Atomic Number 3 Element Facts Atomic Number 3 Element Facts Atomic Number 3 Element Facts What Element is Atomic Number 3? Each atom of atomic number 3 has three protons. Lithium may have a different number of neutrons and electrons, depending on the isotope or ion. ROGER HARRIS/SCIENCE PHOTO LIBRARY / Getty Images By Anne Marie Helmenstine, Ph.D. Lithium is the element that is atomic number 3 on the periodic table. These means each atom contains 3 protons. Lithium is a soft, silvery, light alkali metal denoted with the symbol Li. Here are interesting facts about atomic number 3: Lithium is the lightest metal and the lightest solid element at ordinary temperature and pressure. It also has the highest specific heat capacity of a solid element. Element number 3 is soft enough to cut with shears and light enough to float on water. Among its uses, lithium is used in medications for bipolar disorder, to make lithium ion batteries, and to add a red color to fireworks . It's also used in glass and ceramics and to make high temperature lubricant grease.

It is a coolant in breeder reactors and a source of tritium when atomic number 3 is bombarded with neutrons. Lithium is the only alkali metal that reacts with nitrogen. Yet, it is the least reactive metal in its element group. While it burns in water, it does not do so as vigorously as sodium or potassium. Lithium metal will burn in air and should be stored under kerosene or in an inert atmosphere, like argon. Don't try to extinguish a lithium fire with water -- it will only make it worse! continue reading below our video 10 Facts About the Titanic That You Don't Know Because the human body contains a lot of water, lithium will also burn skin. It is corrosive and should not be handled without protective gear. The name for the element comes from the Greek word "lithos", which means "stone". Lithium was discovered in the mineral petalite (LiAISi4O10). Brazilian naturalist and statesman, Jozé Bonifácio de Andralda e Silva found the stone on the Swedish isle Utö. Although the mineral looked like an ordinary gray rock, it flared red when thrown into a fire.

Swedish chemist Johan August Arfvedson determined the mineral contained a previously unknown element. He couldn't isolate a pure specimen, but did produce a lithium salt from petalite in 1817. The atomic mass of lithium is 6.941. It's density is 0.52 grams per cubic centimeter, which means it not only floats on water, but is just over half as dense! Lithium is believed to be one of only three chemical elements produced in the Big Bang that formed the universe. The other two elements are hydrogen and helium . Several isotopes of lithium are known, but the natural element is a mix of two stable isotopes. Li-7 (92.41 percent natural abundance) and Li-6 (7.59 percent natural abundance). Lithium readily loses its outer electron to form the Li+ ion. This leaves the atom with a stable inner shell of two electrons. The lithium ion readily conducts electricity. Because of its high reactivity, lithium is not found in nature as a pure element, but the ion is abundant in sea water. Lithium compounds are found in clay.

Mankind's first fusion reaction involved atomic number 3, in which lithium was used to make hydrogen isotopes for fusion by Mark Oliphant in 1932. Periodic Table of Elements: Los Alamos National Laboratory 7, 6, 5, 4, 3, 0 History Plutonium is the second transuranium element of the actinide series. Element 93 was discovered in 1940/41 by Glenn T. Seaborg , Edwin M. McMillan , J. W. Kennedy , and A. C. Wahl by deuteron bombardment of uranium-238 in the 60-inch cyclotron at the University of California, Berkeley Lab. They first synthesized neptunium-238 (half-life 2.1 days) which subsequently beta-decayed to form a new heavier element with atomic number 94 and atomic weight 238 (half-life 87.7 years). It was fitting that element 94 be named after the next planetoid, Pluto following the precedence that uranium was named after the planet Uranus and neptunium after the planet Neptune.

