry with the 4 mrem/yr annual dose limit is 36 pCi/L. The MCL for total uranium is 

30 μg/L. Applying the specific activity of uranium-238 to this MCL produces the value of 10 pCi/L for 

that isotope alone. Uranium-234 (and a very small amount of uranium-235) would be expected to 

accompany uranium-238 in an introduced source, so the total uranium activity would be about 20 pCi/L 

for uranium in its natural isotopic ratios at the MCL concentration of 30 μg/L. If someone drank 2 L 

every day for a year with this total uranium level, the resulting estimated dose would be 4 mrem (on the 

order of 1% of average background). For the other nine isotopes, an MCL equivalent has been calculated 

from the EPA annual dose limit for drinking water of 4 mrem assuming an ingestion rate of 2 L/d. 

Although this dose limit is specifically for radionuclides that emit beta particles and photons, it was also 

used here to derive concentrations for alpha emitters for which isotope-specific MCLs do not exist 

(americium-241, californium-252, and plutonium-238); note this aligns with the dose comparison for 

uranium above. The MCL for these isotopes would be 15 pCi/L as gross alpha, so levels estimated from 

the 4 mrem/yr dose are below this MCL – at 1.5, 5.1, and 1.7 pCi/L, respectively. 

For inhalation, EPA has identified an annual dose limit of 10 mrem for the air pathway in NESHAPs. 

Indicated as the effective dose equivalent (EDE) (in 40 CFR 61, Subpart H), the standard NESHAPs 

compliance code CAP88-PC integrates internal EDEs for 50 years, which results in the 50-year CEDE. 

The NESHAP-equivalent values in Table 1 correspond to this annual dose limit as CEDEs, assuming an 

inhalation rate of 20 m3

/d. This approach is also consistent with the CEDE basis of the MCL equivalents. 

As further context for the radium-226 decay product, radon-222 in air, the EPA limit for indoor radon is 

4 pCi/L, or 4,000 pCi/m3

. 

In addition to these dose-based calculations that use the dose conversion factors in Federal Guidance 

Report 11, radiological risk estimators (as risk per pCi) in the more recent Federal Guidance Report 13 

can also be used to back-calculate isotope-specific concentrations for a given target risk level. To 

illustrate for two isotopes, concentrations corresponding to a 10-4 risk level for cancer incidence from tap 

water ingestion for one year are 4,500 pCi/L for cesium-137 and 1,900 pCi/L for strontium-90. For 

longer exposures, these numbers are divided by the number of years to obtain a duration-specific value. 

 

References 

EPA, 1988, Limiting Values of Radionuclide Intake and Air Concentration and Dose Conversion Factors 

for Inhalation, Submersion, and Ingestion, Federal Guidance Report No. 11, EPA-520/1-88-020, Office 

of Radiation Programs, Washington, DC (September). 

EPA, 1999, Cancer Risk Coefficients for Environmental Exposure to Radionuclides, Federal Guidance 

Report No. 13, EPA 402-R-99-001, Office of Radiation and Indoor Air, Washington, DC (September). 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: As 

Atomic Number: 33

(protons in nucleus)

Atomic Weight: 75

Arsenic 

What Is It? Inorganic and organic arsenic occur naturally in the environment, 

with inorganic forms being most abundant. Inorganic arsenic is associated with 

other metals in igneous and sedimentary rocks, and it also occurs in combination 

with many other elements, especially oxygen, chlorine, and sulfur. Organic arsenic 

contains carbon and hydrogen. Both inorganic and organic forms exist naturally in 

soils, plants, animals, and humans. Most pure, inorganic arsenic compounds are 

white or colorless powders with no specific smell or taste. Because it is an 

element, arsenic does not degrade nor can it be destroyed. 

How Is It Used? Arsenic has been recognized as a poison since ancient times. In past centuries it was used to 

treat syphilis, and decades ago it was a common active ingredient in pesticides and was also a common wood 

preservative. Today, about 90% of arsenic produced is used as a wood preservative (chromated copper 

arsenate). Although organic arsenicals continue to be used as pesticides, primarily on cotton, inorganic 

compounds can no longer be used. Arsenic is also used as a feed additive for poultry and swine and in 

cattle and sheep dips to control lice and ticks. In addition, arsenic is used in alloys (primarily in leadacid batteries for automobiles) and in semiconductors and light-emitting diodes. 

What’s in the Environment? Arsenic occurs everywhere in the environment. Weathering of rock is the major 

natural source of inorganic arsenic, and it is also released by human activities. For example, arsenic is emitted as a 

fine dust when arsenic-containing ores are heated at smelters to process copper or lead. The concentration of 

arsenic in the earth’s crust ranges from 2 to 5 milligrams per kilogram (mg/kg), or parts per million 

(ppm). The mean natural soil concentration is 5 mg/kg, and it ranges from about 1 to 40 mg/kg. 

Water-soluble arsenites (the trivalent form, As III) and arsenates (the pentavalent form, As V) are 

the most common forms. Arsenites especially can be relatively mobile, with a typical concentration 

associated with soil particles estimated to be 10 to 200 times higher than in the interstitial water 

(water in the pore spaces between the soil particles). Levels in drinking water are commonly 2 to 20 parts per 

billion (ppb). Bacteria, fungi, and some plants methylate inorganic arsenic, converting it to organic compounds. 

Many methylated forms are volatile, such as dimethylarsine. Some organisms (notably in aquatic systems) can 

accumulate nontoxic, organic forms of arsenic; for example, levels of arsenobetaine in shrimp are often high. 

However, the typical ratio of the arsenic concentration in plants to that in soil is low, estimated at 0.006 (or 0.6%). 

What Happens to It in the Body? Arsenic can be taken in by eating food, drinking water, or breathing air, and 

to a limited degree via skin contact. Children, and to a lesser extent adults, can also be exposed by ingesting soil. 

When ingested, dissolved arsenic compounds are readily absorbed (80-90%) through the gastrointestinal tract and 

distributed in the blood to the liver, kidney, lung, spleen, aorta, and skin. Two processes are involved in arsenic 

metabolism: (1) oxidation/reduction reactions that interconvert arsenate and arsenite, and (2) methylation of 

arsenite to form monomethyl arsenic acid and dimethyl arsenic acid. The methylated forms are less toxic and more 

easily excreted in the urine. Most arsenic is eliminated in the urine within a week (75-90%, depending on the 

compound), especially from the liver, kidney, and spleen, while that in the skin, brain, skeleton, and especially hair 

and nails, remains somewhat longer. When arsenic is inhaled and deposited in the lungs, about 80% is absorbed 

into the bloodstream and distributes throughout the body as above. Arsenic in soil or dissolved in water does not 

readily penetrate the skin (only 3% is estimated to be absorbed), so dermal exposures are not typically a concern. 

What Are the Primary Health Effects? Depending on the amount ingested, arsenic can be beneficial (animal 

studies suggest that low levels of arsenic in the diet are essential) or adverse (high levels can be toxic). The acute 

lethal dose to humans can be about 2 to 20 mg/kg body weight per day (mg/kg-day). Ingesting high doses of 

arsenic irritates the stomach and intestines, with symptoms including nausea, vomiting, diarrhea and liver swelling. 

However, wide recognition of its toxicity makes arsenic poisoning today very rare. Ingesting small amounts over 

time produces chronic effects such as skin darkening and formation of corns, damage to peripheral nerves, 

cardiovascular system effects, hair and appetite loss, and mental disorders. Effects from inhaling arsenic dust 
Chemical Toxicity Values 

Cancer Risk Non-Cancer Effect

Inhalation UR Oral SF Oral RfD 

4.3 per 

mg/m3

1.5 per 

mg/kg-day 

0.0003

mg/kg-day 

include respiratory irritation, rhinitis, pharyngitis, laryngitis, and sometimes nasal perforation. Skin contact with 

inorganic arsenic dusts can cause dermatitis, allergic hypersensitivity, and conjunctivitis. Occupational exposure 

studies show a correlation between chronic arsenic exposure and lung cancer. 

Arsenic can also cause reproductive/developmental effects, including spontaneous 

abortions and reduced birth weights. Epidemiological studies indicate an 

association between arsenic concentrations in drinking water and increased 

incidences of skin, liver, kidney, lung, and bladder cancers. Studies also show an 

association between inhaling arsenic and lung cancer. From these sets of data, the 

U.S. Environmental Protection Agency (EPA) has classified inorganic arsenic as a 

known human carcinogen. Limited information is available on the joint toxicity of 

arsenic with other chemicals. For neurological effects, the predicted direction of 

joint toxicity of arsenic and lead is greater than additive, whereas the joint toxicity 

of these metals is predicted to be less than additive for the kidney and 

hematopoietic (blood-forming) system. The joint toxicity of arsenic and cadmium 

on the kidney, hematopoietic system, and male reproductive system is predicted to 

be less than additive. Additional information on joint toxicity is provided in the 

companion chemical mixtures fact sheet. 

What Is the Risk? The EPA has developed toxicity values (see box below) to estimate the risk of getting cancer 

or other adverse health effects as a result of inhaling or ingesting inorganic arsenic. These toxicity values have 

been developed based on studies of workers exposed to arsenic in occupational settings, workers applying arsenical 

pesticides, and populations consuming drinking water containing high concentrations of arsenic. The toxicity value 

for estimating the risk of getting cancer following oral exposure is called a slope factor (SF), and the value for 

estimating risk following inhalation exposure is called an inhalation unit risk (UR). An SF is an estimate of the 

chance that a person exposed to the chemical will get cancer 

from ingesting 1 mg/kg-day for a lifetime. The UR is an 

estimate of the chance that a person will get cancer from 

continuous exposure to a chemical in air at a unit concentration 

of 1 mg per cubic meter (m3

). Using the oral SF, the EPA 

estimates that a person would have a one-in-a-million chance 

of developing cancer if exposed daily to 2 liters (L) of drinking 

water containing 0.02 microgram (μg)/L arsenic. Similarly, 

using the inhalation UR, EPA estimates that a person would have a one-in-a-million chance of developing cancer if 

exposed daily over a lifetime to air containing 0.0002 μg/m3

. The toxicity value for non-cancer effects from oral 

exposure is called a reference dose (RfD), which is an estimate of the highest dose that can be taken in every day 

without causing adverse non-cancer health effects. To illustrate how the RfD is applied, a 150-pound (lb) person 

could safely ingest 0.02 mg arsenic every day without expecting any adverse effects (2.2 lb = 1 kg, or 1,000 g, or 

1 million mg). A toxicity value for non-cancer effects following inhalation exposure to arsenic is not available. 

What Are Current Limits for Environmental Releases and Human Exposure? To help track facility 

releases to the environment, the Superfund amendments addressing emergency planning and community right-toknow require immediate reporting of a release of 1 lb (0.454 kg) or more of any arsenic compound that occurs in a 

24-hour period, and also require normal releases to be reported annually and entered into a nationwide Toxics 

Release Inventory. For drinking water, EPA revised its maximum contaminant level of 0.05 mg/L to 0.01 mg/L in 

2001; public water systems are to comply with this limit by January 2006. For air, the Occupational Safety and 

Health Administration has established limits of 0.01 mg/m3

 for inorganic and organic arsenic compounds. 

Where Can I Find More Information? More information on arsenic can be found in the primary information 

source for this overview, the Toxicological Profile for Arsenic, prepared by the Agency for Toxic 

Substances and Disease Registry (ATSDR) and available at http://www.atsdr.cdc.gov/toxpro2.html. 

Other web-based sources of information include the ATSDR ToxFAQs 

(http://www.atsdr.cdc.gov/toxfaq.html), the EPA’s Integrated Risk Information System 

(http://www.epa.gov/iris/subst/index.html), and the National Library of Medicine Hazardous 

Substances Data Bank (http://toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB). 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Be 

Atomic Number: 4 

(protons in nucleus)

Atomic Weight: 9

Beryllium

What Is It? Beryllium is a hard, grayish metal that occurs naturally as a 

component of certain rocks, soil, coal, oil, and volcanic dust. Beryllium 

minerals have been known since ancient times as the gemstones emerald, 

aquamarine, and beryl. Compounds of beryllium are either white or colorless 

and do not have a particular smell. Because it is an element, beryllium does 

not degrade nor can it be destroyed. 

How Is It Used? Beryllium is used primarily in metal alloys (mainly with copper) for instruments, aircraft 

parts, springs, electrical connectors, and other industrial components. It is also incorporated into 

ceramics used in electrical insulators, microwave ovens, and rocket nozzles. Pure beryllium 

metal is used in missile and rocket parts, aircraft, heat shields, mirrors, and nuclear weapons. 

What’s in the Environment? The concentration of beryllium in the earth’s crust generally ranges from 

1 to 15 milligrams per kilogram (mg/kg), or parts per million (ppm). The average concentration of naturally 

occurring beryllium in U.S. soils is 0.6 ppm, and levels typically range from 0.1 to 40 ppm. Concentrations 

in sandy soil are estimated to be up to 250 times higher than in interstitial water (the water in the pore space 

between the soil particles), with much higher concentration ratios in loam and clay soils. 

Beryllium naturally enters waterways through the weathering of rocks and soils that 

contain this metal. It can also be released to surface waters from industrial waste 

discharges. Beryllium levels in drinking water range from 0.01 to 0.7 parts per billion 

(ppb). It is also naturally present in various foods, with a median concentration of 

22.5 μg/kg reported across 38 different food types, ranging from less than 0.1 μg/kg to 

2,200 μg/kg (in kidney beans). One cigarette contains about 0.5 to 0.7 μg beryllium, with 

about 5 to 10% escaping into sidestream smoke. 

The major source of environmental releases from human activities is combustion of coal and fuel oil. Air 

concentrations of beryllium are typically less than 0.0005 microgram per cubic meter (μg/m3

). The form 

usually released to the atmosphere is beryllium oxide, which eventually falls on land or water either in rain 

and snow or as dry particles. This oxide does not dissolve easily and does not move readily in soil or water, 

and it is considered unlikely to accumulate in plants and animals. The typical ratio of the beryllium 

concentration in plants to that in soil estimated to be low, at 0.0015 (or 0.15%). 

What Happens to It in the Body? Beryllium can enter the body by eating food, breathing air, or smoking 

a cigarette. Children, and to a lesser extent adults, can also be exposed to a limited extent by ingesting soil. 

The internal fate of beryllium depends on the form that enters the body. Most beryllium compounds do not 

dissolve easily and are not well absorbed (less than 1%) from the gastrointestinal tract, so they are generally 

excreted in the feces. Dermal absorption is also expected to be very low, although information is limited. 

The beryllium that is absorbed is excreted very slowly, and it tends to accumulate in bone as well as other 

tissues and organs. Dusts and fumes containing beryllium can be inhaled and deposited in the lungs. Some 

deposited particles can clear from the lungs slowly, and soluble beryllium compounds can be converted to 

less soluble compounds. The biological half-life of inhaled, soluble beryllium compounds is about 2 to 

8 weeks. Inhaled beryllium is excreted mainly in the urine. 

What Are the Primary Health Effects? Inhalation of beryllium can result in two types of respiratory 

disease, acute beryllium disease and chronic beryllium disease (also referred to as berylliosis). Both forms 

can be fatal. The acute disease usually occurs after exposure to high levels (more than 1 mg/m3

) of the 

relatively soluble forms of beryllium, with symptoms ranging from inflammation of the nasal passages to 

severe chemical pneumonia. Some people can get chronic beryllium disease from breathing low levels, 

occurring in less than 15% of those exposed to more than 0.0005 mg/m3

. This disease is a type of immume 
Chemical Toxicity Values 

Cancer Risk Non-Cancer Effect

Inhalation UR Oral RfD Inhalation RfC 

2.4 per 

mg/m3

0.002

mg/kg-day 

0.00002

mg/m3

response only observed in sensitized individuals, and it involves the formation 

of granuloma and development of fibrosis of the lung. There can be a 

protracted latency period (up to 25 years) before the onset of any symptoms 

following exposure. In contrast, ingesting beryllium has generally not been 

reported to cause effects in humans because very little is absorbed into the 

body. Contact dermatitis is the most common effect of beryllium on the skin, 

and contact with scraped or cut skin can cause rashes or ulcers. In its current 

narrative for the cancer weight of evidence, the U.S. Environmental Protection 

Agency (EPA) describes beryllium as a likely human carcinogen for the 

inhalation pathway and states that the carcinogenic potential of ingested 

beryllium cannot be determined. Under the previous 1996 cancer guidelines, 

EPA classified beryllium as a probable human carcinogen. 

What Is the Risk? The EPA has developed toxicity values (see box below) to estimate the risk of getting 

cancer or other adverse health effects as a result of inhaling or ingesting beryllium. The toxicity value for 

estimating the risk of getting cancer following inhalation exposure is called an inhalation unit risk (UR), 

which is an estimate of the chance that a person will get cancer from continuous exposure to a chemical in air 

at a unit concentration of 1 mg/m3

. The toxicity value for non-cancer effects from inhalation exposure is 

called a reference concentration (RfC), which is an 

estimate of the highest concentration in air that could be 

breathed every day without causing an adverse effect. The 

value for evaluating the possibility of non-cancer effects 

from oral exposure is the reference dose (RfD). The UR 

and RfC for beryllium are based on studies of humans 

exposed to beryllium in the workplace, and the RfD was 

developed by studying test animals given relatively high 

doses over their lifetimes, then adjusting and normalizing those results to a mg/kg-day basis for humans. To 

illustrate how the UR is applied, the EPA estimates that a person would have a one-in-a-million chance of 

developing cancer if exposed to air containing beryllium at a concentration of 0.0004 μg/m3

 every day over a 

lifetime. (A microgram is one millionth of a gram.) Using the RfD, it is estimated that 150-pound (lb) 

person could safely ingest 0.14 mg of beryllium every day without experiencing any adverse effects 

(2.2 lb = 1 kg, or 1,000 g, or 1 million mg). 

What Are Current Limits for Environmental Releases and Human Exposure? To help track 

facility releases to the environment, the Superfund amendments that address emergency planning and 

community right-to-know require immediate reporting of a release of 10 lb (4.54 kg) or more of any 

beryllium compound that occurs within a 24-hour period, and also require normal releases to be reported 

annually and entered into a nationwide Toxics Release Inventory. For drinking water supplies, the EPA has 

established a maximum beryllium level of 4 μg/liter. For workers exposed over regular work days, the 

National Institute for Occupational Safety and Health (NIOSH) and the Occupational Safety and Health 

Administration (OSHA) have established protective levels of 0.0005 and 0.002 mg/m3

 of beryllium and 

beryllium compounds, respectively. In addition, the EPA restricts industrial releases of beryllium to 10 g in 

a 24-hour period or an amount that would result in atmospheric levels of 0.01 μg/m3

 or less averaged over 

30 days. 

Where Can I Find More Information? More information on beryllium can be found in the primary 

information source for this overview, the Toxicological Profile for Beryllium, prepared by the 

Agency for Toxic Substances and Disease Registry (ATSDR) and available on the Internet at 

http://www.atsdr.cdc.gov/toxpro2.html. Other web-based sources of information include the 

ATSDR ToxFAQs (http://www.atsdr.cdc.gov/toxfaq.html), EPA’s Integrated Risk Information 

System Database (http://www.epa.gov/iris/subst/index.html), and the Hazardous Substances 

Data Bank (http://toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB). 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

 

Symbol: Cr

Atomic Number: 24 

(protons in nucleus)

Atomic Weight: 52 

Chromium 

What Is It? Chromium is found naturally in rocks, soil, plants, and animals, 

including people. It occurs in combination with other elements as chromium salts, 

some of which are soluble in water. The pure metallic form rarely occurs naturally. 

Chromium does not evaporate, but it can be present in air as particles. Because it is an 

element, chromium does not degrade nor can it be destroyed. 

How Is It Used? Chromium is used to make steel and other alloys, for chrome plating, and as an additive to limit 

corrosion. Named for its colored compounds, chromium has also been used to make dyes and pigments 

for paints, and to make bricks in furnaces, tan leather, and preserve wood. 

What’s in the Environment? Chromium is present everywhere and can be found in three forms: metal ore, 

trivalent chromium (Cr III), and hexavalent chromium (Cr VI). The trivalent form occurs naturally in many fresh 

vegetables and fruits, meat, grains, and yeast. Relatively insoluble, it is the most prevalent form in 

surface soils where oxidation processes (which convert chromium from the hexavalent to trivalent 

form) are most common. Hexavalent chromium also occurs naturally, notably in water-saturated 

(reducing) conditions, and it is an indicator of human pollution. This form is relatively soluble and 

can move more readily through soil to groundwater. The concentration of naturally occurring 

chromium in U.S. soil ranges from 1 to 2,000 parts per million (ppm), with an average of 54 ppm. 

The concentration in sandy soil particles is estimated to be 70 times higher than in interstitial water (the water in the 

pore spaces between the particles), and concentration ratios are higher (e.g., 1,500) for clay soil. The typical ratio 

of chromium in plants to chromium in soil is estimated at 0.0045 (or 0.45%). 

Chromium concentrations in air and water are very low. In air, concentrations generally range between 0.01 and 

0.03 microgram per cubic meter (μg/m3

), and drinking water levels are generally less than 2 parts per billion (ppb). 

Absent information on the specific form, a general assumption often made for chromium that has been present in 

surface soil over time is that 10 to 17% of the total is in the hexavalent form. Groundwater can be contaminated at 

certain industrial facilities where sodium dichromate solutions (hexavalent) were used to prevent corrosion in 

piping. For groundwater, it is often assumed that most is hexavalent. 

What Happens to It in the Body? Chromium can be taken in by breathing air, drinking water, or eating food. 

Children, and to a lesser extent adults, can also be exposed by ingesting soil. It can also be absorbed through the 

skin to a limited extent. Hexavalent chromium is more readily absorbed than trivalent chromium, regardless of the 

route of exposure. When air containing chromium is inhaled, chromium particles can be deposited in the lungs. 

Those deposited in the upper part of the lungs are usually coughed up and swallowed. Some that deposit deep in 

the lungs can dissolve, which allows chromium to pass through the lining of the lungs and enter the bloodstream. 

