Atomic Number 8 Element Facts Atomic Number 8 Element Facts Atomic Number 8 Element Facts What Element is Atomic Number 8? Oxygen is atomic number 8 on the periodic table. Each oxygen atom has 8 protons. ROGER HARRIS/SCIENCE PHOTO LIBRARY / Getty Images By Anne Marie Helmenstine, Ph.D. Updated April 03, 2016. Oxygen, element symbol O, is the element that is atomic number 8 on the periodic table. This means every atom of oxygen has 8 protons. Varying the number of electrons forms ions, while altering the number of neutrons makes different isotopes of the element, but the number of protons remains constant. Here is a collection of interesting facts about atomic number 8. Atomic Number 8 Element Facts While oxygen is a colorless gas under ordinary conditions, element 8 is actually quite colorful ! Liquid oxygen is blue, while the solid element may be blue, pink, orange, red, black, or even metallic. Oxygen is a nonmetal belonging to the chalcogen group . It is highly reactive and readily forms compounds with other elements.

It is found as a pure element in nature as oxygen gas (O2) and ozone (O3). Excited oxygen atoms produce the green and red colors of the aurora . Although air consists mainly of nitrogen, atomic number 8 is responsible for most of the colors that we see. continue reading below our video Learn About Molecules and Moles Today, oxygen makes up about 21% of the Earth's atmosphere . However, air was not always so highly oxygenated! A 2007 NASA-funded study determined oxygen has been present in air for about 2.3 billion to 2.4 billion years, with levels starting to rise 2.5 billion years ago. Photosynthetic organisms, such as plants and algae, are responsible for maintaining the high oxygen levels necessary for life. Although hydrogen atoms are the most numerous type of atom in the human body , oxygen accounts for about two-thirds of the mass of most living organisms, mainly because cells contain a lot of water. 88.9% of the weight of water comes from oxygen.

Swedish pharmacist Carl Wilhelm Scheele, French chemist Antoine Laurent Lavoisier, and British scientists and clergyman Joseph Priestly researched and discovered oxygen between 1770 and 1780. Lavoisier first called element number 8 by the name "oxygen" in 1777. Oxygen is the third most abundant element in the universe . The element is made by stars around 5x more massive than the Sun when they reach the point where they burn carbon or a combination of helium in carbon in fusion reactions. Over time, the abundance of oxygen in the universe will increase. Until 1961, atomic number 8 was the standard for the atomic weight of the chemical elements. In 1961, the standard was switched over to carbon-12. It's a common misconception that hyperventilation is caused by breathing in too much oxygen. Actually, hyperventilating is caused by exhaling too much carbon dioxide. Although carbon dioxide can be toxic at high levels, its needed in the blood to prevent it from becoming too alkaline.

Breathing too quickly causes blood pH to rise, which constricts blood vessels in the brain, leading to headache, slurred speach, dizziness, and other symptoms. Essential Element 8 Information State of Matter at Room Temperature: Gas Atomic Weight: 15.9994 Density: 0.001429 grams per cubic centimeter Isotopes: At least 11 isotopes of oxygen exist. 3 are stable. Most Common Isotope: Oxygen-16 (accounts for 99.757% of the natural abundance) Melting Point: -218.79 °C Atomic oxygen | Article about Atomic oxygen by The Free Dictionary Atomic oxygen | Article about Atomic oxygen by The Free Dictionary Related to Atomic oxygen: Dioxygen , Diatomic Oxygen oxygen, gaseous chemical element; symbol O; at. no. 8; interval in which at. wt. ranges 15.99903–15.99977; m.p. −218.4&degC;; b.p.

−182.962&degC;; density 1.429 grams per liter at STP; valence −2. The existence and properties of oxygen had been noted by many scientists before the announcement of its isolation by Priestley in 1774. Scheele had also succeeded in preparing oxygen from a number of substances, but publication of his findings was delayed until after that of Priestley's. As a result, Priestley and Scheele are credited with the discovery of the element independently. The fact that the gas is a component of the atmosphere was finally and definitely established by Lavoisier a few years later. In 1929, W. F. Giaque and H. L. Johnston announced the discovery of two isotopes of oxygen, of mass numbers 17 and 18. Properties and Compounds Oxygen is a colorless, odorless, tasteless gas; it is the first member of Group 16 of the periodic table periodic table, chart of the elements arranged according to the periodic law discovered by Dmitri I. Mendeleev and revised by Henry G. J. Moseley.

In the periodic table the elements are arranged in columns and rows according to increasing atomic number (see the table entitled ..... Click the link for more information. . It is denser than air and only slightly soluble in water. A poor conductor of heat and electricity, oxygen supports combustion but does not burn. Normal atmospheric oxygen is a diatomic gas (O2) with molecular weight 31.9988. Ozone ozone , an allotropic form of the chemical element oxygen (see allotropy). Pure ozone is an unstable, faintly bluish gas with a characteristic fresh, penetrating odor. The gas has a density of 2.144 grams per liter at STP. ..... Click the link for more information. is a highly reactive triatomic (O3) allotrope of oxygen (see allotropy allotropy [Gr.,=other form]. A chemical element is said to exhibit allotropy when it occurs in two or more forms in the same physical state; the forms are called allotropes. ..... Click the link for more information. ). When cooled below its boiling point oxygen becomes a pale blue liquid; when cooled still further the liquid solidifies, retaining its color.

Oxygen is paramagnetic in its solid, liquid, and gaseous forms. Although eight isotopes of oxygen are known, atmospheric oxygen is a mixture of the three isotopes with mass numbers 16, 17, and 18. Oxygen is extremely active chemically, forming compounds with almost all of the elements except the inert gases. Oxygen unites directly with a number of other elements to form oxides oxide, chemical compound containing oxygen and one other chemical element. Oxides are widely and abundantly distributed in nature. Water is the oxide of hydrogen. Silicon dioxide is the major component of sand and quartz. ..... Click the link for more information. . It is a constituent of many acids and of hydroxides, carbohydrates, proteins, fats and oils, alcohols, cellulose, and numerous other compounds such as the carbonates, chlorates, nitrates and nitrites, phosphates and phosphites, and sulphates and sulphites. The common reaction in which it unites with another substance is called oxidation (see oxidation and reduction oxidation and reduction, complementary chemical reactions characterized by the loss or gain, respectively, of one or more electrons by an atom or molecule. Originally the term oxidation .....

Click the link for more information. ). The burning of substances in air is rapid oxidation or combustion combustion, rapid chemical reaction of two or more substances with a characteristic liberation of heat and light; it is commonly called burning. The burning of a fuel (e.g., wood, coal, oil, or natural gas) in air is a familiar example of combustion. ..... Click the link for more information. . The respiration respiration, process by which an organism exchanges gases with its environment. The term now refers to the overall process by which oxygen is abstracted from air and is transported to the cells for the oxidation of organic molecules while carbon dioxide (CO2 ..... Click the link for more information. of animals and plants is a form of oxidation essential to the liberation of the energy stored in such food materials as carbohydrates and fats. The rusting of iron and the corrosion of many metals results from the action of the oxygen in the air. Natural Occurrence and Preparation Oxygen is the most abundant element on earth, constituting about half of the total material of its surface. Most of this oxygen is combined in the form of silicates, oxides and water.