Seaborg submitted a paper to the journal Physical Review in March 1941 documenting the discovery, but the paper was quickly withdrawn when it was found that an isotope of plutonium, Pu-239 could undergo nuclear fission making it useful in developing an atomic bomb. Pu-239 had a fission cross-section 50% greater than that of 235U, the best fissioning element known at that time. Seaborg was called away from Berkeley to lead the Plutonium Production Lab or "Met Lab" at the University of Chicago. The Met Lab was to produce useful quantities of plutonium as part of the secret Manhattan Project during World War II to develop an atomic bomb. On August 18, 1942, a trace quantity of plutonium was isolated and measured at the Met Lab for the first time. About 50 micrograms of Pu-239 combined with uranium and fission products was produced and only about 1 microgram was isolated. This was enough material for chemists to determine the new element's atomic weight. In November 1943 a few milligrams of PuF3 was reduced to create the first sample of plutonium metal.

Enough plutonium was produced to make it the first man-made element to be visible to the unaided eye. The nuclear properties of plutonium-239 were also being studied and researchers found that when hit with a neutron it fissions by releasing energy and more neutrons. These neutrons can hit neighboring atoms of Pu-239 and so on, in an exponentially fast chain-reaction, releasing a tremendous amount of energy. This energy could result in an explosion large enough to destroy a city or fuel a nuclear reactor. During WW II the three primary research and production sites of the Manhattan Project were the Plutonium Production Facility at what is now the Hanford Site , Washington, the Uranium Enrichment facilities at Oak Ridge, Tennessee , and the weapons research and design laboratory, now known as Los Alamos National Laboratory . In 1943, the first production reactor that made Pu-239 was the X-10 Graphite Reactor built at a facility in Oak Ridge, Tennessee that later became the Oak Ridge National Laboratory. The Manhattan Project produced the plutonium for the "Trinity Test" conducted in New Mexico by Los Alamos Laboratory Director Robert Oppenheimer and Army General Leslie Groves.

The world’s first atomic bomb ("The Gadget") was exploded near Socorro, New Mexico on July 16, 1945, resulting in an explosion with an energy equivalent of approximately 20,000 tons of TNT . The first atomic bomb used in war had a uranium core and was dropped on Hiroshima, Japan on August 6, 1945. The second atomic bomb used had a plutonium core and was nicknamed " Fat Man " because of its round shape. It was used to destroy Nagasaki, Japan in August 9, 1945, which put an end to WW II. Publication of the discovery and the naming of the new element plutonium was delayed until a year after the end of World War II . Seaborg originally considered the name "plutium", but later thought that it did not sound as good as "plutonium." Later, during the Cold-War era, large stockpiles of weapons-grade plutonium were built up by both the Soviet Union and the United States. Each year about 20 tons of plutonium is still produced as a by-product of the nuclear power industry.

As of 2007 it was estimated that the plutonium stockpile was about 500 tons, world-wide. Since the end of the Cold War these stockpiles have become a focus of nuclear proliferation concerns. In 2000, the United States and the Russian Federation mutually agreed to each dispose of 34 tons of weapon grade plutonium before the end of 2019 by converting it to a mixed uranium-plutonium oxide (MOX) fuel to be used in commercial nuclear power reactors. Today plutonium-239 remains an important component of nuclear weapons, and the United States maintains plutonium-related capabilities in support of national defense and global nuclear deterrence. Pu-239 for civilian nuclear power plants provides energy for many nations. Plutonium-238 continues to be vital to space exploration pushing the limits beyond which manned space exploration is possible and satisfying our quest for knowledge. Isotopes Twenty-three radioactive isotopes of plutonium have been characterized from mass numbers 228 to 247. Nine of these exhibit metastable states , though these all have half-lives less than one second.