The finding of toxic effects following dermal exposure suggests that chromium is 

absorbed through the skin, although information on the percent absorbed is limited. 

Once in the bloodstream, chromium moves to all parts of the body. It is not 

metabolized, but hexavalent chromium is reduced by enzymatic reactions to trivalent 

chromium in the body. Inhaled chromium is excreted both in the urine and the feces. 

Ingestion of food is the major source of chromium exposure for most people in the 

United States. On average, adults take in an estimated 60 μg of trivalent chromium 

every day with their food. If taken in as hexavalent chromium, it is rapidly converted 

to the trivalent form after entering the stomach. When swallowed, most chromium 

leaves the body within a few days through the feces. A small amount – about 0.5% 

of Cr III and 10% of Cr VI – will pass through the lining of the intestines and enter 

the bloodstream. From there, chromium is distributed to all parts of the body. It 

then passes through the kidneys and is eliminated in the urine in a few days. The 

trivalent form in food can attach to other compounds that make it easier for 

chromium to be absorbed and enter the bloodstream from the stomach and intestines. 
What Are the Primary Health Effects? The trivalent form of chromium is an essential nutrient in our diet and 

is needed for many important functions, including lipid, protein, and fat metabolism. Even at levels above those 

required to maintain health, it exhibits very low toxicity and it is not known to cause cancer. In contrast, hexavalent 

chromium can be toxic, including causing cancer if it is inhaled; the lethal dose is estimated at about 7- milligrams 

hexavalent chromium per kilogram (mg/kg) body weight. When inhaled, hexavalent chromium can damage the 

lining of the nose and throat, and irritate the lungs as well as the gastrointestinal tract. Nasal irritation has been 

observed following acute exposure at levels less than 0.01 mg/m3

. When swallowed, it can upset the stomach and 

damage the liver and kidneys. Some people have an allergic skin reaction after touching material containing 

chromium. Hexavalent chromium is one of a small set of chemicals the U.S. Environmental Protection Agency 

(EPA) has classified as a known human carcinogen, based on studies of workers in chromium processing factories 

who developed lung cancer after chronic inhalation exposures. However, hexavalent chromium does not cause 

cancer when ingested, most likely because it is rapidly converted to the trivalent form after entering the stomach. 

Information on joint toxicity with other chemicals is provided in the companion chemical mixtures fact sheet. 

What Is the Risk? The EPA has developed toxicity values (see box below) to estimate the risk of getting cancer 

or experiencing other adverse health effects as a result of inhaling or ingesting chromium. The toxicity value for 

estimating the risk of getting cancer following inhalation exposure is called an inhalation unit risk (UR), which is 

an estimate of the chance that a person will get cancer from continuous exposure to a chemical in air at a unit 

concentration of 1 mg/m3

. A reference dose (RfD) is an estimate of the highest dose that can be taken in every day 

without causing an adverse non-cancer effect, and a reference concentration (RfC) is an estimate of the highest 

concentration in air that could be breathed every 

day without causing an adverse effect. These 

toxicity values have been developed by studying 

test animals given relatively high doses over their 

lifetimes, then adjusting those results to a mg/kgday basis for humans, or directly from studies of 

humans exposed to chromium in the workplace. 

To illustrate, a 150-pound (lb) person could safely 

ingest 100 mg (about three ounces) of trivalent 

chromium or 0.2 mg of hexavalent chromium 

every day without expecting any adverse 

effects. Using the UR, EPA estimates that a person 

would have a one-in-a-million chance of developing 

cancer if continuously exposed for a lifetime to air containing 0.00008 μg/m3

 hexavalent chromium, as particulates. 

What Are Current Limits for Environmental Releases and Human Exposures? To help track facility 

releases to the environment, the Superfund amendments that address emergency planning and community right-toknow require the immediate reporting of releases of 10 lb (4.54 kg) or more of chromic acid and 1,000 lb (454 kg) 

or more for all other regulated chromium compounds that occur within a 24-hour period, and also require normal 

releases to be reported annually and entered into a nationwide Toxics Release Inventory. The EPA has established 

a maximum contaminant level in drinking water of 0.1 ppm. The Occupational Safety and Health Administration 

has established protective levels of 0.5 and 1 mg/m3 in air for water-soluble and -insoluble trivalent compounds, and 

0.1 mg/m3

 for hexavalent compounds. 

Where Can I Find More Information? More details can be found in the primary information source for this 

overview, the Toxicological Profile for Chromium, prepared by the Agency for Toxic Substances 

and Disease Registry (ATSDR) and available at http://www.atsdr.cdc.gov/toxpro2.html. Other 

web-based sources of information include the ATSDR ToxFAQs 

(http://www.atsdr.cdc.gov/toxfaq.html), EPA’s Integrated Risk Information System 

(http://www.epa.gov/iris/subst/index.html), and the National Library of Medicine Hazardous 

Substances Data Bank (http://toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB).

Chemical Toxicity Values 

Cancer 

Effect Form Non-Cancer Effect

of 

Chromium Inhalation 

UR

Oral 

RfD Inhalation RfC

Cr III None 

established 

1.5 

mg/kg-day 

None 

established 

Cr VI 12 

per mg/m3 0.003 mg/kg-day 

0.0001 mg/m3

(as particulates) 

0.000008 mg/m3

(as chromic acid 

mists and dissolved 

Cr VI aerosols)
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Cu

Atomic Number: 29

(protons in nucleus)

Atomic Weight: 64

 Copper 

What Is It? Copper is a reddish metal that can be easily molded or shaped. 

It is an essential element for all living things, including humans. Because 

copper is an element, it does not degrade nor can it be destroyed. 

How Is It Used? Copper is extensively mined and processed in the United States. It is primarily used as 

the metal or alloy in electrical wiring, sheet metal, pipes, and other metal products, including 

pennies (which are made with copper and zinc). Copper compounds are also commonly used in 

farming to treat plant diseases like mildew. In addition, they are used to treat water to eliminate 

algae, and as a preservative for wood, leather, and fabrics. 

What’s in the Environment? Copper is very common in the environment. It occurs naturally in rock, 

soil, water, sediment, and air, as well as in plants and animals. Its average concentration in the earth’s crust 

is about 50 parts per million (ppm). Soil concentrations commonly range from 2 to 100 ppm and up to 

250 ppm, but concentrations as high as 7,000 ppm have been found near copper production facilities. High 

soil concentrations also occur in areas where waste from sewage treatment plants or mining and other copper 

industries are disposed of on the soil. Considerable data indicate that copper does not biomagnify in the food 

chain. The typical ratio of the concentration of copper in plants to that in soil is estimated at 0.25 (or 25%). 

Most plants contain less than 10 ppm copper on a dry-weight basis, and concentrations in animal foods are 

2 to 4 ppm, with dairy products containing less than 1 ppm. 

Copper is prevalent in surface water throughout the United States, with an average concentration in lakes and 

rivers of 4 parts per billion (ppb). Copper in aquatic systems can strongly attach to suspended particles and 

sediment, and it also converts to forms that are not easily absorbed after uptake by 

organisms. The average concentration in groundwater is similar to that in surface water, 

with higher levels in areas with higher naturally occurring copper in geologic materials. 

Although it can leach through somewhat acidic, sandy soil, copper is relatively immobile in 

most soils. It binds more strongly to soil particles than do other divalent cations, and it 

precipitates out of soil water by forming solids with ions such as carbonate and hydroxide. 

Concentrations of copper in soil particles can range between 35 and 100 or more times higher than 

concentrations in water in the pore spaces between the particles. Copper is also present in drinking water, 

with U.S. concentrations generally ranging from 20 to 75 ppb. However, many households with older 

plumbing have concentrations above 1,000 ppb (or 1 ppm) because copper in the pipes and brass faucets 

dissolves into the water as it sits overnight. 

What Happens to It in the Body? Exposure to copper can occur by breathing air, drinking water, eating 

food, and to a limited degree, by skin contact with soil or water. Children, and to a lesser extent adults, can 

also be exposed by ingesting soil. Absorption of copper following ingestion is normally regulated by 

homeostatic mechanisms so that the balance between copper intake and excretion is controlled. Studies of 

humans indicate that, in general, about 50% of ingested copper is absorbed into the bloodstream, although 

reported values range from 15 to 97%. Water-soluble forms of copper are more readily absorbed than 

insoluble forms. Zinc, molybdenum, and some other metals can decrease dietary copper absorption. Most 

absorbed copper is transported to the liver, with minor amounts going to bone and other tissues. Although it 

is known that copper can enter the body through the lungs or skin, information on the amount absorbed is 

limited. Most copper is excreted in the bile and feces, with only 2 to 4% of absorbed copper excreted in the 

urine. It may take several days for copper to leave the body, although studies have shown that injected 

copper has a half-life in the body of about 4 weeks. 

What Are the Primary Health Effects? In humans, copper is necessary for overall good health. People 

in the United States are estimated to take in 1 to 10 milligrams (mg) or more each day in their diets. The 

National Academy of Sciences has recommended 2 to 3 mg of copper as a safe and adequate daily intake for 
adults. Copper has been shown to have a protective effect against cadmium poisoning, and people who do 

not have enough copper in their diet can be more susceptible to adverse effects from lead. 

Drinking water with concentrations of 30 ppm or greater can cause vomiting, 

diarrhea, stomach cramps, and nausea. Large intakes can cause liver or kidney 

damage, or even death in cases of extreme exposure. People with Wilson’s 

disease have a genetic defect that results in the accumulation of copper in 

tissues, including the liver, kidney, and cornea. The excess copper in these 

people can cause damage to the kidney, liver, and brain; hemolytic anemia; and 

other effects. 

Acute exposure via inhalation of copper dusts or fumes can cause a condition 

called “metal fume fever,” characterized by chills, fever, dry throat and aching 

muscles. Long-term exposure to copper dust in air can irritate the nose, mouth, 

and eyes, and cause headaches, dizziness, nausea, and diarrhea. Adverse effects 

on the lungs of animals have been reported at concentrations of 0.1 to 3 ppm in 

air. In some people, skin contact with copper can result in an allergic reaction 

that is expressed as a skin irritation or rash. No data exist indicating that copper 

can cause cancer, and the International Agency for Research on Cancer (IARC) has stated that copper is not a 

human carcinogen. Similarly, data do not exist to indicate that copper can cause birth defects in humans. 

What Is the Risk? The U.S. Environmental Protection Agency (EPA) develops toxicity values to estimate 

the risk of non-cancer health effects from ingesting chemicals. The toxicity value used to estimate a noncancer effect is called a reference dose (RfD), which is an estimate of the highest dose that can be taken in 

every day without causing an adverse non-cancer effect. For copper, no formal RfD has been developed. 

(And as explained above, there is no indication of cancer risk from copper exposures; thus no cancer slope 

factor has been defined for this metal.) For comparison, a reference value could be derived from the 

maximum contaminant level goal (MCLG) identified for drinking water (see next section), although this 

level is not based on health effects. To illustrate how this could be applied, if it is assumed that a 150-lb 

person drinks 2 liters of water every day containing a given amount of copper, the amount corresponding to 

the MCLG would be 2.5 mg (2.2 lb = 1 kg, or 1,000 g). 

What Are Current Limits for Environmental Releases and Human Exposure? To protect against 

the aquatic organisms (such as fish), the EPA established a water quality criterion of 1 ppm for copper in 

lakes and streams. The EPA has identified an MCLG of 1.3 ppm for copper in drinking water supplies, 

which serves as an action level such that if more than 10% of tap water samples exceed this value, additional 

steps are taken to improve water quality. (Copper is regulated by a treat technique approach that considers 

corrosiveness.) For air, the Occupational Safety and Health Administration (OSHA) has established limits of 

0.2 milligram per cubic meter (mg/m3

) of copper fume (vapor generated from heating copper) and 1.0 mg/m3

of copper dusts (fine metallic copper particles) and mists in workroom air, to protect workers during an 

8-hour workday over a 40-hour work week. The National Institute for Occupational Safety and Health 

(NIOSH) recommends that average concentrations in workplace air be limited to 0.1 mg/m3

 for copper fumes 

and 1.0 mg/m3

 for copper mist. 

Where Can I Find More Information? More information on copper can be found in the primary 

information source for this overview, the Toxicological Profile for Copper, prepared by the 

Agency for Toxic Substance and Disease Registry (ATSDR) and available on the Internet at 

http://www.atsdr.cdc.gov/toxpro2.html. Other web-based sources of information include the 

ATSDR ToxFAQs (http://www.atsdr.cdc.gov/toxfaq.html) and the Hazardous Substances Data 

Bank (http://toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB). 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Lead 

What Is It? Lead is an element found naturally in rocks, soil, plants and animals. It 

typically occurs in combination with other elements as lead salts, some of which are 

soluble in water. The pure metallic form of lead is bluish-gray, but metallic lead 

rarely occurs naturally. Lead does not evaporate, but it can be present in air as 

particles. Because it is an element, lead does not degrade nor can it be destroyed. 

Several radioactive isotopes are naturally present in the environment, with lead-210 

being the isotope of most concern. (Information on radioactive isotopes is presented in the companion fact 

sheets for radium, thorium, and natural decay series.) 

How Is It Used? Lead has been used for thousands of years for a variety of purposes. Today, its major use is 

in the production of certain types of batteries. Lead is also used to make ammunition, metal products (sheet 

metal, solder, and pipes), medical equipment (radiation shields and surgical equipment), 

paints, ceramic glazes, caulking, scientific equipment (circuit boards for computers), and highprecision glass for lasers and other optical equipment. In recent years, the amount used in 

products such as paints and ceramics has decreased significantly to help minimize exposures of 

people and animals. Tetraethyl and tetramethyl lead (volatile organic forms) were used for 

many years in gasoline to increase octane rating. In the United States, this use was phased out 

during the 1980s, and lead was banned from use in gasoline for transportation in 1996. 

What’s in the Environment? Lead occurs everywhere in the environment. Concentrations in U.S. soil 

typically range from less than 10 to 30 milligrams of lead per kilogram of soil (mg/kg). However, amounts in 

the top layers vary widely and can be much higher due to human activities. For example, 

concentrations near roadways can be 30 to 2,000 mg/kg higher than natural levels due to past use 

of leaded gasoline. In air, concentrations typically range from 0.001 to 0.002 microgram per 

cubic meter (μg/m3

) in remote areas and from 0.2 to 0.4 μg/m3 in urban areas. Levels in surface 

water and groundwater typically range from 5 to 30 μg/liter. Lead is relatively immobile in soil 

but can leach to groundwater over time. Concentrations in sandy soil particles are estimated to be 270 times 

higher than in the water in pore spaces between the soil particles; it binds even more tightly to clay and loam 

soils, with concentration ratios of about 500 to more than 16,000. Reported concentrations of lead in various 

foods range from 0.002 to 0.65 mg/kg, with higher levels generally found in vegetables. The typical ratio of the 

concentration of lead in plants to that in the soil on which they grow is estimated at roughly 0.04 (or 4%). 

What Happens to It in the Body? Lead can be taken in by eating food, drinking water, or breathing air. 

Children, and to a lesser extent, adults, can also be exposed by ingesting soil. Lead can also be absorbed 

through the skin, although this is usually a less important route of exposure. If air containing lead particles is 

inhaled, particles deposited in the lungs can lead to about 90% being absorbed. Particles deposited in the upper 

parts of the lung are usually coughed up and swallowed, while those deposited deep in the lungs can dissolve, 

allowing lead to enter the bloodstream. If lead is swallowed with food, the amount absorbed into the 

bloodstream is about 10 to 15% in a typical adult; however, about 60 to 80% is absorbed in adults who have not 

eaten for a day. In general, if adults and children ingest the same amount of lead, children will absorb a higher 

percentage (about 50%). After lead enters the bloodstream, it travels to three main compartments: blood, soft 

tissue, and bone. About 95% and 73% of lead in the body is stored in bones and teeth for adults and children, 

respectively. Lead has a half-life in blood of about 1 month, whereas lead in bone has a half-life of greater than 

20 years. Inorganic lead is not metabolized in the body, but it can be conjugated with glutathione. About 75% 

of absorbed lead is excreted in urine and about 25% in feces; lead can also be excreted in breast milk. 

What Are the Primary Health Effects? Lead can affect almost every organ and system in the body, 

including the gastrointestinal tract, the hematopoietic system (blood-forming tissues), cardiovascular system, 

central and peripheral nervous systems, kidneys, immune system, and reproductive system. Young and unborn 

children can be extremely sensitive. Exposure of pregnant women to high levels of lead can result in premature 

births and smaller babies, followed by learning difficulties and reduced growth. The latter effects are also seen 

in young children exposed to lead after birth, as are effects on other organ systems. In adults, peripheral nerve 

damage has been observed at 40 to 60 micrograms of lead per deciliter of blood (μg/dL) anemia at 80 μg/dL, 

Symbol: Pb 

Atomic Number: 82

(protons in nucleus)

Atomic Weight: 207 
and encephalopathy at 100 μg/dL. Although studies indicate that lead 

acetate and lead phosphate cause cancer in laboratory animals, we do not 

know if lead can cause cancer in humans after being ingested or inhaled. 

On the basis of the animal studies, the U.S. Environmental Protection 

Agency (EPA) has classified lead as a probable human carcinogen. The 

joint toxicity of lead with other chemicals, including essential nutrients, has 

been studied more extensively than for most chemicals. Depending on the 

endpoint and chemical, the joint toxicity can be additive, higher than 

additive, or less than additive. For example, higher toxicity to the nervous 

system is predicted in combination with arsenic, cadmium, or manganese. 

In contrast, kidney toxicity is predicted to be less than additive for these 

same metal pairs. Zinc can protect against lead toxicity by reversing its 

enzyme-inhibiting effects, whereas iron deficiency appears to increase the 

gastrointestinal absorption of lead leading to increased toxicity to the 

hematopoietic system as well as other effects. Additional information is 

provided in the companion chemical mixtures fact sheet. 

What Is the Risk? Unlike most other chemicals, the potential for adverse 

health effects from inorganic lead is based on predicted or measured levels of lead in blood rather than on 

toxicity values. The EPA has developed a mathematical model (the Integrated Exposure Uptake Biokinetic 

Model, IEUBK), to predict concentrations of lead in the blood of children resulting from exposure to lead in 

soil, air, drinking water, food, and other sources. Predicted blood-lead concentrations are compared to a 

concentration of 10 μg/dL to evaluate the health risk to children. Using the IEUBK Model, the EPA estimated 

that the blood concentration of lead in children could exceed 10 μg/dL if the concentration of lead in soil at 

residences exceeds 400 mg of lead per kg of soil. Similarly, EPA used the Adult Lead Model to predict a soil 

concentration of 800 mg/kg that would be protective of the fetus of a female worker in an occupational setting; 

this concentration would also be protective for male or female adult workers. Although certain lead compounds 

have been shown to cause cancer in animals, the risk of cancer is not typically evaluated for lead because people 

are more sensitive to its non-cancer effects. (For radioactive isotopes of lead, the cancer risks are included in 

the risks for radium and thorium, as indicated in those companion fact sheets.) The organic compound tetraethyl 

lead is very toxic, and the potential for non-cancer effects from exposure to this form is evaluated using toxicity 

values, similar to the approach used for other chemicals. The toxicity value for evaluating oral exposure is the 

reference dose (RfD), which is an estimate of the highest dose that can be taken in every day without causing an 

adverse non-cancer effect. The RfD for tetraethyl lead is 0.0000001 mg/kg-day. It was developed by studying 

test animals given relatively high doses and then adjusting those results to a mg/kg-day basis for humans. A 

toxicity value for assessing non-cancer effects of tetraethyl lead following inhalation exposure is not available. 

What Are Current Limits for Environmental Releases and Human Exposures? To help track facility 

releases to the environment, the Superfund amendments addressing emergency planning and community rightto-know require that releases of eleven lead compounds to air, water, or land be reported annually and entered 

into a nationwide Toxics Release Inventory. For lead arsenate, a release of over 1 lb (0.454 kg) must be 

reported immediately, while the quantity for the other lead compounds is 10 lb (4.54 kg). The EPA requires that 

lead in air not exceed 1.5 μg/m3

 averaged over three months. The drinking water action level for lead is 

15 μg/L. The EPA has defined hazardous concentrations of lead as: 40 μg per square foot (ft2

) in dust on floors 

and 250 μg/ft2

 for interior windowsills of homes; 400 mg/kg in bare soil in children’s play areas; and 

1,200 mg/kg in bare soil in other parts of the yard. For the workplace, the Occupational Safety and Health 

Administration has established a permissible exposure limit (PEL) of 0.1 mg/m3

 for metallic lead and 

0.05 mg/m3

 for lead from soluble compounds. If the concentration of lead in the blood of a worker exceeds 

50 μg/dL, the worker is not allowed to remain in that work area. Many other regulations and recommendations 

have been developed for lead to protect public health. 

Where Can I Find More Information? More information can be found in the primary information source for 

this overview: the Toxicological Profile for Lead, prepared by the Agency for Toxic Substances 

and Disease Registry (ATSDR) and available at http://www.atsdr.cdc.gov/toxpro2.html). 

Other web-based sources of information include the ATSDR ToxFAQs 

(http://www.atsdr.cdc.gov/toxfaq.html), EPA’s Integrated Risk Information System 

(http://www.epa.gov/iris/subst/index.html), and the National Library of Medicine Hazardous 

Substances Data Bank (http://www.toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB). 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

 

Mercury 

What Is It? Mercury is an element that occurs naturally in the environment, 

usually in combination with other elements as mercury compounds or salts. 

Metallic mercury is a shiny, silver-white metal that is a liquid at room 

temperature. If heated, it enters the atmosphere as a colorless, odorless gas. 

Mercury combines with other elements to form inorganic mercury 

compounds, some of which are soluble in water. It also combines with carbon 

to form organic mercury compounds, such as methylmercury. Because it is an element, mercury does not 

degrade nor can it be destroyed. 