It makes up about 90% of water, two thirds of the human body and one fifth by volume of air. It is found in the sun, and has a role in the stellar carbon cycle (see nucleosynthesis nucleosynthesis or nucleogenesis, in astronomy, production of all the chemical elements from the simplest element, hydrogen, by thermonuclear reactions within stars, supernovas, and in the big bang at the beginning of the universe (see nucleus; nuclear energy). ..... Click the link for more information. ). Oxygen is prepared for commercial use by the liquefaction and fractional distillation of air and more expensively by the electrolysis of water; it is stored and transported under high pressure in steel cylinders. It can also be obtained by heating certain of its compounds, such as barium peroxide, potassium chlorate, and the red oxide of mercury. Uses Oxygen is of great importance in the chemical and the iron and steel industries. Its major use is in steel production, for example in the Bessemer process Bessemer process [for Sir Henry Bessemer], industrial process for the manufacture of steel from molten pig iron.

The principle involved is that of oxidation of the impurities in the iron by the oxygen of air that is blown through the molten iron; the heat of oxidation raises the ..... Click the link for more information. . The oxyacetylene torch oxyacetylene torch , tool that mixes and burns oxygen and acetylene to produce an extremely hot flame. This torch can be used for cutting steel and for welding iron and various other metals. The temperature of the flame can reach as high as 6,300&degF; (3,480&degC;). ..... Click the link for more information. is another important industrial application. Oxygen is utilized in medicine in the treatment of respiratory diseases and is mixed with other gases for respiration in submarines, high-flying aircraft, and spacecraft. Liquid oxygen is used as an oxidizer in the fuel systems of large rockets. Oxygen was formerly the official standard for the atomic weights atomic weight, mean (weighted average) of the masses of all the naturally occurring isotopes of a chemical element, as contrasted with atomic mass, which is the mass of any individual isotope. Although the first atomic weights were calculated at the beginning of the 19th cent. .....

Click the link for more information. of elements. The chemists used natural oxygen, a mixture of three isotopes, to which the value of 16 was assigned while the physicists assigned the value of 16 specifically to the oxygen isotope 16. In 1961 carbon-12 replaced oxygen as the standard. Oxygen O, a chemical element of group VI of Mendeleev’s periodic system. Atomic number, 8; atomic mass, 15.9994. Under normal conditions, oxygen is a colorless, tasteless, and odorless gas. It would be difficult to name another element that plays as important a role on earth as oxygen. History. The processes of combustion and respiration have long attracted the attention of scientists. The first hints that not all of the air but only its “active” part supports combustion were found in Chinese eighth-century manuscripts. Much later, Leonardo da Vinci (1452–1519) regarded air as a mixture of two gases, of which only one is consumed during cumbustion and respiration.

The ultimate discovery of the two principal constituents of air, nitrogen and oxygen, a discovery of epochal importance in science, was made only in the late 18th century. Oxygen was prepared almost simultaneously by K. Scheele (1769–70) by calcination of nitrates (KNO3, NaNO3), manganese dioxide (MnO2), and other substances and by J. Priestley (1774) by the heating of red lead oxide, Pb3O4, and mercuric oxide, HgO. D. Rutherford discovered nitrogen in 1772. A. Lavoisier, performing a quantitative analysis of air, found that it “consists of two [gases] of different and, so to speak opposite character,” namely of oxygen and nitrogen. On the basis of extensive experimental studies, he correctly interpreted combustion and respiration as processes of the interaction of various substances and oxygen.

Since oxygen in a constituent of acids, Lavoisier called it oxygène, that is, “acid former” (from the Greek oxys [acidic]) and gennao [I generate]; hence the Russian name kislorod, from kislyi [acidic] and rodif [to generate]. Distribution in nature. Oxygen is the most widespread element on earth. Combined oxygen constitutes about six-sevenths of the earth’s hydrosphere (85.82 percent by weight) and almost one-half of the lithosphere (47 percent by weight). Only in the atmosphere, where oxygen is present in the free state, is it the second most abundant element (23.15 percent by weight), after nitrogen. Oxygen is also in first place in the number of minerals formed by it (1,364). The most widespread minerals containing oxygen are silicates (feldspars, micas), quartz, iron oxides, carbonates, and sulfates. Living organisms on the average contain about 70 percent oxygen; it is a constituent of most of the important organic compounds (proteins, fats, carbohydrates) and a component of the inorganic constituents of the skeleton.

The role of free oxygen is especially important in biochemical and physiological processes, particularly in respiration. With the exception of some anaerobic microorganisms, all animals and plants receive the energy required by life processes from the biological oxidation of various substances with oxygen. The entire mass of the free oxygen on earth originated from and is being maintained by the life activities of green plants, which evolve oxygen in the course of photosynthesis, on land and in the oceans. Photosynthesis and the dominance of free oxygen on the earth’s surface give rise to strongly oxidizing conditions. Conversely, reducing media are formed where oxygen is absent, as in magma, in deep levels of subterranean water, in the silt of the oceans and lakes, and in marshes. Reduction-oxidation processes with the participation of oxygen determine the concentration of many elements and the formation of such mineral deposits as coal, oil, sulfur, iron, and copper ores. The oxygen cycle is also changed by the economic activities of man. In some industrial countries more oxygen is consumed during fuel combustion than the amount evolved by the photosynthesis of plants. The annual oxygen consumption for combustion of fuels on the earth is about 9 X 199 tons. Isotopes, atoms, and molecules.

Oxygen has three stable isotopes—16O, 17O, and 18O—the average content of which constitutes, respectively, 99.759 percent, 0.037 percent, and 0.204 percent of the total number of oxygen atoms on the earth. The predominance of the lightest isotope 16O in the isotope mixture is due to the 16O nucleus consisting of eight protons and eight neutrons. The theory of the atomic nucleus indicates that such nuclei are particularly stable. In accordance with oxygen’s position in Mendeleev’s periodic system of elements, the electrons of the oxygen atom are arranged in two shells: two in the inner shell and six in the outer shell (configuration Is22s22p4). Since the outer shell is not filled and the ionization potential and the electron affinity are 12.61 and 1.46 eV, respectively, the oxygen atom usually acquires electrons in the course of formation of chemical compounds and has a negative effective charge.

Conversely, rare are the compounds in which electrons are torn away (or more precisely, pulled away) from the oxygen atom (as in F2O and F2O2). In the past, proceeding solely from the position of oxygen in the periodic system, a negative charge ( — 2) was ascribed to the oxygen atom. However, experimental data have indicated that the ion O2- does not exist either in the free state or in compounds and that the effective negative charge of the oxygen atom practically never exceeds unity. Under ordinary conditions, the oxygen molecule is diatomic (O2). A triatomic molecule, ozone (O3), is formed in a silent electric discharge; small amounts of tetratomic molecules (O4) have been detected at high pressures. The electronic structure of O2 is of great theoretical interest. The molecule has two unpaired electrons in the ground state; the usual “classical” structural formula O ⇄ O with two two-electron bonds is inapplicable to it. An exhaustive explanation of this fact is provided by the theory of molecular orbitals.