The longest-lived isotopes are plutonium-244, with a half-life of 80.8 million years, plutonium-242, with a half-life of 373,300 years, and plutonium-239, with a half-life of 24,110 years. All of the remaining radioactive isotopes have half-lives less than 7,000 years. The primary decay modes of isotopes with mass numbers lower than plutonium-244, are spontaneous fission and α emission , mostly forming uranium and neptunium isotopes as decay products along with a variety of daughter fission products. The primary decay mode for isotopes with mass numbers higher than plutonium-244 is by β emission , mostly forming americium isotopes as daughter decay products. Plutonium-241 is the parent isotope of the neptunium decay series , decaying to americium-241 via β decay.

By far of greatest importance is the isotope 239Pu produced in extensive quantities in nuclear reactors from natural uranium: 238U (n, gamma) → 239U (beta) → 239Np (beta) → 239Pu Plutonium-238 with a half-life of 87.7 years is another important isotope. Both Pu-239 and Pu-238 have many practical applications as discussed below. Properties Plutonium is unique among the elements in its physicochemical complexities by virtue of its position at a transitional location in the periodic table where the 5f electrons are at the border between delocalized (not associated with a single atom) and localized (associated with a single atom) behavior and it is considered one of the most complex of the elements. Plutonium also sits near the juncture where the actinide series transitions from main d-block element chemistry to rare earth like behavior as a result of the actinide contraction. Because of its defense and commercial importance, plutonium is one of the most intensely investigated of elements.

Plutonium metal has a bright silvery appearance at first and takes on a dull gray, yellow or olive green tarnish when oxidized in air. A relatively large piece of plutonium is warm to the touch because of the energy given off by alpha decay. Larger pieces will produce enough heat to boil water. The metal readily dissolves in concentrated mineral acids. Plutonium metal normally has six allotropes or crystal structures; alpha (α), beta (β), gamma (γ), delta (δ), delta prime (δ') and epsilon (ε). It forms a seventh phase (zeta, ζ) under high temperature and a limited pressure range. These allotropes have very similar energy levels but significantly varying densities (from 16.00 to 19.86 grams/cm3) and crystal structures . This makes plutonium very sensitive to changes in temperature, pressure, or chemistry, and allows for dramatic volume changes following phase transitions . At room temperature plutonium is in its alpha (α) form , the most common structural form of the element. It is as hard and brittle as cast iron unless alloyed with other metals to form the room-temperature stabilized delta (δ) phase which makes it soft and ductile.

Unlike most metals, it is not a good conductor of heat or electricity . It has a low melting point (640 °C) and an unusually high boiling point (3,228 °C). A ring of 99.96% pure electro-refined plutonium (Los Alamos National Laboratory) Plutonium can form alloys and intermediate compounds with most other metals. Gallium, aluminum, americium, scandium and cerium can stabilize the δ phase of plutonium metal. Nuclear fuel pellets can be formed by alloying plutonium with various metals such as: aluminum; zirconium; cerium; cerium-cobalt; uranium-titanium, uranium-zirconium and uranium-molybdenum. Thorium-plutonium-uranium alloys were investigated as a nuclear fuel for fast breeder reactors. A plutonium-gallium-cobalt alloy (PuCoGa5) was found to be an unconventional superconductor, showing superconductivity below 18.5 Kelvin, an order of magnitude higher than the highest between heavy fermion systems known. Plutonium forms compounds with a variety of other elements. Plutonium reacts with pure hydrogen, forming plutonium hydride s.

It also reacts readily with oxygen, forming PuO and PuO2 as well as intermediate and sub-stoichiometric oxides. The metal reacts with the halogens , giving rise to trivalent Pu compounds with the general formula PuX3 where X can be F , Cl , Br or I and tetravalent plutonium compounds such as PuF4 . The following oxyhalides are observed: PuOCl, PuOBr and PuOI. Plutonium reacts with carbon to form PuC, nitrogen to form PuN and silicon to form PuSi2. Pu3+ and Pu4+ oxalates are important intermediates that are calcined to form oxides as a step in plutonium processing. Other important compounds in reprocessing are fluoride, peroxide, acetylacetone, carbonate and hydroxide. The color displayed by plutonium solutions depends on both the oxidation state and the extent of complexation by various ligands.