How Is It Used? Metallic mercury is used to produce chlorine gas and caustic soda and to extract gold from 

ores. It is used in thermometers, dental fillings, and batteries. Some inorganic mercury compounds are used in 

skin-lightening creams, as antiseptic creams and ointments, and as anti-mildew agents. 

What’s in the Environment? The most common forms of mercury that occur naturally in the environment 

are metallic mercury; the inorganic salts, mercuric sulfide and mercuric chloride; and methylmercury. 

Microorganisms and various natural processes can change these forms of mercury: metallic mercury can 

combine with other elements to form inorganic mercury compounds, and inorganic mercury compounds can 

combine with carbon to form organic mercury compounds which, in turn, can be changed to 

inorganic compounds. Methylmercury is the most common form created by these natural 

processes. This compound is of particular concern because it can bioaccumulate through the 

food chain. That is, small organisms and plants take up methylmercury as they feed, animals 

higher up the food chain eat these plants and organisms, and the process continues with 

methylmercury levels increasing up the food chain. Of particular concern is bioaccumulation 

in fish and shellfish, in which mercury concentrations in fish high in the food chain can be 

much higher than concentrations in the surrounding water or than in fish lower in the food chain. However, 

mercury does not bioconcentrate in terrestrial plants; the typical ratio of the concentration in plants to that in soil 

is estimated at 0.2 (or 20%). In the United States, the concentration of mercury in soil containing natural levels 

of mercury ranges from less than 0.02 to about 6 milligrams of mercury per kilogram of soil (mg/kg). Mercury 

leaches somewhat slowly, with the concentration in soil particles estimated to be about 10 to 100 times higher 

than in the water between the soil particles. In air, concentrations of mercury range from about 0.01 to 

0.02 microgram per cubic meter (μg/m3

). 

What Happens to It in the Body? Mercury can be taken into the body 

primarily by breathing air or eating food. How much mercury enters the 

body and what happens depends on the form of mercury and the route of 

exposure (that is, through air, food, or skin). When vapors of metallic 

mercury are inhaled, as much as 80% is absorbed. Following ingestion, 

absorption is less than 0.01% for metallic mercury, less than 10% for 

inorganic mercury, and greater than 95% for organic mercury. Mercury can 

also be absorbed through the skin, but the amount is small compared to 

breathing or swallowing it. 

Once in the body, metallic mercury and organic mercury compounds easily 

reach most tissues, including the brain; mercury also reaches the fetus of a 

pregnant woman. Mercury accumulates in the kidneys. In the brain, metallic 

mercury and methylmercury can be converted to an inorganic form that is 

then trapped in the brain. Mercury tends to stay in the body for weeks or 

months; elemental mercury and mercury vapor have a half-life of 35 to 

90 days, mercury salts have a half-life of about 40 days, and methylmercury 

has a half-life of about 90 days. Eventually, mercury leaves the body 

through the urine and feces, but the ratio is dependent on the exposure 

circumstances. For example, urinary excretion accounts for 58% of the body burden following long-term 

exposure to inorganic mercury compared to only 13% of the body burden from short-term exposure. 

Symbol: Hg 

Atomic Number: 80 

(protons in nucleus)

Atomic Weight: 201 
What Are the Primary Health Effects? Exposure to high levels of mercury can damage the brain, kidneys, 

and developing fetus. The nervous system is very sensitive to all forms of mercury, although the brain is most 

sensitive to metallic mercury and methylmercury because they enter the brain more easily than inorganic 

mercury salts. Exposure can cause tremors; memory loss; and changes in personality, vision, and hearing. (As a 

note, the Mad Hatter in the book Alice in Wonderland displays personality changes noticed during the 1800s to 

early 1900s in people engaged in the hat-making business in which mercury was used to process leather.) 

Children and fetuses are particularly sensitive to the harmful effects of mercury on the nervous system; exposure 

can result in mild to severe brain damage, including effects on a child’s behavior and ability to think and learn. 

Mercury also damages the kidneys. The essential element, selenium, may be protective against the toxic effects 

of mercury. 

Although studies in animals indicate that mercuric chloride and methylmercury can cause cancer in laboratory 

animals, we do not know if they can cause cancer in humans ingesting these compounds or breathing them in 

air. On the basis of the animal studies, the Environmental Protection Agency (EPA) identifies both mercuric 

chloride and methylmercury as possible human carcinogens. 

What Is the Risk? The EPA has developed toxicity values to estimate the risk of developing non-cancer 

effects as a result of ingesting or inhaling mercury (see box below). The toxicity value for estimating noncancer effects from oral exposure is called a reference dose (RfD), which is an estimate of the highest dose that 

can be taken in every day without causing an 

adverse non-cancer effect. The toxicity value for 

inhalation exposure is the reference concentration 

(RfC), which is an estimate of the highest 

concentration in air that can be breathed every day 

without causing an adverse effect. These toxicity 

values have been developed based on studies of 

workers exposed to mercury in occupational settings 

(metallic mercury), epidemiological studies of 

mothers and infants (methylmercury), and by 

studying test animals given relatively high doses of 

mercury over their lifetimes, then adjusting and 

normalizing those results to a mg/kg-day basis for humans (mercuric chloride and phenylmercuric acetate). To 

illustrate how the RfD may be applied, a 150-pound (lb) person could safely ingest 0.001 mg of methylmercury 

every day without expecting any adverse effects (2.2 lb = 1 kg, 1,000 g, or 1 million milligrams). 

What Are Current Limits for Environmental Releases and Human Exposures? To help track facility 

releases to the environment, the Superfund amendments addressing emergency planning and community rightto-know require that releases of certain chemicals to air, water, or land be reported immediately if they exceed 

specified levels, and also require normal releases to be reported annually and entered into a nationwide Toxics 

Release Inventory. Several mercury compounds are regulated under those amendments. Immediately 

reportable quantities are 1 lb (0.454 kg) for mercury and mercuric cyanide; 10 lb (4.54 kg) for mercuric nitrate, 

mercuric sulfate, mercuric thiocyanate, mercurous nitrate, and mercury fulminate; and 100 lb (45.4 kg) for 

phenylmercuric acetate. The EPA has established a limit of 0.002 milligram per liter (mg/L) for mercury in 

drinking water. (Mercury is not regulated under the Clean Air Act.) For air in the workplace, the Occupational 

Safety and Health Administration has set a permissible exposure limit (PEL) for mercury vapor of 0.1 milligram 

per cubic meter (mg/m3

) as a ceiling limit, and 0.01 mg/m3 for organo (alkyl) mercury compounds as an 8-hour 

time-weighted average. The Food and Drug Administration has set an action level in fish and shellfish of 1 mg 

of methylmercury per kg of fish. 

Where Can I Find More Information? More information on mercury can be found in the primary 

information source for this overview, the Toxicological Profile for Mercury, prepared by the 

Agency for Toxic Substances and Disease Registry and available on the Internet at 

http://www.atsdr.cdc.gov/toxpro2.html. Other web-based sources of information include the 

ATSDR ToxFAQs (http://www.atsdr.cdc.gov/toxfaq.html), the EPA’s Integrated Risk Information 

System (http://www.epa.gov/iris/subst/index.html), and the National Library of Medicine 

Hazardous Substances Data Bank (http://toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB). 

Chemical Toxicity Values 

Non-Cancer Effect 

Form of Mercury Oral RfD Inhalation RfC 

Metallic mercury None established 0.0003 mg/m3

Mercuric chloride 0.0003 mg/kg-day None established 

Methylmercury 0.0001 mg/kg-day None established 

Phenylmercuric 

 acetate 

0.00008 mg/kg-day None established 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: Zn

Atomic Number: 30 

(protons in nucleus) 

Atomic Weight: 65 

Zinc 

What Is It? Zinc is one of the most common elements in the earth’s crust. It is 

also an essential element for all living things. Pure zinc is a bluish-white, shiny 

metal. Powdered zinc is explosive and can burst into flames if stored in a damp 

place. Because it is an element, zinc does not degrade nor can it be destroyed. 

How Is It Used? Zinc has many commercial and industrial uses. Metallic zinc is used to coat iron and other 

metals to prevent rust, and it is also used in dry cell batteries. Zinc is mixed with other metals to form alloys 

such as brass and bronze, and pennies are made from a copper-zinc alloy. Zinc is also combined 

with other elements such as chlorine, oxygen, and sulfur to form zinc compounds used to make 

white paints, ceramics, rubber, wood preservatives, dyes, and fertilizers. Zinc compounds are also 

used in the drug industry as ingredients in common products like sun blocks, diaper rash ointments, 

deodorants, athlete’s foot preparations, acne and poison ivy preparations, and anti-dandruff shampoos. 

What’s in the Environment? Zinc is found throughout the environment in air, soil, and water, and it is 

present in all foods. It can be released by natural processes, but most results from human 

activities. Releases to air, water, and soil are common in areas where ores are mined, 

processed, and smelted for zinc. Because cadmium and lead are commonly present in zinccontaining ores, they are also typically released during these processes and so areas are often 

jointly contaminated. Zinc can be releases to the atmosphere during the production of steel and 

burning of coal or waste. Surface water can be impacted by discharges of metal manufacturing 

and chemical industry wastes, and also by run-off following precipitation on soils high in zinc, either due to the 

natural setting or human applications, including use of zinc fertilizer on agricultural soils. 

The average concentration of zinc in air (as fine dust particles) is typically less than 1 microgram per cubic 

meter (μg/m3

), although concentrations of 5 μg/m3 have been measured near industrial sources. In lakes and 

rivers, some zinc remains dissolved in water or as fine suspended particles, while other zinc settles to the bottom 

in association with heavier particles. Average concentrations range from 0.02 to 0.05 milligram per liter (mg/L) 

in surface water and 0.01 to 0.1 mg/L in drinking water. Levels U.S. soils typically range from 10 to 

300 mg/kilogram (kg), with an average concentration of about 50 mg/kg. Zinc generally remains in the upper 

layers bound to soil particles, but it can leach to groundwater depending on soil characteristics, moving more 

readily in sandy soil. Concentrations of zinc in sandy soil particles are about 200 times higher than in the water 

between the soil particles, and concentration ratios are even higher (over 1,000) in both loam and clay soils. 

Some fish may accumulate zinc, but it does not build up in plants. The typical ratio of the concentration in 

plants to that in soil is estimated at 0.9 (or 90%). Zinc has been measured in food at concentrations ranging 

from 2 parts per million (ppm) in leafy green vegetables to 29 ppm in meat, 

fish, and poultry. On average, people ingest 7 to 163 mg of zinc every day. 

What Happens to It in the Body? Zinc is one of the most abundant trace 

elements in the human body. It is typically taken in by ingestion of food and 

water, although it can also enter the lungs by inhaling air, including that 

contaminated with zinc dust or fumes from smelting or welding activities. The 

amount of zinc that can pass directly through the skin is very small. 

Absorption of zinc into the bloodstream following ingestion is normally 

regulated by homeostatic mechanisms, so the balance between zinc intake and 

excretion is controlled. Absorption from the gastrointestinal tract is 20 to 30% 

in people with diets containing adequate levels of zinc, but it can reach 80% in 

those with low levels of zinc in their diets or body tissues. Zinc is normally 

excreted in the urine and feces, with small amounts excreted in sweat. About 

90% of what is retained is found in muscles and bones. 
What Are the Primary Health Effects? Zinc is an essential element in our diet, but too little or too much can be 

harmful. Without enough dietary intake, people can experience a loss of appetite, decreased sense of taste and 

smell, decreased immune function, slow healing of wounds, and skin sores. Too little zinc can also result in 

poorly developed sex organs and retarded growth in young men. If pregnant women do not have enough zinc, 

babies might have growth retardation. Harmful effects from too much zinc generally begin at levels from 10 to 

15 times higher than the recommended dietary allowances of 5, 12, and 15 mg per day for infants, women, and 

men, respectively. Eating large amounts of zinc can cause stomach cramps, nausea, and vomiting. Taking in 

large amounts of zinc over an extended period can cause anemia, damage the pancreas, and lower the levels of 

high-density lipoprotein cholesterol (the good form of cholesterol). 

Breathing dust or fumes containing large amounts of zinc can cause a short-term disease called metal fume 

fever. This disease is an immune response affecting the lungs and body temperature. It is not known if there are 

health effects from breathing lower levels of zinc over long periods of time. It is also not known if high levels 

of zinc affect human reproduction or cause birth defects. However, infertility, low birth weight, and skin 

irritation have been observed in laboratory animals such as rats, guinea pigs, mice, and rabbits given high doses 

of zinc. 

The U.S. Environmental Protection Agency (EPA) has stated that adequate information to evaluate the 

carcinogenicity of zinc is not available. However, no studies exist that indicate zinc causes cancer in humans. 

Zinc deficiency may increase the toxic effects of arsenic, copper, cadmium and lead; thus an adequate amount of 

zinc can be considered protective against the toxicity of these elements. However, too much zinc can increase 

the absorption of lead, which has been shown to have an additive effect on the hematological effects of zinc. 

What Is the Risk? The EPA has developed a toxicity value called a reference dose (RfD) (see box below) to 

estimate the risk of adverse health effects as a result of ingesting zinc. An RfD is an estimate of the highest dose 

that can be taken in every day without causing an adverse non-cancer effect. This toxicity value has been 

developed from clinical studies in humans given dietary supplements of 

zinc. To illustrate how the RfD is applied, a 150-pound (lb) person could 

safely ingest 21 mg of zinc every day without expecting any adverse effects 

(2.2 lbs = 1 kg, or 1,000 g, or 1 million mg). 

What Are the Current Limits for Environmental Releases and Human Exposure? To help track 

facility releases to the environment, the Superfund amendments that address emergency planning and 

community right-to-know require the immediate reporting of a release of 1,000 lb (454 kg) or more of any zinc 

compound that occurs within 24 hours, and also require normal releases to be reported annually and entered into 

a nationwide Toxics Release Inventory. For drinking water, EPA has established a maximum contaminant level 

of 5 ppm for zinc and zinc compounds based on taste (not toxicity). The Occupational Safety and Health 

Administration (OSHA) has established a protective level of 1 milligram per cubic meter (mg/m3

) for zinc 

chloride fumes and a level of 5 mg/m3

 for zinc oxide fumes during an 8-hour workday over a 40-hour 

workweek. The National Institute for Occupational Safety and Health (NIOSH) has established the same 

standards for zinc and zinc chloride fumes for up to a 10-hour workday over a 40-hour workweek. 

Where Can I Find More Information? More information on zinc can be found in the primary information 

source for this overview, the Toxicological Profile for Zinc prepared by the Agency for Toxic 

Substances and Disease Registry (ATSDR) and available on the Internet at 

http://www.atsdr.cdc.gov/toxpro2.html. Other web-based sources of information include the 

ATSDR ToxFAQs (http://www.atsdr.cdc.gov/toxfaq.html), EPA’s Integrated Risk Information 

System (http://www.epa.gov/iris/subst/index.html), and the Hazardous Substances Data Bank 

(http://toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB). 

Chemical Toxicity Value 

Non-Cancer Effect: Oral RfD

0.3 mg/kg-day 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

 

What Is It? Carbon tetrachloride, CCl4, is a nonflammable man-made chemical 

that is produced as a liquid but evaporates easily in the environment and is 

commonly found as a gas. This colorless liquid is slightly soluble in water and has 

a sweet odor; most people can smell carbon tetrachloride at concentrations of about 

140 to 580 parts per million (ppm). An organic compound, carbon tetrachloride 

stays in air for a long time, with a half-life of 30 to 100 years. (The chemical halflife is the time it takes half the initial amount to be broken down, e.g., by photolysis 

or a photochemical reaction). When heated to very high temperatures, it 

decomposes to toxic phosgene and hydrogen chloride fumes. 

How Is It Used? Carbon tetrachloride was widely used for decades as a cleaning fluid and solvent, including for 

degreasing equipment and machinery parts at facilities such as the Hanford Site. It was also 

used at the Hanford Site in the refining process during the separation of plutonium. In the past, 

carbon tetrachloride was commonly used in the dry cleaning industry and in homes as a spot 

remover and was also used to make refrigerator fluids and propellants for aerosol cans. It was 

also used in agriculture through the mid-1980s as a fumigant to kill insects in grain. Although 

carbon tetrachloride is still used in some industrial applications, its production is being phased 

out in this country and in many others because of concerns about its effects on the earth’s 

ozone layer. 

What’s in the Environment? Carbon tetrachloride is widespread in the environment due to its extensive past use 

and persistence. Because it is slightly soluble in water and evaporates quickly from surface water and soil, most is 

present in air. Unlike many volatile compounds, carbon tetrachloride is quite stable in air. It does 

not readily dissociate in the lower atmosphere, nor is it easily washed out by rainfall. Only when it 

reaches the upper atmosphere (above 20 kilometers) does photodissociation become important. 

There, CCl4 reacts with ultraviolet light, producing trichloromethyl radicals and chlorine atoms. 

These two photodegradation products are key catalytic species in reaction pathways that destroy the 

ozone layer. Worldwide, its concentration in air is about 0.1 part per billion (ppb), while urban 

levels are higher at 0.2 to 0.6 ppb; reported levels in U.S. cities range from 0.14 to 0.3 ppb. This compound is also 

often found in air inside buildings (from internal materials). 

Carbon tetrachloride is present in many drinking water supplies, usually at less than 0.5 ppb. In groundwater, it can 

undergo reductive dechlorination in the presence of either free sulfide and ferrous ions or naturally occurring 

minerals that provide these ions. Microbial degradation has been shown to occur with estimated half-lives (in 

laboratory studies) of 6-12 months under aerobic conditions and 7-28 days under anaerobic conditions, although 

data are limited. In soil and sediment, carbon tetrachloride attaches to organic matter, preferring this phase 

100 times more than water. It does not appear to accumulate in plants or animals, including fresh- and salt-water 

organisms. 

What Happens to It in the Body? Carbon tetrachloride can enter the body when 

someone breathes air or ingests water or food containing the chemical, and it can 

also be easily absorbed through the skin. When carbon tetrachloride is inhaled or 

ingested (and possibly following absorption through the skin), much of it leaves the 

body within an hour or two in exhaled air. Of the initial amount breathed in, 30 to 

60% can be absorbed across the lungs and retained in the body, notably in fat. The 

rest is removed, primarily in the feces. 

Of the initial amount ingested, most (85% or more) is quickly absorbed into the 

bloodstream from the gastrointestinal tract. About 80 to 85% is then exhaled 

beginning about 8 minutes after ingestion (as blood circulates to the lungs). About 

4% of the amount that stays in the body is converted to carbon dioxide and then 

exhaled, while the remainder is metabolized and degraded, with a half-life in the 

body of about one day. Most of what remains accumulates in fatty tissue such as the 

liver and can take several weeks to be eliminated from the body in urine or feces, 

either as carbon tetrachloride or as degradation products such as chloroform. 

Carbon Tetrachloride

Symbol: CCl4 

 Cl 

 ⏐

 Cl ⎯ C ⎯ Cl 

 ⏐

 Cl 

Molecular Weight: 154
 

What Are the Primary Health Effects? Inhaling high concentrations (20,000 ppb or more) of carbon 

tetrachloride can affect the central nervous system, causing headache and dizziness often accompanied by nausea. If 

someone breathes air with levels ten times higher (200,000 ppb or more), the liver and kidney can be affected. In 

the liver, carbon tetrachloride causes fat to build up, making this organ swollen and tender and impairing its 

function. In the kidney, it reduces the ability to produce urine, causing the body to retain water (especially the 

lungs) and waste products to build up in the blood. Except in severe cases, these effects disappear after exposure 

stops, and the liver and kidney begin functioning normally again within a few days or weeks. 

Eating food or drinking water with high concentrations of carbon tetrachloride can also cause similar effects in the 

liver and kidney. Eating food with 2,500 ppm carbon tetrachloride can cause mild effects in most people, but they 

can be severe, even fatal, in individuals such as heavy drinkers whose liver function is already impaired. Its toxicity 

is also increased by interactions with other chemicals such as ketones (e.g., acetone). Carbon tetrachloride has been 

shown to increase the frequency of liver tumors in animals given relatively high concentrations by mouth (through a 

tube) for a long time. Although data indicate it causes liver cancer in animals, we do not know whether it can cause 

cancer in humans ingesting it in food or water. We also do not know whether it can cause cancer in animals or 

humans if it is inhaled. On the basis of the animal studies, the U.S. Environmental Protection Agency (EPA) has 

identified carbon tetrachloride as a probable human carcinogen. 

What Is the Risk? The EPA has developed toxicity values to estimate the risk of getting cancer or other adverse 

health effects as a result of inhaling or ingesting carbon tetrachloride (see box below). The toxicity value for 

estimating the risk of getting cancer is called a slope factor (SF), and the value for estimating risk following 

inhalation exposure is called an inhalation unit risk (UR). An SF is an estimate of the chance that a person exposed 

to the chemical will get cancer from ingesting one milligram per kilogram (mg/kg-day) for a lifetime. The UR is an 

estimate of the chance that a person will get cancer from continuous exposure to a chemical in air at a unit 

concentration of one milligram per cubic meter (mg/m3). The value for the non-cancer effect is called a reference 

dose (RfD), which is an estimate of the highest dose that can be taken in every day without causing an adverse noncancer effect. These toxicity values have been developed by studying test animals given relatively high doses over 

their lifetimes, then adjusting and normalizing those results to a mg/kg-day basis for humans. 

To illustrate how the RfD is applied, a 150-pound (lb) 

person could safely ingest 0.05 mg carbon tetrachloride 

every day without expecting any adverse effects 

(2.2 lb = 1 kg, or 1,000 g). In contrast to the RfD, which 

represents a “safe daily dose” (and so is compared to the 

amount an individual takes in, as a ratio), the SF is 

multiplied by the amount taken in to estimate the cancer 

risk and the UR is multiplied by the air concentration. 

Using these values, the EPA estimates that a person 

would have a one-in-a-million chance of developing cancer if they drank about two quarts of water containing 

0.3 microgram per liter (μg/L), or inhaled air containing about 0.07 microgram per cubic meter (μg/m3

) every day 

for 30 years. 