The ionization energy of the oxygen molecule (O2 — e → O+2) constitutes 12.2 eV, and the electron affinity (O2 + e → O-2)> 0.94 eV. The dissociation of molecular oxygen into atoms is negligible at ordinary temperatures, becoming noticeable only at 1500°C; at 5000°C the oxygen molecules are almost completely dissociated into atoms. Physical properties. Oxygen is a colorless gas that liquifies at — 182.9°C and normal pressure to a pale blue liquid, which, in turn, solidifies at — 218.7°C forming blue crystals. The density of gaseous oxygen (at 0°C and normal pressure) is 1.42897 g/l. The critical temperature of oxygen is fairly low (tCrit = -118.84°C), that is, lower than that of Cl2, CO2, SO2, and some other gases; Pcrit = 4.97 meganewtons per square meter (49.71 atm).

The heat conductivity (at 0°C) is 23.86 X 10-3 W/(m.°K), or 57 X 10-6 cal/(sec · cm.°C). The molar heat capacities (at 0°C) are Cp = 28.9 and Cv = 20.5 in joules (mole·°K) and Cp = 6.99 and Cv = 4.98 in cal/(mole· °C), respectively; CP/CV = 1.403. The dielectric permeability of gaseous oxygen is 1.000547 (at 0°C), and of liquid oxygen 1.491. The viscosity is 189 millipoises (at 0°C). Oxygen is poorly soluble in water: at 20°C and 1 atm, 0.031 m3 dissolves in 1 m3 of water; 0.049 m3 dissolves at 0°C. Platinum black and activated charcoal are efficient solid absorbents for oxygen. Chemical properties.

Oxygen forms chemical compounds with all other elements except the light inert gases. Being the most active nonmetal (after fluorine), oxygen interacts directly with most elements. The only exceptions are the heavy inert gases, the halogens, gold, and platinum; their compounds with oxygen are obtained by indirect methods. Almost all of the reactions involving oxygen are exothermic oxidation reactions, that is, accompanied by the evolution of heat. Oxygen reacts with hydrogen at ordinary temperatures very slowly, whereas this reaction proceeds explosively above 550°C: 2H2 + O2 = 2H2O. Oxygen reacts with sulfur, carbon, nitrogen, and phosphorus very slowly under ordinary conditions. The reaction rate increases with increasing temperature until at an ignition temperature characteristic for each element combustion occurs. The reaction of oxygen with nitrogen is endothermic because of the particular stability of the N2 molecule and becomes noticeable only above 1200°C or in an electric discharge: N2 + O2 = 2NO. Oxygen actively oxidizes almost all metals and, with particular ease, the alkali and alkaline-earth metals.

The reactivity of a metal with oxygen depends on many factors, such as the condition of the metal surface, the degree of subdivision, and the presence of impurities. The role of water is of particular importance in the interaction of substances with oxygen. For example, such an active metal as potassium does not react at all with oxygen that is completely devoid of moisture, but it ignites in oxygen at room temperature in the presence of the minutest quantities of water vapor. It has been calculated that the annual loss owing to corrosion amounts to as much as 10 percent of the entire metal production. Oxides of some metals form peroxides by the addition of oxygen. The resulting compounds contain two or more oxygen atoms bonded to each other. Thus, the peroxides Na2O2 and BaO2 contain the peroxide ion O22-, the superoxides NaO2 and KO2 contain the ion O2-, and the ozonides NaO3, KO3, RbO3, and CsO3 contain the ion O3-. Oxygen reacts exothermally with numerous complex substances.

Thus, ammonia burns in oxygen in the absence of catalysts, the reaction proceeding according to the equation 4NH3 X 3O2 = 2N2 X 6H2O. Oxidation of ammonia with oxygen in the presence of catalysts yields NO (this process is used in the production of nitric acid). Of particular importance is the combustion of hydrocarbons (natural gas, gasoline, kerosene), which constitutes the most important heat source for consumers and industry; for example, CH4 + 2O2 = CO2 + 2H2O. The reaction of hydrocarbons with oxygen forms the basis of many important industrial processes, such as the reforming of methane, which is used for the production of hydrogen: 2CH4 + O2 + 2H2O = 2CO2 + 6H2. Many organic compounds (hydrocarbons with double and triple bonds, aldehydes, phenols, turpentine, drying oils) vigorously add oxygen. The oxidation of nutrients with oxygen in the cells serves as a source of energy for living organisms. Preparation.

There are three basic methods for preparing oxygen: chemical, electrolytic (electrolysis of water), and physical (separation of air). The chemical method was discovered first. Oxygen may be prepared, for example, from potassium chlorate, KClO3, which decomposes on heating with evolution of O2 in amounts of 0.27 m3 per 1 kg of the salt. Barium oxide, BaO, absorbs oxygen at first, when heated up to 540°C, to give the peroxide BaO2, which decomposes on further heating to 870°C with the evolution of pure oxygen. It may be also obtained from KMnO4, Ca2PbO4, K2Cr2O7, and other substances by heating in the presence of catalysts. The chemical method of preparing oxygen is inefficient and expensive and is used only in laboratory practice. The electrolytic method consists of passing a direct electric current through water containing a solution of sodium hydroxide, NaOH, to increase its conductivity. In this case, water is decomposed into oxygen and hydrogen. The oxygen is collected at the positive electrode of the electrolysis unit, and the hydrogen at the negative electrode.

Oxygen is obtained in this method as a byproduct of hydrogen manufacture. The expenditure of 12–15 kW-hr of electric energy is required for the production of 2 m3 of hydrogen and 1 m3 of oxygen. The principal method for the production of oxygen in modern technology is the separation of air. To separate air in its normal gaseous state is very difficult and, therefore it is first liquified and then separated into components. This method for the production of oxygen is known as the low-temperature method of air separation. The air is first compressed with a compressor and then, after passing through heat exchangers, expanded in an expansion turbine or through a nozzle; this leads to its cooling to 93°K (—180°C) and conversion to liquid air. Further separation of liquid air, consisting primarily of liquid nitrogen and liquid oxygen, is based on the boiling point differences of its components (the boiling point of O2 is 90.18°K [-182.9°C], and that of N2, 77.36°K [-195.8°C]).