What Are Current Limits for Environmental Releases and Human Exposures? To help track facility 

releases to the environment, the Superfund amendments addressing emergency planning and community-right-toknow require the immediate reporting of releases of 10 lb (4.54 kg) or more carbon tetrachloride that occur within a 

24-hour period, and also require normal releases to be reported annually and entered into a nationwide Toxics 

Release Inventory. For drinking water supplies, EPA has established a maximum contaminant level of 5 ppb and 

recommends that the level not exceed 300 ppb for adults or 70 ppb for children for chronic exposures (those 

extending more than seven years). For air in the workplace, the Occupational Safety and Health Administration has 

identified a limit of 10,000 ppb for an 8-hour work day over a 40-hour work week. 

Where Can I Find More Information? More information can be found in the primary information source used 

to prepare this overview: the Toxicological Profile for Carbon Tetrachloride, prepared by the 

Agency for Toxic Substances and Disease Registry (ATSDR) and available at 

http://www.atsdr.cdc.gov/toxpro2.html. Other web-based sources of information include the 

ATSDR ToxFAQs (http://www.atsdr.cdc.gov/toxfaq.html), EPA’s Integrated Risk Information 

System (http://www.epa.gov/iris/subst/index.html), and the Hazardous Substances Data Bank 

(http://toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB). 

Chemical Toxicity Values 

Cancer Risk Non-Cancer Effect 

Oral SF Inhalation UR Oral RfD 

0.13 per 

mg/kg-day 

0.015 per 

mg/m3

0.0007 

mg/kg-day 

The RfD for inhaling carbon tetrachloride has been 

taken to be the same as that developed for ingestion.
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

 

Chloroform 

What Is It? Chloroform, CHCl3, is a non-flammable, colorless liquid also 

known as trichloromethane or methyltrichloride. It is slightly soluble in 

water and has a slightly sweet taste. Chloroform has a pleasant, non-irritating 

odor, and most people can smell it at concentrations of about 133 to 276 parts 

per million (ppm) in air. It decomposes in air with a half-life of about 

80 days (the chemical half-life is the time it takes half the initial amount to be 

broken down), to form phosgene and hydrogen chloride. 

How Is It Used? Chloroform is used primarily in the production of chlorodifluoromethane (a 

refrigerant in home air conditioners and supermarket freezers) and fluoropolymers. It is also 

used as an intermediate in the production of plastics, pharmaceuticals, dyes, and pesticides; as 

a heat transfer medium in fire extinguishers; and as a general solvent. In the past, chloroform 

was used extensively as an anesthetic during surgery, but that use has been discontinued. 

What’s in the Environment? Chloroform commonly enters the environment from manufacturing 

processes at industrial facilities, including certain chemical companies and paper mills. 

It is also formed as a byproduct of chlorinating drinking water and wastewater from 

sewage treatment plants. (Chlorine is added to these waters to kill bacteria.) Chloroform 

can also be released into the air during household use of water that contains it, such as 

during showering. 

Most chloroform that enters the environment eventually reaches the atmosphere. The amount of 

chloroform typically measured in U.S. air ranges from 0.02 to 0.05 part of chloroform per billion parts 

(ppb), but concentrations as high as 610 ppb have been measured in air at a municipal landfill. 

Concentrations of chloroform in public water supplies generally range from about 2 to 40 ppb, although 

concentrations above 300 ppb have been reported. Little information is available on concentrations in 

soil, surface water, and groundwater in uncontaminated areas. However, soil concentrations are expected 

to be very low given the tendency of chloroform to either volatilize or move through soil to groundwater. 

The concentration associated with soil particles has been estimated to be about 30% of that in interstitial 

water (in pore spaces between soil particles). 

What Happens to It in the Body? Chloroform can enter the body 

by breathing air, eating food, or drinking water containing the 

chemical and it can also be absorbed through the skin. When 

chloroform is inhaled, up to 75% is retained in the body and nearly 

all ingested chloroform is retained. Although chloroform is carried 

by the blood to all parts of the body, it tends to accumulate in tissues 

with high concentrations of lipids or fats. The tissues with the 

highest concentrations are adipose, brain, liver, kidney, and blood. 

Chloroform is metabolized in the body, and the major end product is 

carbon dioxide. Studies have shown that over 96% of inhaled 

chloroform is exhaled through the lungs within 8 hours, mainly as 

carbon dioxide and unchanged chloroform. Only a small amount 

leaves the body in urine and feces. 

What Are the Primary Health Effects? At high concentrations, 

chloroform depresses the central nervous system and can induce both 

narcosis and anesthesia (unconsciousness). Inhaling somewhat lower 

concentrations (about 900 ppm) for a short time can cause fatigue, 

dizziness, and headache. Exposure to relatively low levels of 

chloroform in air or water for long periods of time can damage the liver and kidneys, and skin sores have 

been reported following direct contact. Reproductive and birth defects have been reported in laboratory 

Symbol: CHCl3

Molecular Weight: 119 

Cl 

 | 

Cl —C— H 

 | 

Cl 
Chemical Toxicity Values 

Cancer Risk Non-Cancer Effect

Inhalation UR Oral RfD 

0.023 per mg/m3 0.01 mg/kg-day 

animal studies. Miscarriages in pregnant (female) rats and mice as well as abnormal sperm (males) were 

reported following exposure to between 30 and 400 ppm of chloroform in air. It is not known if 

chloroform causes reproductive effects in humans. 

Some studies suggest a relationship between the ingestion of chlorinated water and cancer of the colon 

and bladder in humans. Cancer of the liver and kidney has been observed in rats and mice exposed to 

elevated levels of chloroform in food and drinking water. The U.S. Environmental Protection Agency 

(EPA) considers chloroform likely to be carcinogenic to humans by all routes of exposure under highexposure conditions that lead to cytotoxicity (toxicity to cells) and regenerative hyperplasia (abnormal 

increase in number of cells) in susceptible tissues. However, chloroform is not likely to be carcinogenic 

to humans by any route of exposure under lower exposure conditions that do not cause cytotoxicity and 

regenerative hyperplasia. 

What Is the Risk? The EPA has developed toxicity values (see box below) to estimate the risk of 

getting cancer or other adverse health effects as a result of inhaling or ingesting chloroform. The toxicity 

value for estimating the cancer risk following inhalation exposure is called an inhalation unit risk (UR). 

The UR is an estimate of the chance that a person will get cancer from continuous exposure to a chemical 

in air at a unit concentration of 1 milligram per cubic 

meter (mg/m3

). Using the inhalation UR, the EPA 

estimates that a person would have a one-in-a-million 

chance of developing cancer if exposed daily over a 

lifetime to air containing 0.04 microgram per cubic meter 

(μg/m3

, where 1 mg = 1,000 μg). The toxicity value 

typically used for estimating the risk of getting cancer 

following oral exposure is called a slope factor (SF). However, because the carcinogenicity of 

chloroform following ingestion occurs only at exposures resulting in cytotoxicity, the EPA has 

determined that protecting against the non-cancer effects of oral exposure to chloroform is also protective 

against increased risk of cancer and so has not developed an SF. 

The toxicity value for non-cancer effects from oral exposure is called a reference dose (RfD), which is an 

estimate of the highest dose that can be taken in every day without causing an adverse non-cancer effect. 

To illustrate how the RfD is applied, a 150-pound (lb) person could safely ingest 0.68 mg of chloroform 

every day without expecting any adverse effects (2.2 lbs = 1 kg, or 1000 g, or 1 million mg). A toxicity 

value for non-cancer effects following inhalation exposure to chloroform is not available. 

What Are Current Limits for Environmental Releases and Human Exposures? To help track 

facility releases to the environment, the Superfund amendments addressing emergency planning and 

community right-to-know require the immediate reporting of a release or spill of 10 pounds or more of 

chloroform that occurs within a 24-hour period, and also requires that normal releases be reported 

annually and entered into a nationwide Toxics Release Inventory. For drinking water supplies, the EPA 

has established a maximum contaminant level of 100 ppb for trihalomethanes, a class of chemicals that 

includes chloroform. For air in the workplace, the Occupational Safety and Health Administration has set 

a ceiling limit of 50 ppm that is not to be exceeded at any time. 

Where Can I Find More Information? More information on chloroform can be found in the primary 

information source for this overview, the Toxicological Profile for Chloroform, prepared by 

the Agency for Toxic Substances and Disease Registry (ATSDR) and available on the Internet 

at http://www.atsdr.cdc.gov/toxpro2.html. Other web-based sources of information include the 

ATSDR ToxFAQs (http://www.atsdr.cdc.gov/toxfaq.html), EPA’s Integrated Risk Information 

System (http://www.epa.gov/iris/subst/index.html), and the Hazardous Substances Data Base 

(http://toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB). 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

What Are They? Polychlorinated biphenyls (PCBs) are a family of 209 chlorinated 

hydrocarbon compounds (known as congeners) with the same general chemical structure. 

In the United States, commercial mixtures of these man-made chemicals were commonly 

sold under the trade name Aroclor, including Aroclors 1248, 1254, and 1260. For most 

Aroclors, the last two digits indicate the approximate percent of chlorine (e.g., Aroclor 1254 

is 54% chlorine by weight.) The PCBs are nonflammable, colorless to light yellow oily 

liquids or waxy solids that have no odor or taste. They break down very slowly in the 

environment (with half-lives of months to years) and tend to cycle between air, water, and 

soil. (The chemical half-life is the time it takes half the initial amount to be broken down.) 

They can travel long distances in air and water and are found all over the world, including remote areas such as the Arctic. 

How Are They Used? The production of PCBs ended in the United States in 1997 due to concerns over their environmental 

persistence and possible health and ecological effects. Previously, these compounds were used in 

hundreds of commercial and industrial applications due to their chemical stability, high heat 

capacity, low flammability, and insulating properties. Prior to 1974, PCBs were used in both 

closed and open system applications, with the latter including flame-retardants, inks, adhesives, 

dyes, paints, plasticizers, and fluorescent lighting fixtures. After 1974, the use of PCBs was 

restricted to closed system applications such as coolants and lubricants in transformers, capacitors, 

and other electrical equipment. PCBs are still found in closed systems, where, for example, the 

volume of PCBs can range from only a few milliliters in small capacitors to over several thousand 

liters in large capacitors. 

What’s in the Environment? PCBs are found throughout the environment, although concentrations within the United 

States have been decreasing over time because of restrictions on their use and disposal. They have been released into air, 

water, and soil from contaminated facilities, including from incineration of PCB-containing wastes, leakage of old electrical 

equipment and improper disposition of spills. PCBs do not degrade easily, so they tend to persist in the environment as 

they cycle among air, water and soil. Their movement among these media depends on factors such as their degree of 

chlorination and climatic conditions. Low-chlorinated PCBs can volatilize from water and soil into air, with the highest 

release rates occurring in the summer due to higher temperatures. These lighter PCBs can be found far 

away from their original points of release. The low-chlorinated PCBs are also more reactive in the 

atmosphere and can be degraded to carboxylic acids in the presence of highly reactive molecules. 

Highly chlorinated PCBs can stick tightly to airborne particulates, but they do not tend to travel as far 

as lighter PCBs. PCBs can remain in the air for long periods of time, with half-lives ranging from 

months to years. In the United States, the average concentration in air ranges from about 0.02 to 

3.36 nanograms per cubic meter (1 nanogram is a billionth of gram), with the highest concentrations 

found in urban areas. Highly chlorinated PCBs also stick strongly to sediments and soils, especially those with high 

organic content. Because of their strong affinity for soil, PCBs typically do not migrate to groundwater. Concentrations in 

soil are typically less than 100 micrograms per kg (μg/kg), but concentrations can be more than 10,000 times higher in 

highly contaminated soils. PCBs are widely detected in surface waters in the low parts per trillion range. Similar to air, the 

half-lives of PCBs in soil or water ranges from months to years. Of particular concern is bioaccumulation, in which PCB 

concentrations in fish and animals high in the food chain can be much higher than concentrations in fish or animals lower in 

the food chain. Levels of PCBs in fish have been found in the parts per million (ppm) range, while levels in meat and dairy 

products are typically in the low parts per billion range. 

What Happens to Them in the Body? PCBs can be taken into the body primarily by eating contaminated food or 

breathing contaminated air, and to a more limited degree through ingestion of or dermal contact with contaminated water or 

soil. When PCBs are ingested, up to 100% can be absorbed into the blood; the amount absorbed following inhalation is not 

known. PCBs tend to accumulate in tissues with high concentrations of lipids and fats and can remain in the body for a 

long time. The tissues with the highest PCB concentrations are adipose, blood, brain, liver, and lungs. Due to the complex 

structure and number of different congeners, PCBs can undergo a variety of metabolic reactions. The rate of metabolism is 

generally lower for the more highly chlorinated congeners, but the number and position of the chlorine atoms is also 

important. These compounds are excreted in feces and urine. Some metabolites may be as harmful or more harmful than 

the unchanged congener. Reported half-lives of PCB mixtures in the body range from 2.6 to 8 years. 

Polychlorinated Biphenyls (PCBs)

General PCB Structure 

Chlorine atoms can be 

located in any of the 

numbered positions. 
 

Skin

Primary organs affected when 

PCBs are inhaled or ingested. 

Liver

What Are the Primary Health Effects? The most commonly observed effects in 

people (primarily workers) exposed to relatively high concentrations of PCBs are skin 

conditions such as chloracne and rashes and effects on liver function. Liver effects are 

more severe in people with impaired liver function, such as alcoholics. Studies in 

animals indicate that exposure to PCBs can harm the liver, thyroid, and immune and 

endocrine systems. Women exposed to large amounts of PCBs during pregnancy were 

found to have children with lower birth weights, abnormal behavior, irregular 

development of the immune system, and other effects. Because PCBs concentrate in 

breast milk, nursing infants might be at higher risk depending on exposure levels. In 

animals, PCBs have been shown to reduce conception rates and live birth rates. 

Animal studies indicate that exposure to relatively low levels over a long time (years) 

can result in liver and possibly other cancers, with suggestive evidence from human 

studies. The U.S. Environmental Protection Agency (EPA) and International Agency 

for Research on Cancer have determined that PCBs are probably carcinogenic to 

humans. 

What Is the Risk? The EPA has developed toxicity values (see box below) to estimate the risk of getting cancer or other 

adverse health effects as a result of ingesting or inhaling PCBs. The toxicity values were developed based on studies of 

animals exposed to PCB mixtures and individual PCB congeners. The toxicity value for estimating the risk of getting 

cancer following oral exposure is a slope factor (SF), and the value for estimating risk following inhalation exposure is a 

unit risk (UR). An SF is an estimate of the chance that a person exposed to a chemical will get cancer from ingesting 

1 milligram per kilogram per day (mg/kg-day) for a lifetime, and the UR is an estimate of the chance that a person will get 

cancer from continuous exposure to a chemical in air at a concentration of 1 mg/m3

. Because PCBs are mixtures with 

different toxicities that are 

also subject to different 

environmental processes 

(e.g., volatilization and 

degradation), EPA 

developed a range of SFs, 

corresponding to high, low, 

and lowest risk and 

persistence conditions. The criteria for selection of the appropriate SF depend on the available information and exposure 

context. For example, the high value is used to evaluate food chain exposure and soil ingestion, and the low value is 

typically used to evaluate inhalation of volatile congeners. To illustrate how the UR might be applied, using the low risk 

and persistence UR, a person would have a one-in-a-million chance of developing cancer if exposed daily over a lifetime to 

air containing 0.01 μg/m3

 of PCBs. The toxicity value for estimating non-cancer effects from oral exposure is called a 

reference dose (RfD), which is an estimate of the highest dose that can be taken in every day without causing adverse noncancer health effects. Using the RfD, a 150-pound (lb) person could safely ingest 0.001 mg of Aroclor 1254 or 0.0005 mg 

of Aroclor 1016 every day without expecting any adverse effects. The EPA is developing further guidance for assessing 

PCBs at contaminated sites. 

What Are Current Limits for Environmental Releases and Human Exposures? To help track facility releases to 

the environment, the Superfund amendments that address emergency planning and community right-to-know require 

immediate reporting of a release of 1 lb (0.45 kg) or more of PCBs to air, water, or land, and also require that normal 

releases be reported annually and entered into a nationwide Toxics Release Inventory. For drinking water supplies, the 

EPA has established a maximum contaminant level of 0.5 ppb, with a goal of zero. For air in the workplace, the 

Occupational Safety and Health Administration has identified limits of 1 mg/m3

 and 0.5 mg/m3 for PCBs containing 42% 

and 54% chlorine, respectively, for an 8-hour work day over a 40-hour work week. The Food and Drug Administration and 

Federal Insecticide, Fungicide, and Rodenticide Act have set tolerance levels for PCBs in consumer products such as milk 

and manufactured dairy products (1.5 ppm), poultry (3 ppm), eggs (0.3 ppm), animal feed for food producing livestock 

(0.2 ppm), fish and shellfish (2 ppm), infant and junior foods (0.2 ppm), and paper (10 ppm). 

Where Can I Find More Information? More information on PCBs can be found in the primary information source 

used to prepare this overview, the Toxicological Profile for PCBs, prepared by the Agency for Toxic 

Substances and Disease Registry (ATSDR) and available at http://www.atsdr.cdc.gov/toxprofiles/tp17.html. 

Other web-based sources of information include the ATSDR ToxFAQs 

(http://www.atsdr.cdc.gov/tfacts17.html), the EPA’s Integrated Risk Information System 

(http://www.epa.gov/iris/subst/index.html), and EPA’s report, PCBs: Cancer Dose Response Assessment 

and Application to Environmental Mixtures (http://www.epa.gov/opptintr/pcb/pcb.pdf). 

Chemical Toxicity Values 

Cancer Risk Non-Cancer Effect: Oral RfD

Oral SF Inhalation UR Aroclor 1254 Aroclor 1016 

0.07 per mg/kg-day (lowest)

0.4 per mg/kg-day (low)

2 per mg/kg-day (high)

0.02 per mg/m3 (lowest)

0.1 per mg/m3 (low)

0.6 per mg/m3 (high)

0.00002 

mg/kg-day 

0.00007 

mg/kg-day 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: C2H3Cl3

Molecular Weight: 133

Trichloroethane 

What Is It? Trichloroethane, C2H3Cl3, is a man-made, colorless liquid 

with a sweet, chloroform-like odor. Most people can smell it at levels of 

about 120 to 500 parts per million (ppm). Also known as TCA; 

1,1,1-trichloroethane; and methylchloroform, this compound is slightly 

soluble in water. It evaporates readily and has an atmospheric residence 

time of about six years. Trichloroethane is nonflammable, but notably at 

higher temperatures its vapors form phosgene gas and hydrogen chloride. 

How Is It Used? Trichloroethane production in the United States ended in 1996 as part of the phase-out 

of ozone-depleting chemicals. Although not prohibited, its use has declined since then as extant stores of 

trichloroethane have been depleted. Previously, trichloroethane was widely used as a 

solvent in about a fourth of all U.S. industries. It served as a cleaner and degreaser for 

various equipment, as a precursor in the manufacture of chemicals such as vinylidene 

chloride, and as a coolant during the drilling and tapping of metals. Trichloroethane was 

also used in a variety of consumer products, including drain cleaners, cleaning agents, shoe 

polish, spot removers, wallpaper and carpet glues, food packaging adhesives, photographic films, 

insecticides, rodenticides, and in many arts and crafts supplies such as printing inks and paints. In 

addition, trichloroethane was occasionally used as an aerosol propellant. 

What’s in the Environment? Trichloroethane has been released to the environment by both process 

and fugitive emissions as a result of industrial and consumer applications, although these emissions have 

decreased significantly since production ended nearly 10 years ago. Because of its high volatility, 

trichloroethane released to surface soil and water evaporates quickly, and most is found 

in air. In the atmosphere, trichloroethane is transformed over time to other chemicals 

through reactions with hydroxyl radicals. About 15% migrates to the stratosphere where 

it is converted into chlorine atoms by lower-wavelength ultraviolet light; it is these 

chlorine atoms that contribute to degradation of the ozone layer. When large amounts of 

trichloroethane are released to soil, some can move rapidly down to groundwater while 

some evaporates. The concentration associated with soil particles has been estimated to be about 80% of 

that in interstitial water (in pore spaces between soil particles). What does reach groundwater can 

evaporate and move back up through soil as a gas into the air; its predicted half-life in groundwater is 200 

to 300 days. Organisms naturally present in both soil and groundwater can break down this compound, 

and it does not accumulate in plants or animals. 

Typical concentrations of trichloroethane in air range between 0.1 and 0.9 parts per billion (ppb) outdoors 

and 0.3 and 4.4 ppb indoors, although concentrations as high as 290 ppb have been reported in air in new 

and recently renovated buildings. Higher concentrations in indoor air are likely due to the use of 

trichloroethane in consumer products, although that use continues to decline. Up to 0.01 ppm 

trichloroethane has been detected in rivers and lakes, and drinking water concentrations have been 

reported as 0.01 to 3.5 ppb, with higher levels of 0.8 to 142 ppb in municipal water systems. 

Concentrations in groundwater typically range from 0 to 18 ppb, although highly contaminated wells have 

been found to contain up to 3 grams per liter (g/L). Soil can be contaminated with trichloroethane 

through accidental spills, leaching of contaminated water from landfills or treatment and storage lagoons, 

application of insecticides and rodenticides, wet deposition, and percolation of contaminated rainwater. 

Trichloroethane has been detected in soil at levels up to 120 ppm. It has also been found in raw, 

processed, and prepared food products, with average concentrations ranging from about 0.004 ppb (ice 

from a commercial machine) to 16,000 ppb (allspice). 

Cl 

 | 

Cl —C— CH3

 | 

Cl 
What Happens to It in the Body? Inhalation is the primary route by 

which people are exposed to trichloroethane. It is rapidly absorbed 

from the lungs following inhalation and from the gastrointestinal tract 

following ingestion. Although it is also absorbed through the skin, this 

route is typically less important because trichloroethane evaporates so 

quickly. The highest concentrations of trichloroethane are found in fat, 

liver, kidney and brain. It is rapidly eliminated from the body through 

the lungs and urine, with about 90% of the amount taken in leaving 

unchanged in exhaled air, irrespective of the route of exposure (that is, 

whether inhaled, ingested, or absorbed through the skin). The 

remainder is metabolized and excreted either in the urine as 

trichloroethanol and trichloroacetic acid or exhaled from the lungs as 

carbon dioxide. 