Gradual evaporation of liquid air leads, at first, to the evaporation of mainly nitrogen, and the remaining liquid becomes ever more enriched in oxygen. Continuous repetition of this process on the rectification plates of the air-separator columns yields liquid oxygen of the required purity (concentration). The smallest (capacity of several liters) and largest oxygen air-separation units (35,000 m3/hr of oxygen) in the world have been constructed in the USSR. The units are used in the production of industrial oxygen with a concentration of 95–98.5 percent, technical-grade oxygen with a concentration of 99.2–99.9, and the higher purity oxygen used in medicine. The product may be gaseous or liquified oxygen. The expenditure of electric energy ranges from 0.41 to 1.6 kW-hr/m3. It is also possible to prepare oxygen by separating air using the method of selective permeation (diffusion) through membrane barriers. Air under pressure is passed through barriers made of fluorocarbons, glass, or plastics, the structural lattices of which are capable of allowing some components to pass through while retaining others.

This method of oxygen production has been used to date (1973) only in laboratories. Gaseous oxygen is stored and transported in steel cylinders and receivers at pressures of 15 and 42 meganewtons per square meter (corresponding to 150 and 420 bars, or 150 and 420 atm, respectively). Liquid oxygen is stored and shipped in metal Dewar vessels or in special tanks. Special pipelines are also used for transporting liquid and gaseous oxygen. Oxygen cylinders are painted light blue and are labeled “oxygen” in black. Uses. Technical-grade oxygen is used in the gaseous flame working of metals, in welding, in oxygen cutting, in case hardening, in metallization, and in other processes, as well as in aviation, in underwater navigation, and elsewhere. Industrial oxygen is used in the chemical industry for the production of such products as synthetic liquid fuels, lubricating oils, nitric and sulfuric acids, methanol, ammonia, ammonia fertilizers, and peroxides of metals. Liquid oxygen is used in working with explosives, in rocket engines, and in laboratory applications as a cooling agent.

Pure cylinder oxygen is used for breathing at high altitudes, during space flights, and in underwater navigation. In medicine, oxygen is administered intramuscularly and to aid the respiration of the gravely ill; it is also used in oxygen, water, and air baths (in oxygen tents). V. L. V Oxygen | Define Oxygen at Dictionary.com oxygen [ok-si-juh n] /ˈɒk sɪ dʒən/ Spell noun, Chemistry. 1. a colorless, odorless, gaseous element constituting about one-fifth of the volume of the atmosphere and present in a combined state in nature. It is the supporter of combustion in air and was the standard of atomic, combining, and molecular weights until 1961, when carbon 12 became the new standard. Symbol: O; atomic weight: 15.9994; atomic number: 8; density: 1.4290 g/l at 0°C and 760 mm pressure.

Origin of oxygen 1780-90; < French oxygène, equivalent to oxy- oxy- 1 + -gène -gen Related forms [ok-si-jen-ik] /ˌɒk sɪˈdʒɛn ɪk/ (Show IPA), oxygenous [ok-sij-uh-nuh s] /ɒkˈsɪdʒ ə nəs/ (Show IPA), adjective oxygenicity [ok-si-juh-nis-i-tee] /ˌɒk sɪ dʒəˈnɪs ɪ ti/ (Show IPA), noun Dictionary.com Unabridged Examples from the Web for oxygen Expand Contemporary Examples The answer lies in polyphenol oxidase, an enzyme that combines with oxygen to speed up cellular decomposition. Aether and Gravitation William George Hooper It is as though one were to wish a fish to live out of water, or a mammal in an atmosphere containing no oxygen. British Dictionary definitions for oxygen Expand noun 1. a colourless odourless highly reactive gaseous element: the most abundant element in the earth's crust (49.2 per cent).

It is essential for aerobic respiration and almost all combustion and is widely used in industry. Symbol: O; atomic no: 8; atomic wt: 15.9994; valency: 2; density: 1.429 kg/m³; melting pt: –218.79°C; boiling pt: –182.97°C (as modifier): an oxygen mask Derived Forms oxygenic (ˌɒksɪˈdʒɛnɪk), oxygenous (ɒkˈsɪdʒɪnəs) adjective Collins English Dictionary - Complete & Unabridged 2012 Digital Edition © William Collins Sons & Co. Ltd. 1979, 1986 © HarperCollins Publishers 1998, 2000, 2003, 2005, 2006, 2007, 2009, 2012 Word Origin and History for oxygen Expand n.

gaseous chemical element, 1790, from French oxygène, coined in 1777 by French chemist Antoine-Laurent Lavoisier (1743-1794), from Greek oxys "sharp, acid" (see acrid ) + French -gène "something that produces" (from Greek -genes "formation, creation;" see -gen ). Intended to mean "acidifying (principle)," it was a Greeking of French principe acidifiant. So called because oxygen was then considered essential in the formation of acids (it is now known not to be). The element was isolated by Priestley (1774), who, using the old model of chemistry, called it dephlogisticated air. The downfall of the phlogiston theory required a new name, which Lavoisier provided. Online Etymology Dictionary, © 2010 Douglas Harper oxygen ox·y·gen (ŏk'sĭ-jən) n.

Symbol O An element constituting 21 percent of the atmosphere by volume that occurs as a diatomic gas, O2, combines with most elements, is essential for plant and animal respiration, and is required for nearly all combustion. Atomic number 8; atomic weight 15.9994; melting point -218.8°C; boiling point -183.0°C; gas density at 0°C 1.429 grams per liter; valence 2. A medicinal gas containing not less than 99.0 percent, by volume, of O2. The American Heritage® Stedman's Medical Dictionary Copyright © 2002, 2001, 1995 by Houghton Mifflin Company. Published by Houghton Mifflin Company. (ŏk'sĭ-jən) Symbol O A nonmetallic element that exists in its free form as a colorless, odorless gas and makes up about 21 percent of the Earth's atmosphere. It is the most abundant element in the Earth's crust and occurs in many compounds, including water, carbon dioxide, and iron ore.

Oxygen combines with most elements, is required for combustion, and is essential for life in most organisms. Atomic number 8; atomic weight 15.9994; melting point -218.4°C; boiling point -183.0°C; gas density at 0°C 1.429 grams per liter; valence 2. See Periodic Table . Our Living Language : In 1786, the French chemist Antoine Lavoisier coined a term for the element oxygen (oxygène in French). He used Greek words for the coinage: oxy- means "sharp," and -gen means "producing." Oxygen was called the "sharp-producing" element because it was thought to be essential for making acids. Lavoisier also coined the name of the element hydrogen, the "water-producing" element, in 1788. Soon after, in 1791, another French chemist, J. A. Chaptal, introduced the word nitrogen, the "niter-producing" element, referring to its discovery from an analysis of nitric acid. The American Heritage® Science Dictionary Copyright © 2002.

Published by Houghton Mifflin. All rights reserved. Basic tomic theory The word "atom" comes from the Greek a-tomos, meaning "un-cut-able" Throughout most of history the idea that matter is composed of minute particles had languished as a philosophical abstraction known as atomism , and no clear relation between these "atoms" and the chemical "elements" had been established. This began to change in the early 1800's when the development of balances that permitted reasonably precise measurements of the weight changes associated with chemical reactions ushered in a new and fruitful era of experimental chemistry. This resulted in the recognition of several laws of chemical change that laid the groundwork for the atomic theory of matter: conservation of mass . Laws of chemical change Recall that a "law", in the context of science, is just a relationship, discovered through experimentation, that is sufficiently well established to be regarded as beyond question for most practical purposes. Because it is the nature of scientists to question the "unquestionable", it occasionally happens that exceptions do arise, in which case the law must undergo appropriate modification. Conservation of mass-energy This is usually considered the most fundamental of law of nature.