What Are the Primary Health Effects? Trichloroethane is 

relatively less toxic than most chlorinated solvents (which accounts in 

part for its extensive use before effects on the ozone layer were 

recognized). At elevated concentrations, trichloroethane can depress the central nervous system (slowing 

functions) and cause an irregular heartbeat, loss of balance and coordination, and fluid buildup in the 

lung. In animals, exposure to lower concentrations can damage the nervous system and liver. 

Trichloroethane is one of many solvents that have been intentionally inhaled to alter mood or 

consciousness, and it has been associated with “sudden sniffing death” syndrome. Although some studies 

suggest that this compound is weakly mutagenic, those effects may be associated with another chemical 

present as a contaminant and not trichloroethane itself. The U.S. Environmental Protection Agency 

(EPA) has stated that adequate information is not available to evaluate its carcinogenicity. 

What Is the Risk? The EPA has developed toxicity values called reference doses (RfDs) to estimate the 

risk of non-cancer health effects as a result of ingesting various chemicals. An RfD is an estimate of the 

highest dose than can be taken in every day without causing an adverse health effect. Although the EPA 

previously established an oral RfD for trichloroethane, the agency withdrew that value pending further 

scientific evaluation. A reference concentration (RfC), which is the concentration in air that can be 

inhaled every day without causing non-cancer health effects, has not been established for this compound. 

What Are Current Limits for Environmental Releases and Human Exposure? To help track 

facility releases to the environment, the Superfund amendments that address emergency planning and 

community right-to-know require the immediate reporting of a release of 1,000 pounds (454 kilograms) 

or more of trichloroethane that occurs within 24 hours, and also require normal releases to be reported 

annually and entered into a nationwide Toxics Release Inventory. For drinking water, the EPA has 

established a maximum contaminant level of 0.2 milligram per liter (mg/L). For air in the workplace, the 

Occupational Safety and Health Administration (OSHA) has established a limit of 350 ppm for an 8-hour 

workday over a 40-hour week. 

Where Can I Find More Information? More information on trichloroethane can be found in the 

primary information source for this overview, the Toxicological Profile for 

1,1,1-Trichloroethane, prepared by the Agency for Toxic Substance and Disease 

Registry (ATSDR) and available at http://www.atsdr.cdc.gov/toxpro2.html. Other 

web-based sources of information include the ATSDR ToxFAQs 

(http://www.atsdr.cdc.gov/toxfaq.html), EPA’s Integrated Risk Information System 

(http://www.epa.gov/iris/subst/index.html), and the Hazardous Substances Data Bank 

(http://www.toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB). 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

What Is It? Trichloroethylene, C2HCl3, is a man-made, colorless liquid with a 

sweet odor that most people can detect at levels of about 100 parts per million 

(ppm). Also known as trichloroethene and often called TCE, this compound is 

moderately soluble in water. Trichloroethylene evaporates readily and has an 

atmospheric half-life of about a week. It is converted to phosgene gas and 

hydrogen chloride. 

How Is It Used? Trichloroethylene is an industrial solvent used primarily in metal cleaning and degreasing 

operations, including to remove grease from metal parts used to make automobiles. It is also used to make 

other chemicals such as polyvinyl chloride, and in varnishes, adhesives, paints, and lacquers. Trichloroethylene 

is present in many household products, such as spot removers, carpet cleaning fluids, typewriter 

correction fluids, and paint removers. It was once used as a dry-cleaning solvent, fumigant, and 

general anesthetic, but those uses have been discontinued. Before 1977, it was used to remove 

caffeine from coffee beans. An estimated 200 million pounds of trichloroethylene are used each 

year in the United States. 

What’s in the Environment? Trichloroethylene is mainly present in air and water, with metal degreasing 

representing a major release source for air. Although rapidly broken down in the atmosphere, primarily via 

reactions with hydroxyl radicals to produce phosgene, dichloroacetyl chloride, and formyl 

chloride, TCE is used in such large amounts that those levels continue to be replenished. 

Found throughout the world, the average concentration in U.S. air ranges from about 30 to 

460 parts per trillion (ppt), with higher levels in urban and industrial areas. 

Trichloroethylene evaporates easily from surface water, but it can remain in the subsurface, including 

groundwater, for a long time. In soil and groundwater, it can be broken down by bacteria to produce 

vinylidene chloride (a suspected human carcinogen) and vinyl chloride (a known human carcinogen). The 

concentration of trichloroethylene associated with the soil particles has been estimated to be about 60% of 

that in interstitial water (in pore spaces between soil particles). From a study by the U.S. Environmental 

Protection Agency (EPA), up to a third of drinking water sources in this country contain trichloroethylene. 

The median concentration in waters from groundwater sources is around 0.3 parts per billion (ppb). 

Trichloroethylene has also been found at very low levels in various prepared foods. It does not accumulate 

to significant levels in plants, animals, or fish. 

What Happens to It in the Body? Trichloroethylene can enter the body 

by breathing air, drinking water, or eating food containing this chemical, and 

it can also be absorbed through the skin. When trichloroethylene is inhaled, 

up to 75% can be retained in the body. A substantial fraction is also retained 

following ingestion. Some trichloroethylene is stored in fatty tissue, while 

the rest remains in the blood to circulate throughout the body, and it is rapidly 

metabolized to dichloroacetic acid, trichloroacetic acid, trichloroethanol, and 

other chemicals. Most of these metabolites leave the body in the urine. 

What Are the Primary Health Effects? Exposure to trichloroethylene 

can potentially affect a number of organs and systems, including the nervous 

system, liver, kidney, blood, cardiovascular system, immune system, and 

reproductive system. Skin contact with trichloroethylene can cause rashes, 

while inhalation or ingestion of high concentrations for a short time primarily 

Trichloroethylene

Symbol: C2HCl3

H

C1

C = C

C1

C1

Molecular Weight: 131 
 

affects the central nervous system. Inhalation or ingestion of very high concentrations can lead to loss of 

consciousness and death. At somewhat lower concentrations, people can become sleepy or dizzy or get 

headaches. Except in severe cases, the effects disappear after the exposure stops. Exposure to relatively low 

levels of trichloroethylene in air or water for a long period of time (years) can damage the liver and kidney, 

with more severe effects in people with impaired liver or kidney function, such as alcoholics. 

Abnormal skeletal development and other effects have been observed in the offspring of exposed animals; 

however, it is not known if similar effects can occur in humans. Trichloroethylene has been shown to cause 

cancer in animals, but we do not know if it causes cancer in humans. The International Agency for Research 

on Cancer has determined that trichloroethylene is not classifiable as to human carcinogenicity. The EPA is 

currently reviewing its carcinogenicity, and in a 2001 draft health assessment, characterized trichloroethylene 

as “highly likely to produce cancer in humans” based on the 1999 proposed (and now accepted) cancer 

guidelines and as a probable human carcinogen, based on the former 1986 cancer guidelines. 

What Is the Risk? The EPA has developed toxicity values to estimate the risk of getting cancer or other 

adverse health effects as a result of inhaling or ingesting trichloroethylene. These toxicity values were 

developed by studying test animals given relatively high doses over their lifetimes, then adjusting and 

normalizing those results to a milligram per kilogram per day (mg/kg-day) basis for humans and by studying 

humans exposed in occupational settings. However, at this time, EPA has withdrawn all previously 

published toxicity values as further scientific data have become available and are being evaluated. 

The toxicity value for estimating the risk of getting cancer is called a slope factor (SF) for ingestion exposure 

(also sometimes used for inhalation exposure). An SF is an estimate of the chance that a person exposed to a 

chemical will get cancer from taking in 1 milligram per kilogram of body weight per day (mg/kg-day), for a 

lifetime. The EPA identified draft oral SFs ranging from 0.02 to 0.4 per mg/kg-day in the 2001 draft health 

assessment. The toxicity value used to evaluate the non-cancer effect is called a reference dose (RfD) for 

ingestion exposure, and a reference concentration (RfC) for inhalation exposure. An RfD is an estimate of 

the highest dose that can be taken in every day without causing an adverse non-cancer effect. The RfC is the 

highest concentration in air that can be breathed everyday without causing non-cancer health effects. The 

oral RfD withdrawn from IRIS was 0.006 mg/kg-day. A draft RfD of 0.0003 mg/kg-day and an inhalation 

RfC of 0.04 milligram per cubic meter (mg/m3

) were later derived in EPA’s 2001 draft health assessment. 

What Are Current Limits for Environmental Releases and Human Exposures? To help track 

facility releases to the environment, the Superfund amendments that address emergency planning and 

community right-to-know require the immediate reporting of releases of 100 pounds (45.4 kg) or more of 

trichloroethylene that occur within a 24-hour period, and also require normal releases to be reported annually 

and entered into a nationwide Toxics Release Inventory. For drinking water supplies, the EPA has 

established a maximum contaminant level of 5 ppb for TCE with a goal of zero. For air in the workplace, the 

Occupational Safety and Health Administration has identified a limit of 100 ppm for an 8-hour work day 

over a 40-hour work week. 

Where Can I Find More Information? More information on trichloroethylene can be found in the 

primary information source used to prepare this overview, the Toxicological Profile for 

Trichloroethylene, prepared by the Agency for Toxic Substances and Disease Registry and 

available at http://www.atsdr.cdc.gov/toxpro2.html. Other web-based sources of 

information include the ATSDRs ToxFAQs (http://www.atsdr.cdc.gov/toxfaq.html), the 

EPA’s Integrated Risk Information System (http://www.epa.gov/iris/subst/index.html), and 

the National Library of Medicine Hazardous Substances Data Bank 

(http://toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB).
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

 

What Is It? Cyanide can occur naturally and can also be man-made. The 

most common forms are hydrogen cyanide (HCN) and the sodium and 

potassium salts (NaCN and KCN). Hydrogen cyanide is a colorless or paleblue liquid at room temperature (below 78o

F) and a colorless gas at higher 

temperatures; the gas is highly flammable and potentially explosive. Sodium 

and potassium cyanide are white crystals that can release hydrogen cyanide gas if they contact acids. Both 

the gas and crystals have a faint, bitter almond-like odor that is perceptible at a concentration of about 1 part 

per million (ppm). 

How Is It Used? Cyanide is used in electroplating, metallurgy, and mining. It is also used to make 

synthetic fibers, plastics, dyes, pharmaceuticals, and pesticides, including fumigants. In 

addition, cyanide serves as a chemical intermediate in various production processes. 

(Because of its acute lethality, cyanide has been used in suicide attempts, and hydrogen 

cyanide gas has been used in the past as a chemical weapon, including in gas chambers.) 

What’s in the Environment? Natural cyanides are produced by certain bacteria, fungi, 

and algae, and they are present in a number of plants and foods as cyanogenic glycosides. 

Man-made cyanides typically enter the environment from metal finishing and organic 

chemical industries. Other sources include iron and steel works, municipal waste burning, 

cyanide-containing pesticides, and vehicle exhaust. Cyanide is also present in cigarette 

smoke and smoke generated when certain synthetic materials such as plastics are burned. 

The estimated concentration of hydrogen cyanide in non-urban air is less than 0.2 parts per billion (ppb), 

where its half-life can be from 1 to 3 years. Cyanide concentrations measured in urban runoff have ranged 

from 2 to 33 micrograms/liter (μg/L), and concentrations of 5 to 1,400 μg/L have been measured in leachates 

from U.S. landfills. Most cyanide in surface water forms hydrogen cyanide and evaporates. In drinking 

water, chloramine treatment can produce cyanogen chloride, with reported concentrations ranging from 

0.45 to 0.80 μg/L. Concentrations in mainstream cigarette smoke range from 10 to 550 μg/cigarette, with 

lower concentrations (up to 111 μg/cigarette) in sidestream smoke. Cyanide is present in foods and plants at 

concentrations up to 4,000 ppm. It does not appear to accumulate in organisms, e.g., in fish. 

What Happens to It in the Body? Cyanide typically enters the body 

primarily by breathing air, but people can also be exposed by eating 

foods containing cyanogenic glycosides, breathing cigarette smoke, and 

through contact with the skin. Skin contact is common only among 

people who work in cyanide-related industries. After being ingested or 

inhaled, hydrogen cyanide is rapidly absorbed through the 

gastrointestinal tract and lungs into the bloodstream, while the 

absorption of cyanide salts tends to be slower. Approximately 58% of 

inhaled hydrogen cyanide is absorbed. 

Most cyanide that enters the body is changed to the less toxic form, 

thiocyanate. Some cyanide reacts to form vitamin B12. The primary 

route of elimination is through excretion of thiocyanate in the urine. 

Some free hydrogen cyanide is also excreted unchanged in the breath, 

saliva, sweat, and urine. Most cyanide and its products leave the body 

within 24 hours after exposure. 

What Are the Primary Health Effects? Hydrogen cyanide and its salts are among the most rapidly 

acting poisons known. Cyanide exerts its toxic effects by blocking the use of oxygen by the cells. Within 

minutes, exposure to toxic levels of cyanide causes rapid breathing, weakness, headache, vomiting, and an 

increased heart rate. Exposure to lethal concentrations causes convulsions, low blood pressure, and 

Cyanide

Symbol: CN- 

 (-C ≡ N)

Molecular Weight: 26 

Hematopoietic

System

Cardiovascular 

System

Respiratory

System

Primary organs affected when 

cyanide is inhaled or ingested.
 

respiratory failure leading to death. Because of their slower absorption, death from ingesting cyanide salts 

can be delayed, but effects still usually occur within one hour. Survivors of cyanide poisoning may develop 

heart and brain damage. Exposure to lower concentrations of cyanide over extended periods of time can also 

result in health effects. In tropical regions of Africa, ataxic neuropathy (reduced coordination due to effects 

on the nervous system), goiter, poor vision, and other disorders have been associated with chronic ingestion 

of cassava, notably the root (manioc), a dietary staple containing cyanogenic glycosides that release 

hydrogen cyanide. A decrease in the production of thyroid hormone in newborns is also reported in areas 

where cassava is a basic food. Fetal abnormalities have been observed in animal studies. There are no 

reports that cyanide can cause cancer in animals or people. 

What Is the Risk? The U.S. Environmental Protection Agency (EPA) has developed toxicity values to 

estimate the risk of developing non-cancer effects 

as a result of inhaling or ingesting cyanide 

compounds (see box at right). The toxicity value 

for non-cancer effects from oral exposure is called 

a reference dose (RfD), which is an estimate of the 

highest dose that can be taken in every day without 

causing an adverse non-cancer effect. The toxicity 

value for inhalation exposure is called a reference 

concentration (RfC), which is an estimate of the 

highest concentration in air that can be inhaled 

every day without causing an adverse non-cancer 

effect. The RfDs have been developed by studying 

test animals given relatively high doses of cyanide 

over their lifetimes, then adjusting and normalizing 

those results to a milligram per kilogram per day 

(mg/kg-day) basis for humans. The RfC for 

hydrogen cyanide is based on a study of workers 

exposed in an occupational setting. To illustrate 

how the RfD may be applied, a 150-pound (lb) 

person could safely ingest 2.7 mg of sodium cyanide every day without expecting any adverse effects 

(2.2 lb =1 kg, 1000 g, or 1 million mg). 

What Are Current Limits for Environmental Releases and Human Exposures? To help track 

facility releases to the environment, the Superfund amendments addressing emergency planning and 

community right-to-know require that releases of certain chemicals to air, water, or land be reported 

immediately if they exceed specified levels, and they also require normal releases to be reported annually and 

entered into a nationwide Toxic Release Inventory. Several cyanide compounds are regulated under these 

amendments. Immediately reportable quantities are 1 lb (0.454 kg) for potassium silver cyanide and silver 

cyanide; 10 lb (4.45 kg) for calcium cyanide, copper cyanide, cyanogen chloride, hydrogen cyanide, 

potassium cyanide, and sodium cyanide; 100 lb (45.4 kg) for cyanogen; and 5,000 lb (2,270 kg) for 

ammonium thiocyanate. The EPA has established a limit of 0.2 mg/L for cyanide in drinking water. For air 

in the workplace, the Occupational Safety and Health Administration has set a permissible exposure limit 

(PEL) over an 8-hour work day of 11 milligrams per cubic meter (mg/m3

) for hydrogen cyanide, and a PEL 

of 5 mg/m3

 for potassium cyanide and sodium cyanide (as CN-). 

Where Can I Find More Information? More information on cyanide can be found in the primary 

information source for this overview, the Toxicological Profile for Cyanide, prepared by the 

Agency for Toxic Substances and Disease Registry (ATSDR) and available at 

http://www.atsdr.cdc.gov/toxpro2.html. Other web-based sources of information include the 

ATSDR ToxFAQs (http://www.atsdr.cdc.gov/toxfaq.html), the Hazardous Substances Data 

Bank (http://www.toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB), EPA’s Integrated Risk 

Information System (http://www.epa.gov/iris/subst/index.html), and the Center for Disease 

Control and Prevention’s Emergency Preparedness and Response Chemical Agents List 

(http://www.bt.cdc.gov/agent/cyanide/index.asp). 

Chemical Toxicity Values 

Non-Cancer Effect 

Form of Cyanide Oral RfD Inhalation RfC 

Hydrogen cyanide 0.02 mg/kg-day 0.003 mg/m3

Cyanide, free 0.02 mg/kg-day None established 

Calcium cyanide 

Cyanogen 

Sodium cyanide 

0.04 mg/kg-day None established 

Chlorine cyanide 

Potassium cyanide 

Zinc cyanide 

0.05 mg/kg-day None established 

Copper cyanide 0.005 mg/kg-day None established 

Silver cyanide 0.1 mg/kg-day None established 

Potassium silver 

cyanide 0.2 mg/kg-day None established 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005

Symbol: NO3 / NO2

Molecular 

Weight: 62 / 46 

Nitrate and Nitrite 

What Are They? Nitrate and nitrite are compounds that contain a 

nitrogen atom joined to oxygen atoms, with nitrate containing three oxygen 

atoms and nitrite containing two. In nature, nitrates are readily converted 

to nitrites and vice versa. Both are anions, or ions with a negative charge. 

They tend to associate with cations, or ions with a positive charge, to 

achieve a neutral charge balance. 

How Are They Used? Nitrates are used primarily to make fertilizer, but they are also used to make 

glass and explosives. These compounds also are used in various chemical production and separation 

processes. Nitrites are manufactured mainly for use as a food preservative, and both nitrates and nitrites 

are used extensively to enhance the color and extend the shelf life of processed meats. 

What’s in the Environment? Nitrates are naturally present in soil, water, and food. In the natural 

nitrogen cycle, bacteria convert nitrogen to nitrate, which is taken up by plants and incorporated into 

tissues. Animals that eat plants use the nitrate to produce proteins. Nitrate is returned to the environment 

in animal feces, as well as through microbial degradation of plants and animals after they die. 

Microorganisms can convert nitrate or the ammonium ion (which is a nitrogen atom 

combined with four hydrogen atoms) to nitrite; this reaction occurs in the environment 

as well as within the digestive tract of humans and other animals. After bacteria 

convert (reduce) nitrate to nitrite in the environment, the nitrogen cycle is completed 

when they then convert the nitrite to nitrogen. Normally, this natural cycling process 

does not allow excessive amounts of nitrates or nitrites to accumulate in the 

environment. However, human activities have increased environmental nitrate concentrations, with 

agriculture being the major source. This includes increased use of nitrogen-containing fertilizers as well 

as concentrated livestock and poultry farming; the latter two produce millions of tons of nitrate-containing 

manure each year. Nitrate and nitrite compounds are very soluble in water and quite mobile in the 

environment. They have a high potential for entering surface water when it rains, as nitrates in applied 

fertilizers can dissolve in runoff that flows into streams or lakes; they also have a high potential for 

entering groundwater through leaching. The concentration associated with soil particles has been 

estimated to be about half the concentration in interstitial water (the water in the pore spaces between the 

soil particles). 

What Happens to It in the Body? Nitrate is a normal component of the human diet, with the average 

daily intake from all sources estimated at 75 milligrams (mg), or about 0.0026 ounce. Upon ingestion, 

about 5% of the nitrate taken in by healthy adults is converted (reduced) to nitrite by bacteria in saliva. 

Further nitrate is converted by bacteria inside the alimentary tract. Certain conditions in the stomach can 

increase the conversion of nitrate to nitrite, specifically when the pH of the gastric fluid is high enough 

(above 5) to favor the growth of nitrate-reducing bacteria. This process is of major concern for infants, 

whose gastrointestinal systems normally have a higher pH than those of adults. Nitrites in the stomach 

can react with food proteins to form N-nitroso compounds; these compounds can also be produced when 

meat containing nitrites or nitrates is cooked, particularly using high heat. While these compounds are 

carcinogenic in test animals, evidence is inconclusive regarding their potential to cause cancer (such as 

stomach cancer) in humans. 

What Are the Primary Health Effects? Nitrates themselves are relatively nontoxic. However, when 

swallowed, they are converted to nitrites that can react with hemoglobin in the blood, oxidizing its 

divalent iron to the trivalent form and creating methemoglobin. This methemoglobin cannot bind oxygen, 

which decreases the capacity of the blood to transport oxygen so less oxygen is transported from the lungs 

to the body tissues, thus causing a condition known as methemoglobinemia. 
Normal individuals have low levels (0.5 to 2%) of methemoglobin in 

their blood. When this level increases to 10%, the skin and lips can take 

on a bluish tinge (cyanosis), and levels above 25% can cause weakness 

and a rapid pulse. At levels above 50 to 60%, a person can lose 

consciousness, go into a coma, and die. Infants are much more sensitive 

than adults to nitrates/nitrites, and essentially all deaths from 

nitrate/nitrite poisoning have been in infants. Long-term exposure to 

lower levels of nitrates and nitrites can cause diuresis (an increase in the 

amount of urine, and starchy deposits and hemorrhaging of the spleen). 

What Are the Risks? The U.S. Environmental Protection Agency 

(EPA) has developed toxicity values to estimate the risk of non-cancer 

health effects from ingesting nitrates and nitrites. The toxicity value 

used to estimate a non-cancer effect following ingestion is called a 

reference dose (RfD). An RfD is an estimate of the highest dose that 

can be taken in every day without causing an adverse effect. 

The RfD for nitrate was developed considering the concentration at which methemoglobinemia was 

indicated at levels above 10% for 0- to 3-month-old infants. This was based on a daily intake of formula 

made with water containing 10 mg per liter (mg/L) of nitrate as nitrogen. (This reflects the amount of 

nitrogen within a nitrate molecule, where 1 mg nitrate as nitrogen = 4.4 mg nitrate as measured in the 

water.) Most cases of infant methemoglobinemia are associated with exposure to formula prepared with 

drinking water at nitrate-nitrogen levels at least two times higher, or exceeding 20 mg/L nitrate-nitrogen. 

For nitrite, the RfD is based on a 10-kg (22-pound [lb]) child drinking 1 liter, or about 1 quart, of water 

every day. The RfD represents a “safe daily dose” and so is 

compared to the amount an individual is estimated to take in 

every day, as a ratio. No adverse effects have been linked 

with inhaling nitrates or nitrites, so reference concentrations 

(RfCs), which are used to assess inhalation toxicity, have not 

been developed. The contribution of nitrites, and indirectly 

nitrates, to potential human carcinogenicity and the 

magnitude of the associated risk are unclear. Nitrites react with secondary amines in food to form 

nitrosamines, many of which are carcinogenic in experimental animals and exert other toxic effects. 

While the EPA has not developed slope factors (toxicity values used to estimate cancer risk) for nitrates 

or nitrites, some are available for nitrosamines.

What Are Current Limits for Environmental Releases and Human Exposure? The EPA 

requires that sodium nitrite releases of more than 100 lb (45.4 kg) and nitrate releases of more than 

10,000 lb (4,540 kg) be reported immediately and that normal releases be reported annually for inclusion 

in the nationwide Toxics Release Inventory. The limits for nitrosamines range from 1 to 10 pounds. The 

EPA primary drinking water standards for nitrate and nitrite are 10 and 1 parts per million (ppm), 

respectively. The Food and Drug Administration allows these compounds to be used as food additives as 

long as they are of food grade and are added only in the amount needed. The maximum amount of nitrite 

allowed in smoked and cured fish and meat is 200 ppm. 

Where Can I Find More Information? More information on nitrates and nitrates can be found in the 

primary information sources used to prepare this overview: (1) Consumer Fact Sheet on 

Nitrates/Nitrites, EPA Office of Groundwater and Drinking Water, available through 

http://www.epa.gov/OGWDW/dwh/c-ioc/nitrates.html; (2) the EPA’s Integrated Risk 

Information System, http://www.epa.gov/iris/subst/0076.htm (nitrate) and /0078.htm 

(nitrite); and (3) the National Library of Medicine Hazardous Substances Data Bank 

(http://toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB). 

Chemical Toxicity Values 

Non-Cancer Effect 

Oral RfD: NO3 Oral RfD: NO2

1.6 mg/kg-day 0.1 mg/kg-day 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Why Are Mixtures an Issue? Every day we are exposed to different combinations of chemicals in our homes, 

at work, in food, and in the environment (air, water, and soil). These mixtures can be composed of tens, 

hundreds, or even thousands of different chemicals. To further complicate matters, each chemical could have 

several toxic effects, ranging from cancer to different kinds of noncancer or systemic effects. A question 

commonly asked is: How could the combined chemicals affect my health? Risk 

assessment has historically looked at effects from single chemicals then simply added the 

risks together, so in a way this is looking at the “tip of the iceberg” for chemical 

mixtures. To better evaluate the many cumulative exposures and effects that could be 

possible, approaches are being developed to assess risks more comprehensively. The 

U.S. Environmental Protection Agency (EPA) first published risk assessment guidelines 

for mixtures in 1986, with an update in 2000. The Agency formally defined a cumulative 

risk policy and offered planning and scoping guidance in 1997, and a cumulative risk 

assessment framework was published in 2003 following extensive input from many 

people. Three types of cumulative risk “icebergs” are worth noting: interactions, exposures, and effects. 

What Are Toxic Interactions? We might understand the toxic effects of certain chemicals in the mixtures to 

which we are exposed, but we simply don’t have this information for all possible mixtures. Where appropriate 

data exist for the mixture itself, such as diesel exhaust or cigarette smoke, it is evaluated as a whole mixture. 

Where such data are not available, information about the individual chemicals in the mixture is used, considering 

their proportions and relative toxicities. We usually assume each chemical would cause the same harm 

regardless of whether it was in a mixture or by itself. Thus, dose addition is our basic assumption, where the 

toxicity of a mixture is predicted by adding the toxicities caused by the doses of its individual chemicals, as 

adjusted to account for their relative toxicities. For example, if a mixture has two chemicals in equal amounts 

and the first is twice as toxic as the second, the toxicity of the mixture would be the same as adding three doses 

of the first. The special case where the chemicals are toxicologically independent (i.e., cause harm in different 

ways) is described by response addition. Here, the combined toxic response is the same as if the responses 

caused by the individual chemicals were added. For instance, tranquilizers and alcohol both depress the central 

nervous system but by different means. If a person is exposed to both, the effect on the central nervous system is 

the same as the sum of the effects caused by each separately. 

In fact, the toxicity of a mixture might be lower or higher than predicted from the known effects of each chemical 

acting alone. The influence one chemical has on the toxic effect of another is called a toxic interaction. The 

EPA mixtures guidance identifies toxic interaction as being 

something other than our default assumption of simple addition. 

Four types of interactions have been defined, relative to dose 

addition (see box). A synergistic interaction can be illustrated by 

alcohol and the solvent carbon tetrachloride, each of which harms 

the liver. Exposure to both damages the liver much more than 

predicted by simple dose addition. An antagonistic interaction can 

be illustrated by arsenic and lead. Exposure to both causes less harm 

to the kidney and blood than predicted by dose addition. 

Potentiation can be illustrated by carbon tetrachloride and 

isopropanol, which is rubbing alcohol (also found in perfume) and is 

not considered a liver toxin. Inside the body, isopropanol interacts 

with carbon tetrachloride and causes it to damage the liver more than 

it would have alone. Inhibition is the basis of some antidotes, where 

you take a dose of a chemical that does not harm you to reduce the 

harmful effect of another. When risks are estimated by assuming only dose or response addition without 

considering toxic interactions, perhaps only the tip of the iceberg is being addressed. 

Interaction Definition 

Synergism The combined effect of two or 

more chemicals is > predicted 

by dose addition 

Antagonism The combined effect of two or 

more chemicals is < predicted 

by dose addition 

Potentiation Exposure to one chemical that 

is not toxic itself increases the 

toxicity of second chemical 

when exposed to both 

Inhibition Exposure to one chemical at a 

nontoxic dose decreases the 

toxic effect of another 

Basic Concepts for Mixtures Risk Assessment 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

What Are Mixture Exposures? Some mixtures are manufactured before they are released to the environment 

(like a specific pesticide formulation), while others are formed during combustion (such as in our car engines) or 

by other chemical transformations. Chemical interactions in the environment can alter the composition and 

behavior of a mixture over time. We are exposed to mixtures everywhere. For example, when we breathe 

outside we can be taking in a mixture of all the air emissions from point sources (such as chimneys and 

smokestacks) and mobile sources (such as cars, buses, and planes) in the area. 

This has led many federal programs to look at cumulative exposures at the 

community scale, by studying releases alone or by combining modeling with 

emissions measurements. For example, a regional air impact modeling 

initiative (RAIMI) launched by EPA in 1999 considers risks at the 

neighborhood level from exposure to multiple air contaminants from multiple 

sources and pathways. An initial RAIMI study has been completed for an area 

in Jefferson County, Texas. This type of integrated assessment is being 

pursued under several national programs. These include the Cumulative Exposure Project (CEP), the NationalScale Air Toxics Assessment (NATA), the Residual Risk Report to Congress, and the Integrated Urban Air 

Toxics Strategy. These types of assessments look at combined doses from many sources rather than only 

considering the chemicals from a single location. A full look at this second exposure “iceberg” could include 

lifestyle-related exposures, such as to exposures to household pesticides, carpeting, or cigarettes. 

What Are Mixtures Effects? In addition to considering the combination of chemicals in environmental 

mixtures and the many ways we can be exposed to them, we also have to realize that each chemical in a mixture 

could cause multiple health effects. Under the classic risk assessment approach, the 

critical or primary effect is generally the most sensitive effect (that first seen as the 

dose is increased above the level where no adverse effects are observed), and it is 

often determined from animal studies; this generally serves as the basis for the 

toxicity value we use to estimate human health risks. However, virtually all 

chemicals can also have secondary effects, which do not occur until we are exposed 

to larger amounts. Information about primary effects and some secondary effects is 

contained in EPA’s Integrated Risk Information System (IRIS) and in the Agency 

for Toxic Substances and Disease Registry (ATSDR) toxicological profiles, both of 

which are available on-line. As science progresses, we will continue to improve our 

ability to evaluate the role of secondary effects. This is important because in theory, 

risks calculated based only on critical effects might not be considered a problem if 

the chemicals in a mixture each impact different organs (e.g., liver) or systems (e.g., 

nervous system) at tolerable levels. However, the evaluation of possible adverse health effects should extend 

beyond the chemicals’ primary targets to also consider their cumulative impact on other organs or systems that 

could be harmed as exposures increase. This requires an extensive toxicity evaluation. In a sense, assessing 

only the primary effects might miss important cumulative secondary effects – a third “iceberg” to keep in mind. 

Where Are We Heading? In the past, risk assessments have typically evaluated risks from distinct chemicals, 

operations, processes, waste streams, or contaminated media. As assessment methods keep pace with emerging 

scientific knowledge, EPA and others will continue to further evaluate contributions from (1) other pollutants 

and processes beyond those at the facility being assessed, (2) many sources of a single pollutant in a given 

community, and (3) the combination of sources, chemicals, and exposures that affect a given community. 

Interactions among chemicals and the cumulative result of secondary effects are being considered as part of 

these enhanced assessments. Additional research will better illuminate the often “hidden” components of 

cumulative health risk from environmental exposures. 

Where Can I Find More Information? Key sources include: the EPA mixtures guidance 

(http://www.epa.gov/ncea/raf/pdfs/chem_mix/chem_mix_08_2001.pdf), EPA cumulative risk 

framework (http://cfpub.epa.gov/ncea/raf/recordisplay.cfm?deid=54944.), EPA IRIS database 

(http://www.epa.gov/iris), and RAIMI (http://www.epa.gov/earth1r6/6pd/rcra_c/raimi/raimi.htm), 

and the ATSDR interaction profiles (http://www.atsdr.cdc.gov/iphome.html). 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

Mixtures of Arsenic, Cadmium, Chromium, and Lead 

(This fact sheet summarizes information for combinations of metals in this set, with a focus on ingestion 

of inorganic forms, to support analyses at contaminated sites. Companion fact sheets provide chemicalspecific information on common use, general environmental levels, and toxicity, also by other routes.) 

Do These Metals Naturally Coexist? Arsenic, cadmium, chromium, and lead all occur 

naturally in the environment and are found in all materials – soils, plants, animals, and 

humans – typically as salts. These metals cannot be destroyed, nor do they degrade; 

however they can be converted to organic forms by biological action both in the 

environment and within the body. (This fact sheet focuses on inorganic forms.)

What Common Uses Could Result in a Combined Presence? These metals share many industrial 

uses, including in metallurgy. Arsenic and 

cadmium are byproducts of lead 

production, and various combinations are 

used in materials and processes as 

highlighted at right. The Agency for Toxic 

Substances Disease Registry (ATSDR) has identified 

these four metals as soil contaminants at about 15% of 

the waste sites reviewed. Amounts and proportions of 

each vary in content and concentration across different sites. 

What Health Effects Are Indicated? 

Chronic oral exposures often represent a 

route of concern for contaminated sites as 

contaminants can migrate to groundwater 

over time. Thus, this fact sheet focuses on 

ingestion-related effects. Key organs or 

systems affected by intermediate 

(subchronic) to long-term (chronic) oral 

exposures to the individual metals are 

highlighted at right. Shown are the critical 

organs and systems (where the first adverse 

effect is observed, in bold) and those affected 

at higher levels that are common across two 

or more metals, or that are known to be 

affected by another metal in the mixture. 

(No critical effect has been defined for 

chromium by the oral route; the respiratory 

system is the primary target for inhalation.) 

What Are The Joint Toxicities? Joint 

toxicity refers to the outcome of two or more 

chemicals acting together; three categories of 

joint toxicity are: greater than additive 

(synergism and potentiation); additive (no 

interaction); and less than additive 

(antagonism and inhibition). Additivity is 

the default assumption for evaluating health 

effects of multiple chemicals. Toxicological 

Uses Combinations 

Metal products, alloys All four 

Dyes, pigments, paints Chromium, cadmium, lead 

Batteries Cadmium; lead and arsenic 

Anti-corrosion/rust coat Chromium and arsenic 

Solder Lead and arsenic 

Wood preservative Arsenic and chromium 
interactions can either increase or decrease the apparent toxicity of a mixture relative to that expected 

from simple addition. No specific health study has been conducted for the quaternary mixture of arsenic, 

cadmium, chromium, and lead. A few studies have investigated the effects of three metals in 

combination, but most have studied pairs within this set. The lead-cadmium pair has been studied most, 

including in human epidemiological studies and oral animal studies. Data from key studies and evidence 

of interactions as given in the draft ATSDR interaction profile are summarized in the table below. 

Joint Toxicity for Selected Organ/System Endpoints Following Ingestion 

Arsenic on Cadmium on Chromium on Lead on

Endpoint Cd Cr Pb As Cr Pb As Cd Pb As Cd Cr 

Nervous system ↑ ↑ ↑ ↔

Kidney ↔ ↓ ↓ ↔ ↔ ↓ ↓ ↔ ↓ ↔

Hematological system ↓ ↓ ↓ ↓ ↓ ↔

Reproductive: male ↓ na ↑ na na ↑

Skin na na na na na ↑ na na na 

Cancer na na na na na ↑ na na na na 

As = arsenic, Cd = cadmium, Cr = chromium, Pb = lead; ↑ = interactive effects are more than additive or 

one metal enhances an effect induced only by the other metal; ↓ = interactive effects are less than additive or 

one metal protects against an effect induced only by the other metal; ↔ = results are inconclusive or do not 

suggest that effects are more or less than additive; blank = relevant information is not available; na = not 

applicable because oral exposure to this metal does not cause the indicated health endpoint. Note that for the 

cardiovascular system, results were inconclusive or unavailable for all pairs. 

For neurological effects, the predicted direction of joint toxic action for the mixture is greater than 

additive for several pairs (arsenic-lead, cadmium-lead, lead-arsenic), which indicates the health hazard of 

those mixtures might be somewhat greater than that estimated by endpoint-specific hazard indices for this 

endpoint; the same higher-than-additive effect is indicated for the male reproductive system (testes) for 

cadmium and lead acting on each other. In contrast, for effects on the kidney and blood, the predicted 

direction of joint toxic action is less than additive for several metal pairs, indicating that the health hazard 

might be somewhat less than estimated by endpoint-specific hazard indices. 

What Is the Joint Risk? No specific data exist to quantify the joint risks of the mixture of arsenic, lead, 

cadmium and chromium. Endpoints of potential concern for this mixture include critical effects of the 

individual metals as well as the common targets of toxicity that might become significant due to 

additivity (considering secondary effects) or certain interactions, as indicated in the table above (noting 

again that additional interactions protect against other adverse effects). The ATSDR recommends using a 

hazard index method with the target-organ toxicity dose (TTD) modification and qualitative weight-ofevidence (WOE) method to assess the additive and interactive actions of the mixture components. These 

methods are suggested only when exposures are significant, i.e., only if the hazard quotients for two or 

more metals are 0.1 or greater. If only one or none of the metals have a hazard quotient at or above that 

level, then no further assessment of joint toxic action is needed because additivity and/or interactions are 

unlikely to result in a significant health hazard. 

Where Can I Find More Information? More information can be found in the primary information 

source for this overview: the draft interaction profile for arsenic, cadmium, chromium and 

lead prepared by ATSDR (http://www.atsdr.cdc.gov/iphome.html). Additional information 

can be found in the companion fact sheets for each metal and in the sources listed in those 

fact sheets. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

 

TABLE 1. Mortality and Morbidity Risk Coefficients for Selected Radionuclidesa 

Lifetime Cancer Risk 

Isotope Mortality Morbidity 

Inhalation Ingestion External Inhalation Ingestion External 

Americium-241 2.4 × 10-8 9.5 × 10-11 1.9 × 10-8 2.8 × 10-8 1.3 × 10-10 2.8 × 10-8

Americium-242m 1.3 × 10-8 6.8 × 10-11 - 1.6 × 10-8 9.0 × 10-11 - 

Americium-243 2.3 × 10-8 9.8 × 10-11 4.3 × 10-7 2.7 × 10-8 1.4 × 10-10 6.4 × 10-7

Berkelium-247 4.0 × 10-8 1.2 × 10-10 2.1 × 10-7 4.8 × 10-8 1.6 × 10-10 3.1 × 10-7

Cadmium-109 2.0 × 10-11 4.2 × 10-12 - 2.2 × 10-11 6.7 × 10-12 - 

Cadmium-113 8.1 × 10-11 2.0 × 10-11 - 1.1 × 10-10 2.9 × 10-11 - 

Cadmium-113m 9.3 × 10-11 2.5 × 10-11 - 1.3 × 10-10 3.6 × 10-11 - 

Californium-248 2.4 × 10-8 3.8 × 10-11 - 2.6 × 10-8 6.2 × 10-11 - 

Californium-249 4.0 × 10-8 1.2 × 10-10 9.3 × 10-7 4.8 × 10-8 1.6 × 10-10 1.4 × 10-6

Californium-250 3.5 × 10-8 8.0 × 10-11 - 3.7 × 10-8 1.1 × 10-10 - 

Californium-251 4.1 × 10-8 1.3 × 10-10 2.6 × 10-7 4.9 × 10-8 1.7 × 10-10 3.8 × 10-7

Californium-252b 2.1 × 10-8 4.1 × 10-11 - 2.2 × 10-8 5.7 × 10-11 - 

Carbon-14 6.5 × 10-12 1.4 × 10-12 - 7.1 × 10-12 2.0 × 10-12 - 

Cesium-134 1.1 × 10-11 3.5 × 10-11 4.8 × 10-6 1.6 × 10-11 5.1 × 10-11 7.1 × 10-6

Cesium-135 1.3 × 10-12 4.0 × 10-12 - 1.9 × 10-12 5.9 × 10-12 - 

Cesium-137 8.1 × 10-12 2.5 × 10-11 1.7 × 10-6 1.2 × 10-11 3.7 × 10-11 2.5 × 10-6

Chlorine-36 9.6 × 10-11 2.9 × 10-12 - 1.0 × 10-10 4.4 × 10-12 - 

Cobalt-57 1.8 × 10-12 9.0 × 10-13 2.4 × 10-7 2.1 × 10-12 1.5 × 10-12 3.6 × 10-7

Cobalt-60 3.0 × 10-11 1.4 × 10-11 8.5 × 10-6 3.6 × 10-11 2.2 × 10-11 1.2 × 10-5

Curium-242 1.4 × 10-8 3.2 × 10-11 - 1.5 × 10-8 5.5 × 10-11 - 

Curium-243 2.4 × 10-8 8.5 × 10-11 2.9 × 10-7 2.7 × 10-8 1.2 × 10-10 4.2 × 10-7

Curium-244 2.3 × 10-8 7.5 × 10-11 - 2.5 × 10-8 1.1 × 10-10 - 

Curium-245 2.4 × 10-8 9.5 × 10-11 1.6 × 10-7 2.8 × 10-8 1.3 × 10-10 2.4 × 10-7

Curium-246 2.4 × 10-8 9.3 × 10-11 - 2.8 × 10-8 1.3 × 10-10 - 

Curium-247 2.2 × 10-8 9.1 × 10-11 9.3 × 10-7 2.5 × 10-8 1.3 × 10-10 1.4 × 10-6

Curium-248c 8.8 × 10-8 3.4 × 10-10 - 1.0 × 10-7 4.8 × 10-10 - 

Curium-250c 5.0 × 10-7 2.0 × 10-9 9.7 × 10-7 5.8 × 10-7 2.8 × 10-9 1.4 × 10-6

Europium-150 2.1 × 10-10 3.6 × 10-12 4.4 × 10-6 2.6 × 10-10 6.1 × 10-12 6.5 × 10-6

Europium-152 1.5 × 10-10 5.0 × 10-12 3.6 × 10-6 1.9 × 10-10 8.7 × 10-12 5.3 × 10-6

Europium-154 1.7 × 10-10 8.5 × 10-12 4.0 × 10-6 2.1 × 10-10 1.5 × 10-11 5.8 × 10-6

Europium-155 1.7 × 10-11 1.6 × 10-12 8.4 × 10-8 1.9 × 10-11 2.8 × 10-12 1.2 × 10-7

Iodine-129 6.2 × 10-12 2.0 × 10-11 - 6.1 × 10-11 1.9 × 10-10 - 

Iodine-131 2.1 × 10-12 6.9 × 10-12 1.1 × 10-6 1.9 × 10-11 6.5 × 10-11 1.6 × 10-6

Iridium-192 2.1 × 10-11 6.0 × 10-12 2.3 × 10-6 2.4 × 10-11 1.1 × 10-11 3.4 × 10-6

Iridium-192m 1.7 × 10-11 8.7 × 10-13 3.7 × 10-7 2.0 × 10-11 1.3 × 10-12 5.4 × 10-7