It is also a good example of a law that had to be modified; it was known simply as Conservation of Mass until Einstein showed that energy and mass are interchangeable. However, the older term is perfectly acceptable within the field of ordinary chemistry in which energy changes are too small to have a measurable effect on mass relations. Within the context of chemistry, conservation of mass can be thought of as "conservation of atoms". Chemical change just shuffles them around into new arrangements. Mass conservation had special significance in understanding chemical changes involving gases, which were for some time not always regarded as real matter at all. (Owing to their very small densities, carrying out actual weight measurements on gases is quite difficult to do, and was far beyond the capabilities of the early experimenters.) Thus when magnesium metal is burned in air, the weight of the solid product always exceeds that of the original metal, implying that the process is one in which the metal combines with what might have been thought to be a "weightless" component of the air, which we now know to be oxygen. More importantly, as we will see later, this experimental result tells us something very important about the mass of the oxygen atom relative to that of the magnesium atom.

Law of definite proportions This law is also known as the law of constant composition. It states that the proportion by weight of the element present in any pure substance is always the same. This enables us to generalize the relationship we illustrated above. Problem Example 1 How many kilograms of metallic magnesium could theoretically be obtained by decomposing 0.400 kg of magnesium oxide into its elements? Solution: The mass ratio of Mg to O in this compound is 1/1.66 = 0.602, so 0.400 kg of the oxide contains (0.400 kg) x 0.602 = 0.241 kg of Mg. The fact that we are concerned with the reverse of the reaction cited above is irrelevant. The laws of definite and of multiple proportions, along with conservation of mass, are known collectively as the laws of chemical composition .2857 4 Next, identify the smallest weight ratio (.0714 here), and, in the rightmost column, divide each of the values in column 3 by this smallest ratio. What you get are the “small whole numbers” that illustrate the law of simple multiple proportions.

Note that, in some cases, the values in column 4 will not all initially work out to integer numbers. Suppose, for example, that the bottom row refers to some compound other than ammonium chloride, and that the numbers in column 4 (reading from top to bottom), work out to {2 - 1- 1.33}. If you recognize that the last number corresponds to 4/3, then multiplying each value by 3 yields the whole numbers {6 3 4} — not quite as “small”, but at least”whole” as the law requires. How Dalton's interpretation of the laws of chemical change established the atomic theory The idea that matter is composed of tiny "atoms" of some kind had been around for at least 2000 years. Dalton's accomplishment was to identify atoms with actual chemical elements.

If Nobel prizes had existed in the early 1800's, the English schoolteacher/meteorologist/chemist John Dalton (1766-1844) would certainly have won one for showing how the experimental information available at that time, as embodied in the laws of chemical change that we have just described, are fully consistent with the hypothesis that atoms are the smallest units of chemical identity. These points of Dalton's atomic theory provided satisfactory explanations of all the laws of chemical change noted above: Explanation of the law of conservation of mass This is really a consequence of "conservation of atoms" which are presumed to be indestructible by chemical means. In chemical reactions, the atoms are simply rearranged, but never destroyed. Explanation of the law of constant composition If compounds are made up of definite numbers of atoms, each of which has its own characteristic mass, then the relative mass of each element in a compound must always be the same. Thus the elements must always be present in a pure sample of a compound in the same proportions by mass. Explanation of the law of multiple proportions A given set of elements can usually form two or more compounds in which the numbers of atoms of some of the elements are different.

Because these numbers must be integers (you can't have "half" an atom!), the mass of one element combined with a fixed mass of any other elements in any two such compounds can differ only by integer numbers. Thus, for the series of nitrogen-hydrogen compounds cited in the Problem Example above, we have the following relations: Compound (AFM), an exquisitely sensitive electromechanical device in which the distance between the tip of a submicroscopic wire probe and the surface directly below it is recorded as the probe moves along a surface to which atoms are adsorbed. Right: Schematic diagram of an atomic force microscope using beam deflection detection. As the cantilever is displaced as it is drawn across a surface, so too will the reflection of the laser beam be displaced on the surface of the photodiode. (Wikimedia Commons) The general principle of the AFM is quite simple, but its realization in an actual device can appear somewhat intimidating! This highly specialized atomic force microscope is one of several similar devices described at the Argonne National Laboratory site → which carries a positive electric charge whose value identifies each element and is known as the atomic number of the element.

• Almost all of the volume of an atom consists of empty space in which electrons, the fundamental carriers of negative electric charge, reside. The extremely small mass of the electron (1/1840 the mass of the hydrogen nucleus) causes it to behave as a quantum particle , which means that its location at any moment cannot be specified; the best we can do is describe its behavior in terms of the probability of its manifesting itself at any point in space. It is common (but somewhat misleading) to describe the volume of space in which the electrons of an atom have a significant probability of being found as the electron cloud. The latter has no definite outer boundary, so neither does the atom. The radius of an atom must be defined arbitrarily, such as the boundary in which the electron can be found with 95% probability. Atomic radii are typically 30-300 pm. The nucleus is itself composed of two kinds of particles. Protons are the carriers of positive electric charge in the nucleus; the proton charge is exactly the same as the electron charge, but of opposite sign.

This means that in any [electrically neutral] atom, the number of protons in the nucleus (often referred to as the nuclear charge) is balanced by the same number of electrons outside the nucleus. Because the electrons of an atom are in contact with the outside world, it is possible for one or more electrons to be lost, or some new ones to be added. The resulting electrically-charged atom is called an ion The other nuclear particle is the neutron . As its name implies, this particle carries no electrical charge. Its mass is almost the same as that of the proton. Most nuclei contain roughly equal numbers of neutrons and protons, so we can say that these two particles together account for almost all the mass of the atom. Atomic number What single parameter uniquely characterizes the atom of a given element? It is not the atom's relative mass, as we will see in the section on isotopes below. It is, rather, the number of protons in the nucleus, which we call the atomic number and denote by the symbol Z. Each proton carries an electric charge of +1, so the atomic number also specifies the electric charge of the nucleus. In the neutral atom, the Z protons within the nucleus are balanced by Z electrons outside it.