Iridium-194m 4.0 × 10-11 7.3 × 10-12 6.9 × 10-6 4.6 × 10-11 1.3 × 10-11 1.0 × 10-5

Neptunium-235 1.0 × 10-12 2.8 × 10-13 - 1.2 × 10-12 5.1 × 10-13 - 

Neptunium-236 2.6 × 10-9 1.5 × 10-11 2.2 × 10-7 3.0 × 10-9 2.3 × 10-11 3.2 × 10-7

Neptunium-237 1.5 × 10-8 5.8 × 10-11 5.4 × 10-7 1.8 × 10-8 9.1 × 10-11 8.0 × 10-7

Nickel-59 3.6 × 10-13 2.3 × 10-13 - 4.7 × 10-13 3.9 × 10-13 - 

Nickel-63 1.4 × 10-12 5.7 × 10-13 - 1.6 × 10-12 9.5 × 10-13 - 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

 

Lifetime Cancer Risk 

Isotope Mortality Morbidity 

Inhalation Ingestion External Inhalation Ingestion External 

Plutonium-236 2.1 × 10-8 6.9 × 10-11 - 2.3 × 10-8 9.9 × 10-11 - 

Plutonium-238 3.0 × 10-8 1.3 × 10-10 - 3.4 × 10-8 1.7 × 10-10 - 

Plutonium-239 2.9 × 10-8 1.3 × 10-10 - 3.3 × 10-8 1.7 × 10-10 - 

Plutonium-240 2.9 × 10-8 1.3 × 10-10 - 3.3 × 10-8 1.7 × 10-10 - 

Plutonium-241 2.8 × 10-10 1.9 × 10-12 - 3.3 × 10-10 2.3 × 10-12 - 

Plutonium-242 2.8 × 10-8 1.3 × 10-10 - 3.1 × 10-8 1.7 × 10-10 - 

Plutonium-244d 2.7 × 10-8 1.3 × 10-10 1.0 × 10-6 3.1 × 10-8 1.6 × 10-10 1.5 × 10-6

Polonium-210 1.0 × 10-8 1.6 × 10-9 - 1.1 × 10-8 2.3 × 10-9 - 

Potassium-40 2.1 × 10-10 2.2 × 10-11 5.5 × 10-7 2.2 × 10-10 3.4 × 10-11 8.0 × 10-7

Protactinium-231 2.5 × 10-7 6.0 × 10-10 1.1 × 10-6 2.6 × 10-7 8.8 × 10-10 1.6 × 10-6

Radium-226 2.4 × 10-8 2.9 × 10-9 5.8 × 10-6 2.5 × 10-8 4.0 × 10-9 8.5 × 10-6

Radium-228 9.0 × 10-8 1.3 × 10-9 8.4 × 10-6 9.7 × 10-8 1.9 × 10-9 1.2 × 10-5

Samarium-146 1.2 × 10-8 4.0 × 10-11 - 1.4 × 10-8 5.3 × 10-11 - 

Samarium-151 8.6 × 10-12 4.6 × 10-13 - 9.2 × 10-12 8.1 × 10-13 - 

Selenium-79 2.3 × 10-12 6.7 × 10-12 - 3.3 × 10-12 9.7 × 10-12 - 

Strontium-90 1.0 × 10-10 7.5 × 10-11 - 1.1 × 10-10 9.5 × 10-11 - 

Technetium-97 7.6 × 10-13 2.3 × 10-13 - 8.5 × 10-13 3.9 × 10-13 - 

Technetium-98 2.6 × 10-11 6.0 × 10-12 4.4 × 10-6 3.0 × 10-11 1.0 × 10-11 6.5 × 10-6

Technetium-99 1.3 × 10-11 2.3 × 10-12 - 1.4 × 10-11 4.0 × 10-12 - 

Thorium-229 2.2 × 10-7 4.7 × 10-10 7.8 × 10-7 2.3 × 10-7 7.2 × 10-10 1.2 × 10-6

Thorium-230 2.7 × 10-8 8.0 × 10-11 - 2.9 × 10-8 1.2 × 10-10 - 

Thorium-232 4.1 × 10-8 9.1 × 10-11 - 4.3 × 10-8 1.3 × 10-10 - 

Tin-121m 4.1 × 10-11 2.9 × 10-12 - 4.4 × 10-11 5.1 × 10-12 - 

Tin-126 3.9 × 10-10 3.0 × 10-11 8.8 × 10-6 4.2 × 10-10 5.3 × 10-11 1.3 × 10-5

Tritium (H-3) 3.9 × 10-14 4.4 × 10-14 - 5.6 × 10-14 6.5 × 10-14 - 

Uranium-232 1.8 × 10-8 2.7 × 10-10 - 1.9 × 10-8 3.9 × 10-10 - 

Uranium-233 1.1 × 10-8 6.3 × 10-11 - 1.2 × 10-8 9.7 × 10-11 - 

Uranium-234 1.1 × 10-8 6.1 × 10-11 - 1.1 × 10-8 9.5 × 10-11 - 

Uranium-235 9.5 × 10-9 6.2 × 10-11 3.7 × 10-7 1.0 × 10-8 9.8 × 10-11 5.4 × 10-7

Uranium-236 9.9 × 10-9 5.8 × 10-11 - 1.0 × 10-8 9.0 × 10-11 - 

Uranium-238 8.8 × 10-9 7.5 × 10-11 - 9.3 × 10-9 1.2 × 10-10 - 

Zirconium-93 8.4 × 10-12 1.7 × 10-12 - 9.2 × 10-12 2.6 × 10-12 - 

a

 This table provides selected risk coefficients for inhalation and dietary ingestion of various radionuclides, and for 

external gamma irradiation where that entry is appropriate. (Source: Cancer Risk Coefficients for Environmental 

Exposure to Radionuclides, Federal Guidance Report 13, U.S. Environmental Protection Agency, EPA 402-R-99-

001, September 1999.) The mortality risk represents the lifetime risk of incurring a fatal cancer, and the 

morbidity risk represents the risk of incurring all cancers (fatal and non-fatal). Values are averaged over all ages 

and both genders. (For context, 10-9 is a billionth, 10-12 is a trillionth, and a pCi is a picocurie, or a trillionth of a 

curie.) To convert to standard international units, multiply by 27 pCi per becquerel (Bq). Values shown here 

include the contributions from short-lived decay products, as indicated in the radionuclide-specific fact sheets. 

(For example, strontium-90 includes the contribution from yttrium-90, and uranium-238 includes the contribution 

from thorium-234). 

 For ingestion and inhalation, units are risk per pCi. For inhalation, the values corresponding to the recommended 

default absorption type for particulates are shown; the maximum value is given if no absorption type was 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

 

recommended. For ingestion, the dietary values shown are the highest for ingestion exposures; values for tap 

water ingestion are typically 70 to 80% of those for diet. The values for tritium are for tritiated water. 

 

 For external exposure, risk coefficients are given for those radionuclides having gamma-ray energies in excess of 

0.03 MeV per decay, accounting for the fraction of time that the radioactive decay results in the emission of 

gamma rays. A dash indicates the radionuclide or its decay products does not emit significant gamma radiation

(see the companion fact sheet on Radioactive Properties, Internal Distribution, and Risk Coefficients). Units for 

external gamma risk coefficients shown in the table are risk per pCi/g soil for one year of exposure. Although no 

inhalation or ingestion coefficients are available for krypton isotopes, coefficients do exist for external gamma 

exposures. Submersion in a cloud of krypton gas poses the highest risk, and the following values are in units of 

risk per pCi/cm3 air for one year of exposure. For krypton-81, mortality and morbidity risk coefficients are 

1.5 × 10-5 and 2.3 × 10-5, respectively. For krypton-85, mortality and morbidity risk coefficients are 8.5 × 10-6 and 

1.2 × 10-5, respectively. 

b

 Standard risk coefficients are not available for californium-252. To help address this gap, values shown here have 

been derived by multiplying the standard risk coefficients for californium-250 by the ratios of the dose conversion 

factors for californium-252 to californium-250, for the given exposure pathways and endpoints. 

c

 Standard risk coefficients are not available for curium-248 or curium-250. To help address this gap, values shown 

here have been derived. For curium-248, standard risk coefficients for curium-246 were multiplied by the ratios 

of the standard dose conversion factors for curium-248 to curium-246, for the given exposure pathways and 

endpoints. For curium-250, standard risk coefficients for curium-246 were multiplied by the ratios of the standard 

dose conversion factors for curium-250 to curium-246, for the given exposure pathways and endpoints. The risk 

coefficient for external exposure for curium-250 is attributable to its short-lived radioactive decay products. 

d

 Standard risk coefficients are not available for plutonium-244. To help address this gap, values shown here have 

been derived by multiplying the standard risk coefficients for plutonium-242 by the ratios of the dose conversion 

factors for plutonium-244 to plutonium-242, for the given exposure pathways and endpoints. The risk coefficient 

for external exposure is attributable to its short-lived radioactive decay products, principally neptunium-240m. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

 

TABLE 2. Summary Radioactive Properties for Selected Radionuclidesa 

Radiation Energy (MeV)

Isotope Half-Life 

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ)

Americium-241 430 yr 3.5 α 5.5 0.052 0.033 

Americium-242m 150 yr 9.8 IT 0.025 0.044 0.0051 

Americium-242 16 hr 820,000 β, EC - 0.18 0.018 

Americium-243 7,400 yr 0.20 α 5.3 0.022 0.056 

Neptunium-239 2.4 days 230,000 β - 0.26 0.17 

Berkelium-247 1,400 yr 1.1 α 5.6 0.061 0.11 

Cadmium-109 1.3 yr 2,600 EC - 0.083 0.026 

Cadmium-113e 9.3 × 1015 yr 3.4 × 10-13 β - 0.093 - 

Cadmium-113m 14 yr 240 β - 0.19 - 

Californium-248 330 days 1,600 α 6.3 0.0060 0.0013 

Californium-249 350 yr 4.1 α 5.8 0.044 0.33 

Californium-250 13 yr 110 α 6.0 0.0057 0.0012 

Californium-251 900 yr 1.6 α 5.8 0.20 0.13 

Californium-252 2.6 yr 540 α 5.9 0.0056 0.0012 

Carbon-14e 5.700 yr 4.5 β - 0.049 - 

Cesium-134 2.1 yr 1,300 β - 0.16 1.6 

Cesium-135 2.3 million yr 0.0012 β - 0.067 - 

Cesium-137 30 yr 88 β - 0.19 - 

Barium-137m (95%) 2.6 min 540 million IT - 0.065 0.60 

Chlorine-36 300,000 yr 0.033 β, EC - 0.027 <0.001 

Cobalt-57 270 days 8,600 EC - 0.019 0.13 

Cobalt-60 5.3 yr 1,100 β - 0.097 2.5 

Curium-242 160 days 3,400 α 6.1 0.010 0.0018 

Curium-243 29 yr 52 α 5.8 0.14 0.13 

Curium-244 18 yr 82 α 5.8 0.086 00017 

Curium-245 8,500 yr 0.17 α 5.4 0.065 0.096 

Curium-246 4,700 yr 0.31 α 5.4 0.0080 0.0015 

Curium-247 16 million yr 0.000094 α 4.9 0.021 0.32 

Plutonium-243 5.0 hr 2.6 million β - 0.17 0.026 

Curium-248 340,000 yr 0.0043 α 4.7 0.0060 0.0012 

Curium-250 6,900 yr 0.21 α, β 1.3 0.0016 - 

Plutonium-246 (25%) 11 days 49,000 β - 0.13 0.14 

Berkelium-250 (14%) 3.2 hr 3.9 million β - 0.29 0.89 

Americium-246 (25%) 39 min 20 million β - 0.66 0.70 

Europium-150 34 yr 70 EC - 0.044 1.5 

Europium-152 13 yr 180 β, EC - 0.14 1.2 

Europium-154 8.8 yr 270 β - 0.29 1.2 

Europium-155 5.0 yr 470 β - 0.063 0.061 

Iodine-129 16 million yr 0.00018 β - 0.064 0.025 

Iodine-131 8.0 days 130,000 β - 0.19 0.38 

Iridium-192 74 days 9,200 β, EC - 0.22 0.82 

Iridium-192m 240 yr 7.8 IT - - 0.16 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

 

Radiation Energy (MeV)

Isotope Half-Life 

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ)

Iridium-194m 170 days 4,000 β - 0.16 2.3 

Krypton-81 210,000 yr 0.021 EC - 0.0051 0.012 

Krypton-85 11 yr 400 β - 0.25 0.0022 

Neptunium-235 1.1 yr 1,400 EC <0.001 0.010 0.0071 

Neptunium-236 120,000 yr 0.013 β, EC - 0.21 0.14 

Plutonium-236 (9%) 2.9 yr 540 α 5.8 0.013 0.0021 

Neptunium-237 2.1 million yr 0.00071 α 4.8 0.070 0.035 

Protactinium-233 27 days 21,000 β - 0.20 0.20 

Nickel-59 75,000 yr 0.082 EC - 0.0046 0.0024 

Nickel-63 96 yr 60 β - 0.17 - 

Plutonium-236 2.9 yr 540 α 5.8 0.013 0.0021 

Plutonium-238 88 yr 17 α 5.5 0.011 0.0018 

Plutonium-239 24,000 yr 0.063 α 5.1 0.0067 <0.001 

Plutonium-240 6,500 yr 0.23 α 5.2 0.011 0.0017 

Plutonium-241 14 yr 100 β <0.001 0.0052 <0.001 

Plutonium-242 380,000 yr 0.0040 α 4.9 0.0087 0.0014 

Plutonium-244 83 million yr 0.000018 α 4.6 0.0071 0.0012 

Uranium-240 14 hr 940,000 β - 0.14 0.0076 

Neptunium-240m 7.4 min 110 million β - 0.68 0.34 

Polonium-208 2.9 yr 590 α 5.1 <0.001 <0.001 

Polonium-208 100 yr 17 α 4.9 <0.001 <0.001 

Potassium-40e 1.3 billion yr 0.0000071 β, EC - 0.52 0.16 

Protactinium-231e 33,000 yr 0.048 α 5.0 0.065 0.048 

Actinium-227e 22 yr 73 α, β 0.068 0.016 <0.001 

Thorium-227e (99%) 19 days 31,000 α 5.9 0.053 0.11 

Francium-223e (1%) 22 min 39 million β - 0.40 0.059 

Radium-223e 11 days 52,000 α 5.7 0.076 0.13 

Radon-219e 4.0 sec 13 billion α 6.8 0.0063 0.056 

Polonium-215e 0.0018 sec 30 trillion α 7.4 <0.001 <0.001 

Lead-211e 36 min 25 million β - 0.46 0.051 

Bismuth-211e 2.1 min 420 million α 6.6 0.010 0.047 

Thallium-207e 4.8 min 190 million β - 0.49 0.0022 

Radium-226e 1600 yr 1.0 α 4.8 0.0036 0.0067 

Radon-222e 3.8 days 160,000 α 5.5 <0.001 <0.001 

Polonium-218e 3.1 min 290 million α 6.0 <0.001 <0.001 

Lead-214e 27 min 33 million β - 0.29 0.25 

Bismuth-214e 20 min 45 million β - 0.66 1.5 

Polonium-214e 0.00016 sec 330 trillion α 7.7 <0.001 <0.001 

Lead-210e 22 yr 77 β - 0.038 0.0048 

Bismuth-210e 5.0 days 130,000 β - 0.39 - 

Polonium-210e 140 days 4,500 α 5.3 <0.001 <0.001 

Radium-228e 5.8 yr 280 β - 0.017 <0.001 

Actinium-228e 6.1 hr 2.3 million β - 0.48 0.97 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

 

Radiation Energy (MeV)

Isotope Half-Life 

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ)

Thorium-228e 1.9 yr 830 α 5.4 0.021 0.0033 

Radium-224e 3.7 days 160,000 α 5.7 0.0022 0.010 

Radon-220e 56 sec 930 million α 6.3 <0.001 <0.001 

Polonium-216e 0.15 sec 350 billion α 6.8 <0.001 <0.001 

Lead-212e 11 hr 1.4 million β - 0.18 0.15 

Bismuth-212e 61 min 15 million α, β 2.2 0.47 0.19 

Polonium-212e (64%) 0.00000031 sec 180,000 trillion α 8.8 - - 

Thallium-208e (36%) 3.1 min 300 million β - 0.60 3.4 

Samarium-146 100,000,000 yr 0.000024 α 2.5 - - 

Samarium-151 90 yr 27 β - 0.020 <0.001 

Selenium-79 650,000 yr 0.070 β - 0.056 - 

Strontium-90 29 yr 140 β - 0.20 - 

Yttrium-90 64 hr 550,000 β - 0.94 <0.001 

Technetium-97 2.6 million 0.0014 EC - 0.0056 0.011 

Technetium-98 4.2 million 0.00088 β - 0.16 1.4 

Technetium-99 210,000 0.017 β - 0.10 - 

Thorium-229 7,300 yr 0.22 α 4.9 0.12 0.096 

Radium-225 15 days 40,000 β - 0.11 0.014 

Actinium-225 10 days 59,000 α 5.8 0.022 0.018 

Francium-221 4.8 min 180 million α 6.3 0.010 0.031 

Astatine-217 0.032 sec 1.6 trillion α 7.1 <0.001 <0.001 

Bismuth-213 46 min 20 million α, β 0.13 0.44 0.13 

Polonium-213 (98%) 0.0000042 sec 13,000 trillion α 8.4 - - 

Thallium-209 (2%) 2.2 min 410 million β - 0.69 2.0 

Lead-209 3.3 hr 4.7 million β - 0.20 - 

Thorium-230e 77,000 yr 0.020 α 4.7 0.015 0.0016 

Thorium-232e 14 billion yr 0.00000011 α 4.0 0.012 0.0013 

Tin-121m 55 yr 54 β, IT - 0.035 0.0049 

Tin-121 (78%) 27 hr 970,000 β - 0.11 - 

Tin-126 250,000 yr 0.029 β - 0.17 0.057 

Antimony-126 12 days 85,000 β - 0.28 2.8 

Tritium (H-3)e 12 yr 9,800 β - 0.0057 - 

Uranium-232 72 hr 22 α 5.3 0.017 0.0022 

Uranium-233 160,000 yr 0.0098 α 4.8 0.0061 0.0013 

Uranium-234e 240,000 yr 0.0063 α 4.8 0.013 0.0017 

Uranium-235e 700 million yr 0.0000022 α 4.4 0.049 0.16 

Thorium-231e 26 hr 540,000 β - 0.17 0.026 

Uranium-236 23 million yr 0.000065 α 4.5 0.011 0.0016 

Uranium-238n 4.5 billion yr 0.00000034 α 4.2 0.010 0.0014 

Thorium-234e 24 days 23,000 β - 0.060 0.0093 

Protactinium-234me 1.2 min 690 million β - 0.82 0.012 

Zirconium-93 1.5 million yr 0.0025 β - 0.020 - 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

 

Radiation Energy (MeV)

Isotope Half-Life 

Specific 

Activity 

(Ci/g)

Decay 

Mode Alpha 

(α)

Beta 

(β)

Gamma 

(γ)

Niobium-93m 14 yr 290 IT - 0.028 0.0019 

a This table summarizes key radioactive properties of selected radionuclides and their associated decay products, 

which are indicated in italics. An “e” indicates the isotope exists naturally in the environment. A dash means the 

entry is not applicable, EC = electron capture, IT = isomeric transition, Ci = curie, g = gram, and MeV = million 

electron volts. Values are given to two significant figures. See the radionuclide-specific fact sheets for further 

information, and the companion fact sheet on Radioactive Properties, Internal Distribution, and Risk Coefficients

for an explanation of terms and interpretation of radiation energies. 
Argonne National Laboratory, EVS Human Health Fact Sheet, August 2005 

TABLE 3. Key References for the Radiological Fact Sheetsa 

Isotope References 

Multiple Argonne National Laboratory, 1993, Manual for Implementing Residual Radioactive Material Guidelines 

Using RESRAD, Version 5.0, ANL/EAD/LD-2, working draft for comment, September

 Argonne National Laboratory, 2001, User’s Manual for RESRAD Version 6, ANL/EAD/4, July 

Eisenbud, Merril, 1987, Environmental Radioactivity from Natural, Industrial, and Military Sources, Third 

Edition, Academic Press, Inc., Orlando, Florida

Eisenbud, Merril and Thomas Gesell, 1997, Environmental Radioactivity from Natural, Industrial, and 

Military Sources, Fourth Edition, Academic Press, San Diego, California 

International Commission on Radiological Protection, 1979, Limits for Intakes of Radionuclides by Workers, 

ICRP Publication 30, Part 1, Annals of the ICRP, 2 (3/4)

International Commission on Radiological Protection, 1980, Limits for Intakes of Radionuclides by Workers, 

ICRP Publication 30, Part 2, Annals of the ICRP, 4 (3/4) 

International Commission on Radiological Protection, 1981, Limits for Intakes of Radionuclides by Workers, 

ICRP Publication 30, Part 3, Annals of the ICRP, 6 (2/3)

International Commission on Radiological Protection, 1983, Radionuclide Transformations, Energy and 

Intensity of Emissions, ICRP Publication 38, Annals of the ICRP, 11-13

International Commission on Radiological Protection, 1989, Age-Dependent Doses to Members of the Public 

from Intake of Radionuclides, Part 1, ICRP Publication 56, Annals of the ICRP, 20 (2)

International Commission on Radiological Protection, 1989, Age-Dependent Doses to Members of the Public 

from Intake of Radionuclides, Part 2, ICRP Publication 67, Annals of the ICRP 23 (3/4)

International Commission on Radiological Protection, 1994, Dose Coefficients for Intakes of Radionuclides by 

Workers, ICRP Publication 68, Annals of the ICRP, 24 (4)

National Council on Radiation Protection and Measurements, 1987, Exposure of the Population in the United 

States and Canada from Natural Background Radiation, NCRP Report 94, Washington, D.C., December 30 

Shleien, Bernard (editor), 1992, The Health Physics and Radiological Health Handbook, Revised Edition, 

Scinta, Inc., Silver Spring, Maryland 

U.S. Environmental Protection Agency, 1988, Limiting Values of Radionuclide Intake And Air Concentration 

and Dose Conversion Factors For Inhalation, Submersion, And Ingestion, Federal Guidance Report No. 11, 

Office of Radiation Programs, EPA-520/1-88-020, September 

U.S. Environmental Protection Agency, 1993, External Exposure to Radionuclides in Air, Water, and Soil, 

Federal Guidance Report No. 12, Office of Radiation and Indoor Air, EPA 402-R-93-081, September 

U.S. Environmental Protection Agency, 1999, Cancer Risk Coefficients for Environmental Exposure to 

Radionuclides, Federal Guidance Report No. 13, Office of Radiation and Indoor Air, EPA 402-R-99-001, 

September 

U.S. Nuclear Regulatory Commission, 1979, Final Generic Environmental Impact Statement on Handling 

and Storage of Spent Light Water Reactor Fuel, Office of Nuclear Material Safety and Safeguards, Project 

No. M-4, NUREG-0575, August 

Information was also obtained from the following Internet addresses: 

http://atom.kaeri.re.kr/ 

http://cdfc.rug.ac.be/HealthRisk/default.htm 

http://environmentalchemistry.com/yogi/periodic/index.html 

http://ie.lbl.gov/fission/235ut.txt 

http://micronmetals.com/toc.htm 

http://nautilus.fis.uc.pt/st2.5/index-en.html 

http://nobel.scas.bcit.ca/resource/navptable.htm 

http://pearl1.lanl.gov/periodic/default.htm 

http://www.bayerus.com/msms/fun/pages/periodic/i_table.html 
Isotope References 

Multiple 

(cont’d)

Information was also obtained from the following Internet addresses (cont’d): 

http://www.cancer.org/docroot/STT/stt_0.asp, (Cancer Facts & Figures 2004), and 

http://www.cancer.org/downloads/STT/CAFF_finalPWSecured.pdf 

http://www.census.govhttp://www.principalmetals.com/utilities/periodic.htm 

http://www.chemicalelements.com/ 

http://www.chemsoc.org/viselements/pages/pertable_fla.htm (requires Shockwave plugin) 

http://www.resource-world.net/PerTable.htm 

http://www.speclab.com/elements/ 

http://www.ucc.ie/ucc/depts/chem/dolchem/html/elem/elem000.html 

http://www.umich.edu/~radinfo/

http://www.vcs.ethz.ch/chemglobe/ptoe/index.html 

http://www.webelements.com/ 

Americium International Commission on Radiological Protection, 1986, The Metabolism of Plutonium and Related 

Elements, ICRP Publication 48, Annals of the ICRP, 16(2/3) 

 Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.encarta.msn.com/find/Concise.asp?ti=017E7000 

http://www.world-nuclear.org/info/inf57.htm 

http://www.encyclopedia.com/printable/00435.html 

http://www.britannica.com/seo/a/americium/ 

http://www.uic.com.au/nip35.htm 

Cadmium Radionuclide-specific information was obtained from the following Internet addresses: 

http://minerals.usgs.gov/minerals/pubs/commodity/cadmium 

http://atom.kaeri.re.kr 

http://ntp-server.niehs.nih.gov/htdocs/8_RoC/RAC/Cadmium&cmpds.html 

http://www.ijc.org/boards/iaqab/meyer/cadmium.htm 

http://webmineral.com/data/Cadmium.html 

http://www.speclab.com/elements/cadmium.htm 

Californium International Commission on Radiological Protection, 1986, The Metabolism of Plutonium and Related 

Elements, ICRP Publication 48, Annals of the ICRP, 16(2/3) 

 Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.britannica.com/nobel/micro/98_61.html 

http://education.jlab.org/itselemental/ele098.html 

http://www.chemicalelements.com/elements/cf.html 

http://www.lenntch.com/Periodic-chart-elements/Cf-en.htm 

http://www.vcs.ethz.ch/chemglobe/ptoe/_/98.html 

Carbon-14 Cember, Herman, 1983, Introduction to Health Physics, Pergamon Press, Inc., Elmsford, New York 

 National Council on Radiation Protection and Measurements, 1985, Carbon-14 in the Environment, NCRP 

Report 81, Washington, D.C., May 15 

 Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.phschool.com/atschool/chemistry/AW/Student_Area/AWCHEM_SC5_ACT.html 

http://www.britannica.com/seo/n/nuclear-medicine/ 

http://crystal.biol.csufresno.edu:8080/projectsF98/535.html 

http://www.users.globalnet.co.uk/~freya01/carbon14.htm 

http://seattletimes.nwsource.com/news/health-science/html98/carb_082598.html 

Cesium National Council on Radiation Protection and Measurements, 1977, Cesium-137 from the Environment to 

Man: Metabolism and Dose, NCRP Report 52, Washington, D.C., January 15 

 Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.encyclopedia.com/articles/02502.html 

http://encarta.msn.com/find/Concise.asp?ti=04B4F000 
Isotope References 

Chlorine Radionuclide-specific information was obtained from the following Internet addresses: 

 http://www.encyclopedia.com/articles/02695.html 

 http://www.c3.org/ 

 http://www.cl2.com/benefits/index.html 

 http://www.history.rochester.edu/class/hanford/cdcPaper/onethree.html 

 http://wwwrcamnl.wr.usgs.gov/isoig/period/cl_iig.html 

Cobalt Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.infoplease.com/ce6/sci/A0812692.html 

http://www.eco-usa.net/toxics/cobalt.html 

http://www.atsdr.cdc.gov/tfacts33.html 

Curium International Commission on Radiological Protection, 1986, The Metabolism of Plutonium and Related 

Elements, ICRP Publication 48, Annals of the ICRP, 16(2/3) 

 Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.encyclopedia.com/articles/03336.html 

http://www.britannica.com/seo/c/curium/ 

http://encarta.msn.com/find/Concise.asp?ti=0457F000 

Europium Radionuclide-specific information was obtained from the following Internet addresses: 

 http://www.doegjpo.com/programs/hanf/AXReport/Report/sect03.htm 

http://minerals.usgs.gov/minerals/pubs/commodity/rare_earths/740301.pdf 

Iodine National Council on Radiation Protection and Measurements, 1977, Protection of the Thyroid Gland in the 

Event of Releases of Radioiodine, NCRP Report 55, Washington, D.C., August 1 

 National Council on Radiation Protection and Measurements, 1983, Iodine-129: Evaluation of Releases from 

Nuclear Power Generation, NCRP Report 75, Bethseda, Maryland, December 1 

 Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.doh.wa.gov/hanford/publications/health/mon10.htm 

http://cpmcnet.columbia.edu/dept/thyroid/RAI.html 

Ionizing 

Radiation 

Agency for Toxic Substances and Disease Registry, 1997 Toxicological Profile for Ionizing Radiation, draft 

for public comment, prepared by Research Triangle Institute, September 

 National Research Council, 1990, Health Effects of Exposure to Low Levels of Ionizing Radiation, BEIR V 

Report, report of the Committee on the Biological Effects of Ionizing Radiations, National Academy Press, 

Washington, D.C. 

 U. S. Environmental Protection Agency, Risk Assessment Guidance for Superfund Volume I Human Health 

Evaluation Manual (Part A), Interim Final, EPA/540/1-89/002, Office of Emergency and Remedial Response, 

Washington, D.C., December 

Iridium Radionuclide-specific information was obtained from the following Internet addresses: 

http://education.jlab.org/itselemental/ele077.html 

http://www.ualberta.ca/~slowpoke/1996/r19.html 

http://www.scn.org/~bh162/iridium.html 

http://www.nrc.gov/reading-rm/doc-collections/cfr/paro020/appb/ 

http://www.goodfellow.com/csp/active/static/A/IR00.html 

http://www.uic.com.au/nip26.htm 

http://periodic.lanl.gov/elements/77.html 

http://en.wikipedia.org/wiki/Iridium 

Krypton National Council on Radiation Protection and Measurements, 1975, Krypton-85 in the Atmosphere – 

Accumulation, Biological Significance, and Control Technology, NCRP Report No. 44, Washington, D.C., 

July 1 

 Radionuclide-specific information was obtained from the following Internet addresses: 

 http://encyclopedia.com/articles/07092.html 

 http://encarta.msn.com/index/conciseindex/34/0344A000.htm?z=1&pg=2&br=1 
Isotope References 

Krypton 

(cont’d)

Radionuclide-specific information was obtained from the following Internet addresses 

 http://www.c-f-c.com/specgas_products/krypton.htm 

 http://www.jamals.com/ehsan/electrical.htm 

Neptunium International Commission on Radiological Protection, 1986, The Metabolism of Plutonium and Related 

Elements, ICRP Publication 48, Annals of the ICRP, 16(2/3) 

 Radionuclide-specific information was obtained from the following Internet addresses: 

 http://nobel.scas.bcit.ca/resource/ptable/np.htm 

 http://encarta.msn.com/find/Concise.asp?z=1&pg=2&ti=761564503 

 http://search.eb.com/bol/topic?eu=56688&sctn=1 

 http://www.isis-online.org/publications/fmct/primer/Section_I.html 

Nickel Radionuclide-specific information was obtained from the following Internet addresses: 

http://encarta.msn.com/find/Concise.asp?ti=066E6000 

http://www.nipera.org/pro&use.htm 

Plutonium International Commission on Radiological Protection, 1972, The Metabolism of Compounds of Plutonium and 

other Actinides, ICRP Publication 19, May 

 International Commission on Radiological Protection, 1986, The Metabolism of Plutonium and Related 

Elements, ICRP Publication 48, Annals of the ICRP, 16(2/3) 

 Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.llnl.gov/csts/publications/sutcliffe/ 

http://www.pu.org/ 

http://plutonium-erl.actx.edu/ 

Polonium National Research Council, 1961, The Radiochemistry of Polonium, P.E. Figgens, Mound Laboratory, 

Miamisburg, Ohio, National Academy of Sciences, NAS-NS 3037, Washington, D.C. 

 National Research Council, 1988, Health Effects of Radon and Other Internally Deposited Alpha Emitters, 

BEIR IV Report, report of the Committee on the Biological Effects of Ionizing Radiations, National Academy 

Press, Washington, D.C. 

 Radionuclide-specific Information was obtained from the following Internet addresses: 

http://www.webelements.com/ 

http://www.wsu.edu/~wsurso/Isotopes/Pol210.html 

http://education.jlab.org/itselemental/ele084.html 

http://www.speclab.com/elements/ 

http://pearl1.lanl.gov/periodic/default.htm 

Potassium Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.encyclopedia.com/articles/10439.html 

http://encarta.msn.com/index/conciseindex/1C/01CB4000.htm?z=1&pg=2&br=1 

Protactinium Radionuclide-specific information was obtained from the following Internet addresses: 

 http://education.jlab.org/itselemental/ele091.html 

 http://www.webspinners.futura.net/apurdy/element_91_100.htm 

 http://www.encyclopedia.com/articlesnew/10559.html 

 http://encarta.msn.com/index/conciseindex/5C/05C22000.htm?z=1&pg=2&br=1 

 http://140.198.18.108/periodic/Pa.html 

Radium National Research Council, 1988, Health Risks of Radon and Other Internally Deposited Alpha-Emitters, 

BEIR IV Report, report of the Committee on the Biological Effects of Ionizing Radiations, National Academy 

Press, Washington, D.C. 

 Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.infoplease.com/ce6/sci/A0840951.html 

http://www.atsdr.cdc.gov/ToxProfiles/phs9022.html 

http://www.epa.gov/ttn/uatw/hlthef/radionuc.html 
Isotope References 

Samarium Radionuclide-specific information was obtained from the following Internet addresses: 

http://minerals.usgs.gov/minerals/pubs/commodity/rare_earths/740301.pdf 

http://encarta.msn.com/index/conciseindex/0D/00D52000.htm?z=1&pg=2&br=1 

http://www.ornl.gov/isotopes/s_sm.html 

http://www.encyclopedia.com/articles/11397.html 

http://www.geology.wisc.edu/~jill/samar.html 

Selenium Radionuclide-specific information was obtained from the following Internet addresses: 

 http://www.lbl.gov/MicroWorlds/Wetlands/WetlandsClue1.html 

 http://www.state.sd.us/doa/das/selenium.htm 

 http://www.encyclopedia.com/articles/11665.html 

 http://encarta.msn.com/index/conciseindex/22/0223E000.htm?z=1&pg=2&br=1=1 

 http://sutekh.nd.rl.ac.uk/cgi-bin/CoNquery?nuc=Se79 

Strontium National Council on Radiation Protection and Measurements, 1991, Some Aspects of Strontium Radiobiology, 

NCRP Report No. 110, Bethesda, Maryland., August 31 

 Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.britannica.com/seo/s/strontium/ 

http://www.epa.gov/iris/subst/0550.htm 

Technetium Radionuclide-specific information was obtained from the following Internet addresses: 

http://encarta.msn.com/find/Concise.asp?ti=03A19000 

http://www.encyclopedia.com/articles/12685.html 

Thorium Archer, V.E., J.K. Wagoner, and F.E. Lundin, 1973, Cancer Mortality among Uranium Mill Workers, Journal 

of Occupational Medicine, 15:11-14 

 Burkart, W., 1991, Uranium, Thorium, and Decay Products, in Metals and Their Components in the 

Environment: Occurrence, Analysis, and Biological Relevance, E. Merian (editor), VCH Publications, NY 

National Research Council, 1988, Health Risks of Radon and Other Internally Deposited Alpha-Emitters, 

BEIR IV Report, report of the Committee on the Biological Effects of Ionizing Radiations, National Academy 

Press, Washington, D.C. 

 Polednak, A.P., A.F. Stehney, and H.F. Lucas, 1983, Mortality among Male Workers at a ThoriumProcessing Plant, Health Physics, 44(Supplement 1):239-251 

 Stehney, A.F., et al., 1980, Health Status and Body Radioactivity of Former Thorium Workers, NUREG/CR1420 (ANL-80-37), prepared by Argonne National Laboratory, Argonne, Ill., for U.S. Department of Energy 

and U.S. Nuclear Regulatory Commission, January 

 Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.atsdr.cdc.gov/ToxProfiles/phs9025.html 

http://tonto.eia.doe.gov/FTPROOT/nuclear/047896.pdf 

http://www.thorium-waste.com/Appl.html 

http://www.thorium-waste.com/Radiation.html 

http://www.thorium-waste.com/Gensci.html 

http://encarta.msn.com/find/Concise.asp?ti=01C11000 

http://www.infoplease.com/ce6/sci/A0848573.html 

http://www.britannica.com/seo/t/thorium/ 

http://www.encyclopedia.com/articles/08646.html 

Tin Zhang, S., et al., 1996, J. Radioanal.Nucl Chem., Letters, 212(2):93-9 

 Radionuclide-specific information was obtained from the following Internet addresses: 

 http://encarta.msn.com/find/Concise.asp?ti=02ECC000 

 http://www.itri.co.uk/tinuses.htm 

 http://www.environmentalchemistry.com/yogi/periodic/Sn.html 
Tritium Argonne National Laboratory, 1999, Risk/Impact Technical Report for the Hanford Groundwater/Vadose 

Zone Integration Project, prepared for the U.S. Department of Energy, December 

 Makhijani, A., 1999, Statement on Tritium before the House Committee on Intergovernmental Coordination, 

State of Georgia: Institute for Energy and Environmental Research, October 19, accessed at 

http://www.ieer.org/comments/tritstmt.html 

 National Council on Radiation Protection and Measurements, 1979, Tritium in the Environment, NCRP 

Report No. 62, Washington, D.C., March 9 

 U.S. Department of Energy, 1999, Final Environmental Impact Statement for the Production of Tritium in a 

Commercial Light Water Reactor, DOE/EIS-0288, Appendix C, March U.S. Department of Energy, 1999, 

Final Environmental Impact 

 Zerriffi, Hisham, 1996, Tritium: The environmental, health, budgetary, and strategic effects of the 

Department of Energy’s decision to produce tritium, Institute for Energy and Environmental Research, March 

20, accessed at http://www.ieer.org/reports/tritium.html 

Uranium Agency for Toxic Substances and Disease Registry, 1990, Toxicological Profile for Uranium, TP-90/29, 

prepared by Syracuse Research Corporation, December 

 Archer, V.E., J.K. Wagoner, and F.E. Lundin, 1973, Cancer Mortality among Uranium Mill Workers, Journal 

of Occupational Medicine, 15:11-14 

 Argonne National Laboratory, 1999, Risk/Impact Technical Report for the Hanford Groundwater/Vadose 

Zone Integration Project, prepared for the U.S. Department of Energy, December 

 Berlin, M., and B. Rudell, 1986, Uranium, in Handbook on the Toxicology of Metals, 2nd ed., L. Friberg, 

G.F. Nordberg, and V. Voik (editors), Elsevier Science Publishers, New York 

 Burkart, W., 1991, Uranium, Thorium, and Decay Products, in Metals and Their Components in the 

Environment: Occurrence, Analysis, and Biological Relevance, E. Merian (editor), VCH Publications, New 

York, pp. 1275-1287 

 Dreesen, D.R., et al., 1982, Mobility and Bioavailability of Uranium Mill Tailings Contaminants, 

Environmental Science and Technology, 16(10):702-706 

 Hodge, H.C., J.N. Stannard, and J.B. Hursh (editors), 1973, Uranium-Plutonium Transplutonic Elements, 

Springer-Verlag, New York 

 Maynard, E.A., and H.C. Hodge, 1949, Studies of the Toxicity of Various Uranium Compounds When Fed to 

Experimental Animals, in The Pharmacology and Toxicology of Uranium Compounds, C. Voegtlin and 

H.C. Hodge (editors), McGraw-Hill Book Company, New York, Vol. I, pp. 309-376 

 Morrow, P.E., et al., 1982, Metabolic Fate and Evaluation of Injury in Rats and Dogs Following Exposure to 

the Hydrolysis Products of Uranium Hexafluoride, NUREG/CR-2268, prepared by University of Rochester, 

Department of Radiation Biology and Biophysics, Rochester, N.Y., for U.S. Nuclear Regulatory Commission, 

Office of Nuclear Regulatory Research, Washington, D.C., December 

 National Research Council, 1988, Health Risks of Radon and Other Internally Deposited Alpha-Emitters, 

BEIR IV Report, report of the Committee on the Biological Effects of Ionizing Radiations, National Academy 

Press, Washington, D.C. 

 Radionuclide-specific information was obtained from the following Internet addresses: 

http://www.britannica.com/seo/u/uranium/ 

http://www.atsdr.cdc.gov/ToxProfiles/phs9029.html 

Depleted 

uranium (DU) 

ATSDR (Agency for Toxic Substances and Disease Registry), 1999, Toxicological Profile for Uranium. 

Update, U.S. Department of Health and Human Services, Public Health Service, September 

 EPA (U.S. Environmental Protection Agency), 1992, Overview of Air Pathway Assessments for Superfund 

Sites (Revised), Interim Final, Office of Air Quality Planning and Standards, Research Triangle Park, N.C. 

Air/Superfund National Technical Guidance Study Series, Report ASF-1a, Vol. I. November 

 EPA (U.S. Environmental Protection Agency), 1996, Soil Screening Guidance, Technical Background 

Document, Appendix B: Route-to-Route Extrapolation of Inhalation Benchmarks, Office of Solid Waste and 

Emergency Response, Washington, D.C. EPA/540/R-95/128, July 
DU (cont’d) EPA (U.S. Environmental Protection Agency), 2005, Integrated Risk Information System (IRIS). Online 

database, Substance file for uranium, soluble salts, Office of Research and Development, National Center for 

Environmental Assessment, http://www.epa.gov/ngispgm3/iris/index.html, August 

 Fisher, D. R., et al., 1994, Uranium Hexafluoride Public Risk, Letter Report, PNL-10065, Pacific Northwest 

Laboratory, Health Protection Department, Richland, Washington, August 

 Gilman, A. P., et al., 1998, Uranyl nitrate — 91-Day Toxicity Studies in the New Zealand White Rabbit. 

Toxicol. Sci. 41(1):129-137, January 

 ICRP (International Commission on Radiological Protection), 1994, Human Respiratory Tract Model for 

Radiological Protection, ICRP Publication 66, Pergamon Press, Oxford, United Kingdom 

 Maynard, E. A., and Hodge, H. C., 1949, Studies of the Toxicity of Various Uranium Compounds when Fed 

to Experimental Animals. In: Pharmacology and Toxicology of Uranium Compounds, National Nuclear 

Energy Series (VI), pp. 309-376 (Voegtlin, I. C., and Hodge, H. C., Eds.), New York, McGraw-Hill 

 McGuire, S. A. 1991, Chemical Toxicity of Uranium Hexafluoride Compared to Acute Effects of Radiation,

NUREG-1391, Final Report. U.S. Nuclear Regulatory Commission, Office of Nuclear Regulatory Research, 

Washington, D.C., February 

 U.S. Nuclear Regulatory Commission, 1994, 10 CFR Part 19, Certification of Gaseous Diffusion Plants Final 

Rule, Discussion on Section 76.85, “Assessment of Accidents,” Fed. Reg. 59(184):48944. September 23 

 U.S. Nuclear Regulatory Commission, 2001, 10 CFR Part 20, Standards for Protection Against Radiation, 

Appendix B: Annual Limits on Intake (ALIs) and Derived Air Concentrations (DACs) of Radionuclides for 

Occupational Exposure; Effluent Concentrations; Concentrations for Release to Sewerage 

Zirconium Radionuclide-specific information was obtained from the following Internet addresses: 

http://encarta.msn.com/index/conciseindex/53/053DF000.htm?z=1&pg=2&br=1 

http://www.amm.com/ref/zirc.htm 

http://search.eb.com/bol/search?type=topic&&query=Zirconium&Dbase=Articles 

http://www.ornl.gov/isotopes/s_zr.html 

a This table summarizes key sources of information used to prepare the radiological fact sheets. 

CFR = Code of Federal Regulations.