Moseley searched for a measurable property of each element that increases linearly with atomic number. He found this in a class of X-rays emitted by an element when it is bombarded with electrons. The frequencies of these X-rays are unique to each element, and they increase uniformly in successive elements. Mosely found that the square roots of these frequencies give a straight line when plotted against Z; this enabled him to sort the elements in order of increasing atomic number. Atomic numbers were first worked out in 1913 by Henry Moseley , a young member of Rutherford's research group in Manchester. You can think of the atomic number as a kind of serial number of an element, commencing at 1 for hydrogen and increasing by one for each successive element. The chemical name of the element and its symbol are uniquely tied to the atomic number; thus the symbol "Sr" stands for strontium, whose atoms all have Z = 38. Mass number This is just the sum of the numbers of protons and neutrons in the nucleus. It is sometimes represented by the symbol A, so A = Z + N in which Z is the atomic number and N is the neutron number The term nuclide simply refers to any particular kind of nucleus.

For example, a nucleus of atomic number 7 is a nuclide of nitrogen. Any nuclide is characterized by the pair of numbers (Z ,A). The element symbol depends on Z alone, so the symbol 26Mg is used to specify the mass-26 nuclide of manganese, whose name implies Z=12. A more explicit way of denoting a particular kind of nucleus is to add the atomic number as a subscript. Of course, this is somewhat redundant, since the symbol Mg always implies Z=12, but it is sometimes a convenience when discussing several nuclides. Know the difference between mass number and atomic mass! (TylerDeWitt, 9 min) Because it is not always easy to display a subscript directly beneath a superscript, it is not uncommon to use constructions such as 12Mg26 , which will often be our practice in this document when it is necessary to show both Z and A explicitly. Isotopes Two nuclides having the same atomic number but different mass numbers are known as isotopes .

Most elements occur in nature as mixtures of isotopes, but twenty-three of them (including beryllium and fluorine, shown in the table) are monoisotopic. For example, there are three natural isotopes of magnesium: 24Mg (79% of all Mg atoms), 25Mg (10%), and 26Mg (11%); all three are present in all compounds of magnesium in about these same proportions. Approximately 290 isotopes occur in nature. The best place to find out about the isotopes of individual elements is this page at the Lawrence Berkeley site , which covers both the natural isotopes and the artificially-produced ones. The two heavy isotopes of hydrogen are especially important— so much so that they have names and symbols of their own: Deuterium accounts for only about 15 out of every one million atoms of hydrogen. Tritium, which is radioactive, is even less abundant. All the tritium on the earth is a by-product of the decay of other radioactive elements.

6 Relative atomic masses: the atomic weight scale For historical reasons, the term atomic weight has a special meaning in Chemistry; it does not refer to the actual "weight" of an atom, which would be expressed in grams or kg. Atomic weights, sometimes called relative weights, are more properly known as relative atomic masses, and being ratios, are dimensionless. See this Wikipedia article for a full discussion of these and related terms. Please note that although the terms mass and weight have different meanings, the differences between their values are so small as to be insignificant for most practical purposes, so the terms atomic weight and atomic mass can be used interchangeably. Atoms are of course far too small to be weighed directly; weight measurements can only be made on the massive (but unknown) numbers of atoms that are observed in chemical reactions. The early combining-weight experiments of Dalton and others established that hydrogen is the lightest of the atoms, but the crude nature of the measurements and uncertainties about the formulas of many compounds made it difficult to develop a reliable scale of the relative weights of atoms. Even the most exacting weight measurements we can make today are subject to experimental uncertainties that limit the precision to four significant figures at best.

Atomic weights are average relative masses In the earlier discussion of relative weights of atoms, we explained how Dalton assigned a relative weight of unity to hydrogen, the lightest element, and used combining weights to estimate the relative weights of the others he studied. Later on, when it was recognized that more elements form simple compounds with oxygen, this element was used to define the atomic weight scale. Selecting O = 16 made it possible to retain values very close to those already assigned on the H=1 scale. Finally, in 1961, carbon became the defining element of the atomic weight scale. But because, by this time, the existence of isotopes was known, it was decided to base the scale on one particular isotope of carbon, C-12, whose relative mass is defined as exactly 12.000. Because almost 99% of all carbon atoms on the earth consist of 6C12, atomic weights of elements on the current scale are almost identical to those on the older O=16 scale. Most elements possess more than one stable isotope in proportions that are unique to each particular element. For this reason, atomic weights are really weighted averages of the relative masses of each that are found on earth.

Atomic weights are the ratios of the average mass of the atoms of an element to the mass of an identical number of 6C12 atoms. You can visualize the atomic weight scale as a long line of numbers that runs from 1 to around 280. The beginning of the scale looks like this: Atomic weights (relative atomic masses) of the first ten elements The red vertical lines beneath each element symbol indicate where that element is located on the atomic weight scale. Of these ten elements, only two, beryllium and fluorine, have a single isotope. The other eight atomic weights are weighted averages of the relative masses of the multiple isotopes that these (and most) elements possess. This is especially noticeable in the case of boron, whose average relative mass falls between 10 and 11. (Historically, observations of cases such as these led to the very concept of isotopes.) For many elements, one particular isotope so dominates the natural mixture that the others have little effect on the average mass. For example, 99.99 percent of hydrogen atoms consist of 1H1, whereas 1H2, the other stable isotope, amounts to only 0.01 percent.

Similarly, oxygen is dominated by 8O16 (over 99.7 percent) to the near exclusion of its two other isotopes. Atomic weights are listed in tables found in every chemistry textbook; you can't do much quantitative chemistry without them! The "standard" values are updated every few years as better data becomes available. You will notice that the precisions of these atomic weights, as indicated by the number of significant figures, vary considerably. Atomic weights of the 26 elements having a single stable isotope (monoisotopic elements) are the most precisely known. Two of these, boron and fluorine, appear in the above table. Owing to geochemical isotopic fractionation (discussed farther on), there is always some uncertainty in averaging the atomic weights of elements with two or more stable isotopes. Industrial processes associated mainly with nuclear energy and weapons production require the isolation or concentration of particular isotopes. When the by-product elements or compounds from which these isotopes have been depleted eventually get distributed in the environment or sold into the commercial marketplace, their atomic weights can vary from "official" values. This has occurred, for example, with lithium, whose isotope Li-6 has been used to produce hydrogen bombs.

Naturally-occurring radioactive elements (all elements heavier than 82Pb) all gradually decay into lighter elements, most of which are themselves subject to radioactive decay. These radioactive decay chains eventually terminate in a stable element, the most common of which is one of the three stable isotopes of lead. Subsequent geochemical processes can cause lead ore bodies from such sources to mix with "primeval" Pb (derived from the cosmic dust that formed the solar system), leading to a range of possible average atomic weights. For these reasons, lead, with a listed average atomic weight of 207.2, has the least-certain mass of any stable element. Weighing atoms: mass spectrometry A major breakthrough in Chemistry occurred in 1913 when J.J. Thompson directed a beam of ionized neon atoms through both a magnetic field and an electrostatic field. Using a photographic plate as a detector, he found that the beam split into two parts, and suggested that these showed the existence of two isotopes of neon, now known to be Ne-20 and Ne-22.

This, combined with the finding made a year earlier by Wilhelm Wien that the degree of deflection of a particle in these fields is proportional to the ratio of its electric charge to its mass, opened the way to characterizing these otherwise invisible particles. Thompson’s student F.W. Aston improved the apparatus, developing the first functional mass spectrometer , and he went on to identify 220 of the 287 isotopes found in nature; this won him a Nobel prize in 1921. His work revealed that the mass numbers of all isotopes are nearly integers (that is, integer multiples of the mass number 1 of the protons and neutrons that make up the nucleus. Mass spectrometry begins with the injection of a vaporized sample into an ionization chamber where an electrical discharge causes it to become ionized. An accelerating voltage propels the ions through an electrostatic field that allows only those ions having a fixed velocity (that is, a given charge) to pass between the poles of a magnet. The magnetic field deflects the ions by an amount proportional to the charge-to-mass ratios. The separated ion beams are detected and their relative strengths are analyzed by a computer that displays the resulting mass spectrum.

In modern devices, a computer also controls the accelerating voltage and electromagnet current so as to being successive ion beams into focus on the detector. Neutral atoms, having no charge, cannot be accelerated along a path so as to form a beam, nor can they be deflected. They can, however, be made to acquire electric charges by directing an electron beam at them, and this was one of the major innovations by Aston that made mass spectrometry practical. The mass spectrometer has become one of the most widely used laboratory instruments. Mass spectrometry is now mostly used to identify molecules. Ionization usually breaks a molecule up into fragments having different charge-to-mass ratios, each molecule resulting in a unique "fingerprint" of particles whose origin can be deduced by a jigsaw puzzle-like reconstruction. For many years, "mass-spec" had been limited to small molecules, but with the development of novel ways of creating ions from molecules, it has now become a major tool for analyzing materials and large biomolecules, including proteins . The mass spectrum of magnesium shows that it consists of three isotopes of masses 24 through 26. The height of each peak shows the abundance of each isotope.

University lecture on MS (UC-Irvine, 55 min) Isotopic mixtures and abundances Only 26 of the elements that occur on the Earth exist as a single isotope; these are said to be monoisotopic . The remaining elements consist of mixtures of between two and ten isotopes. The total number of natural isotopes is 339; of these, 254 are stable, while the remainder are radioactive, meaning that they decay into stable isotopes. Recalling that a given isotope (also known as a nuclide ) is composed of protons and neutrons, each having a mass number of unity, it should be apparent that the mass number of a given nuclide will be an integer, as seen in the mass spectrum of magnesium above. It also follows that the relative atomic masses (“atomic weights”) of monoisotopic elements will be very close to integers, while those of other elements, being weighted averages, can have any value. Problem Example 4 Estimate the average atomic weight of magnesium from the isotopic abundance data shown in the above mass spectrometry plot.

Solution: We just take the weighted average of the mass numbers: (0.7899 × 24) + (0.1000 × 25) + (0.1101 × 26) = 24.32 Note: The measured atomic weight of Mg (24.305) is slightly smaller than this because atomic masses of nuclear components are not strictly additive, as will be explained further below. When there are only two significantly abundant isotopes, you can estimate the relative abundances from the mass numbers and the average atomic weight. The following is a favorite exam problem: Problem Example 5 The average atomic weight of chlorine is 35.45 and the element has two stable isotopes 17Cl35 and 17Cl37. Estimate the relative abundances of these two isotopes. Solution: Here you finally get to put your high-school algebra to work! If we let x represent the fraction of Cl35, then (1-x) gives the fraction of Cl37.

The weighted average atomic weight is then 35x + 37(1-x) = 35.45 Solving for x gives 2x = 1.55, x = 0.775, so the abundances are 77.5% Cl35 and 22.5% Cl37. Problem Example 6 Elemental chlorine, Cl2, is made up of the two isotopes mentioned in the previous example. How many peaks would you expect to observe in the mass spectrum of Cl2? Solution: The mass spectrometer will detect a peak for each possible combination of the two isotopes in dichlorine: 35Cl-35Cl, 35Cl-37Cl, and 37Cl-37Cl. Calculate percent isotopic abundance from average atomic weight (MrB, 6½ min) Isotope effects The chemical behavior of an element is governed by the number and arrangement of its electrons in relation to its nuclear charge (atomic number). Because these quantities are identical for all isotopes of a given element, they are generally considered to exhibit identical chemical properties.

However, it turns out that the mass differences between different isotopes can give rise to very slight differences in their physical behavior that can, in turn affect their chemical behavior as well. These isotope effects are most evident in the lighter elements, in which small differences in neutron number lead to greater differences in atomic mass. Wikipedia article on heavy water Thus no element is more subject to isotope effects than hydrogen: an atom of "heavy hydrogen" 1H2 (also known as deuterium and often given the symbol D) has twice the mass of an atom of 1H2. When this isotope is combined with oxygen, the resulting "heavy water" D2O exhibits noticeably different physical and chemical properties: it melts at 3.8° C and boils at 101.4° C. D2O apparently interferes with cell division in organisms; mammals given only heavy water typically die in about a week. When two or more elements whose atoms contain multiple isotopes are present in a molecule, numerous isotopic modifications become possible.

For example, the two stable isotopes of hydrogen and of oxygen (O16 and O18) give rise to combinations such as H2O18, HDO16, etc., all of which are readily identifiable in the infrared spectra of water vapor. The amount of the rare isotopes of oxygen and hydrogen in water varies enough from place to place that it is now possible to determine the age and source of a particular water sample with some precision. These differences are reflected in the H and O isotopic profiles of organisms. Thus the isotopic analysis of human hair can be a useful tool for crime investigations and anthropology research. See also this Microbe Forensics page , and this general resource on water isotopes. Isotopic fractionation Isotope effects manifest themselves in both physical and chemical changes. In general, Chemical bonds involving lighter isotopes tend to be more readily broken, so reactions that depend on the rupture of such bonds will lead to products that are slightly enriched in the lighter isotopes. Phase changes such as vaporization similarly favor the lighter isotopes. Thus when water vapor condenses in clouds to form rain, the heavier water isotopes become slightly more concentrted in the liquid phase.

These two effects give rise to isotopic fractionation as chemical substances move through the environment — or on a much smaller scale, through the various metabolic processes that occur in organisms. Over time, this leads to changes in the isotopic signatures of elements in different realms of the world that can reveal information that would otherwise be hidden. The degree of isotopic fractionation depends on the temperature. This fact has been put to practical use to estimate the global average temperature in past times by measuring the degree of enrichment of the heavier isotopes of oxygen in glacial ice cores and also in ancient sediments containing the shells of microorganisms. Two interesting discussions of isotopic fractionation from Wesleyan U: Atomic weights, molecular weights and formula weights Molecules are composed of atoms, so a molecular weight is just the sum of the atomic weights of the elements it contains. Problem Example 7 What is the molecular weight of sulfuric acid, H2SO4? Solution: The atomic weights of hydrogen and of oxygen are 1.01 and 16.00, respectively (you should have these common values memorized.) From a table, you can find that the atomic weight of sulfur is 32.06.

Adding everything up, we have (2 x 1.01) + 32.06 + (4 x 16.00) = 98.08 Because some solids are not made up of discrete molecules (sodium chloride, NaCl, and silica, SiO2 are common examples), the term formula weight is often used in place of molecular weight. In general, the terms molecular weight and formula weight are interchangeable. 7 Atomic masses and the atomic mass unit Here again is the beginning of the atomic weight scale that you saw above: You understand by now that atomic weights are relative weights, based on a scale defined by 6C12 = 12. But what is the absolute weight of an atom, expressed in grams or kilograms? In other words, what actual mass does each unit on the atomic weight scale represent? The answer is 1.66053886 × 10–27 kg. This quantity (whose value you do not need to memorize) is known as the unified atomic mass unit , denoted by the abbreviation u. (Some older texts leave off the "unified" part, and call it the amu.

Wikipedia article on the unified atomic mass unit The unified atomic mass unit is defined as 1/12 of the mass of one atom of carbon-12. Why such a hard-to-remember number? Well, that's just how Nature sometimes does things. Fortunately, you don't need to memorize this value, because you can easily calculate its value from Avogadro's number, NA, which you are expected to know: 1 u = 1/NA gram = 1 ÷ (1000 NA) Kg ... but more about that in the later lesson on moles . Masses of the subatomic particles Atoms are composed of protons, neutrons, and electrons, whose properties are shown below: particle Two very important points you should note from this table: The mass of the electron is negligible compared to that of the two nuclear particles; The proton and neutron have masses that are almost, but not exactly, identical. Nuclear masses As we mentioned in one of the problem examples above, the mass of a nucleus is always slightly different from the masses of the nucleons (protons and neutrons) of which it is composed.

The difference, known as the mass defect, is related to the energy associated with the formation of the nucleus through Einstein's famous formula e = mc2. This is the one instance in chemistry in which conservation of mass-energy, rather than of mass alone, must be taken into account. But there is no need for you to be concerned with this in this part of the course. For all practical purposes, until you come to the section of the course on nuclear chemistry, you can consider that the proton and neutron have masses of about 1 u, and that the mass of an atom (in u) is just the sum of the neutron and proton numbers. What you should be able to do Make sure you thoroughly understand the following essential ideas which have been presented above. It is especially important that you know the precise meanings of all the highlighted terms in the context of this topic. Give a chemical definition of element © 2004-2016 by Stephen Lower - last modified 2016-05-13 For information about this Web site or to contact the author, please see the Chem1 Virtual Textbook home page.

The Chem1 Virtual Textbook home page is at This work is licensed under a Creative Commons Attribution-Share Alike 3.0 License . Chem1 Atoms, Elements and the Nucleus covers Basic atomic theory for a course in General Chemistry. It is part of the General Chemistry Virtual Textbook, a free, online reference textbook for General Chemistry by Stephen Lower of Simon Fraser University . This chapter covers the following topics: atoms & elements, laws of chemical change, Dalton's atomic theory, combining and relative weights of atoms, nuclei & nuclides, atomic weight scale,atomic masses. It can be accessed directly at . This material is directed mainly at the first-year college level, but much of it is also suitable for high-school students. It is licensed under a Creative Commons Attribution 3.0 Unported License. Chem1 Virtual Textbook home & main menu: ; Download the Chem1 Virtual Textbook from or from the Archive at Atomic Number and Mass Number ATOMIC NUMBER AND MASS NUMBERS After reading this section you will be able to do the following: Define and determine the atomic number of an atom. Define and determine the mass number of an atom. What is an atom's atomic number?

The number of protons in the nucleus of an atom determines an element's atomic number. In other words, each element has a unique number that identifies how many protons are in one atom of that element. For example, all hydrogen atoms, and only hydrogen atoms, contain one proton and have an atomic number of 1. All carbon atoms, and only carbon atoms, contain six protons and have an atomic number of 6. Oxygen atoms contain 8 protons and have an atomic number of 8. The atomic number of an element never changes, meaning that the number of protons in the nucleus of every atom in an element is always the same. What is an atom's mass number? All atoms have a mass number which is derived as follows. Review: An element's or isotope's atomic number tells how many protons are in its atoms. An element's or isotope's mass number tells how many protons and neutrons in its atoms. Periodic Table of Elements: Los Alamos National Laboratory -2 History From the Greek word oxys, acid, and genes, forming.

The behavior of oxygen and nitrogen as components of air led to the advancement of the phlogiston theory of combustion, which captured the minds of chemists for a century. Joseph Priestley is generally credited with its discovery, although Scheele also discovered it independently. Its atomic weight was used as a standard of comparison for each of the other elements until 1961 when the International Union of Pure and Applied Chemistry adopted carbon 12 as the new basis. Sources Oxygen is the third most abundant element found in the sun, and it plays a part in the carbon-nitrogen cycle, the process once thought to give the sun and stars their energy. Oxygen under excited conditions is responsible for the bright red and yellow-green colors of the Aurora Borealis. A gaseous element, oxygen forms 21% of the atmosphere by volume and is obtained by liquefaction and fractional distillation. The atmosphere of Mars contains about 0.15% oxygen. The element and its compounds make up 49.2%, by weight, of the earth's crust. About two thirds of the human body and nine tenths of water is oxygen.

In the laboratory it can be prepared by the electrolysis of water or by heating potassium chlorate with manganese dioxide as a catalyst. Properties The gas is colorless, odorless, and tasteless. The liquid and solid forms are a pale blue color and are strongly paramagnetic. Forms Ozone (O3), a highly active compound, is formed by the action of an electrical discharge or ultraviolet light on oxygen. Ozone's presence in the atmosphere (amounting to the equivalent of a layer 3 mm thick under ordinary pressures and temperatures) helps prevent harmful ultraviolet rays of the sun from reaching the earth's surface. Pollutants in the atmosphere may have a detrimental effect on this ozone layer. Ozone is toxic and exposure should not exceed 0.2 mg/m# (8-hour time-weighted average - 40-hour work week). Undiluted ozone has a bluish color. Liquid ozone is bluish black and solid ozone is violet-black. Compounds Oxygen, which is very reactive, is a component of hundreds of thousands of organic compounds and combines with most elements. Uses Plants and animals rely on oxygen for respiration. Hospitals frequently prescribe oxygen for patients with respiratory ailments. Isotopes Oxygen has nine isotopes.

Natural oxygen is a mixture of three isotopes. Natural occurring oxygen-18 is stable and available commercially, as is water (H2O with 15% 18O). Commercial oxygen consumption in the U.S. is estimated at 20 million short tons per year and the demand is expected to increase substantially. Oxygen enrichment of steel blast furnaces accounts for the greatest use of the gas. Large quantities are also used in making synthesis gas for ammonia and methanol, ethylene oxide, and for oxy-acetylene welding. Air separation plants produce about 99% of the gas, while electrolysis plants produce about 1%.

Question: Which common, gaseous element has the atomic Number 8, and an Atomic Weight of 16?
Answer: OXYGEN