In aqueous solution plutonium exhibits five ionic valence states: Pu+3 (blue lavender), Pu+4 (salmon-colored, when uncomplexed), PuO+ (lavender), PuO+2 (orange-brown) and PuOxOHy (dark green in basic solution). The pentavalent ion, PuO+ is unstable in aqueous solutions and it disproportionates into Pu+4 and PuO+2. However, PuO2+ can be stabilized in aqueous solution in a narrow pH range around 4.5. By virtue of the close proximity of the electrode potentials of the various plutonium redox couples (~ 1 Volt/NHE), four oxidation states can co-exist in solution simultaneously: Pu3+, Pu4+, PuO2+ and PuO22+. Various oxidation states of plutonium in aqueous solution; all in non-complexing perchloric acid, except Pu(VII) in strong base. (D. E. Hobart and P.D. Palmer, Los Alamos National Lab.) Pu4+ is a "hard" (ionic) cation with the largest electronic charge of plutonium ions and it forms complexes with a variety of inorganic and organic ligands.

In dilute perchloric acid, Pu4+ is un-complexed and is salmon-colored. However in concentrated acids, Pu4+ forms anionic complexes such as: Pu(NO3)62- (dark green) and Pu(Cl)62- (brick red). Pu4+, having a high ionic charge readily hydrolyzes (combines with hydroxide ion) at near-neutral pH values forming a green colloidal suspension that behaves like a solution but is actually a solid precipitate that can be separated by ultra-centrifugation. Plutonium-organic complexes are very important for separation, reprocessing, and purification and include: Tributyl phosphate (TBP); Di-(2-ethylhexyl)phosphoric acid (DEHPA or HDEHP); octyl(phenyl)-N,N-diisobutyl-carbamoylmethylphosphine oxide (CMPO); crown-ethers; and many others. Production Plutonium is the heaviest primordial element in existence by virtue of its most stable isotope, plutonium-244, whose half-life of about 80 million years is just long enough for the element to be found in trace quantities in nature.

In 1971 mass spectrometric measurements of plutonium isolated from Precambrian bastnasite (a lanthanide chloride-fluoride mineral) by Hoffman and coworkers confirmed the presence of Pu-244 in nature. Although the existence of Pu-244 as an extinct radioactive element had been postulated to explain the xenon isotope ratios observed in meteorites, this is the first indication of its present existence in nature. Plutonium also exists in trace quantities in naturally occurring uranium ores. It is formed in much the same manner as neptunium: by irradiation of natural uranium with the neutrons followed by beta decay. Exceedingly small amounts of Pu-238, attributed to the extremely rare double-beta decay of U-238, have been found in natural uranium samples. Plutonium was most likely formed by neutron activation of natural U-238 at the Oklo natural reactor but if formed it has long since decayed away. Plutonium is for the most part a by-product of nuclear reactions in reactors where some of the neutrons released by the fission process convert U-238 nuclei into plutonium.

Plutonium-238 and -239 are the most widely synthesized isotopes. Plutonium-239 is synthesized via the following reaction using uranium (U) and neutrons (n) via beta decay (β−) with neptunium as an intermediate: Plutonium-238 is synthesized by bombarding uranium-238 with deuterons in the following first discovery reaction: Uses Plutonium has assumed the position of dominant importance among the transuranium elements because of its use as an explosive ingredient in nuclear weapons and the place which it holds as a key material in the development of industrial use of nuclear power. During fission, a fraction of the binding energy, which holds a nucleus together, is released as a large amount of electromagnetic and kinetic energy which is quickly converted to thermal energy. Fission of a kilogram of plutonium-239 can produce an explosion equivalent to 21,000 tons of TNT which is equivalent to about 22 million kilowatt hours of heat energy. In 1982 it was estimated that about 300,000 kg had accumulated.

The most common chemical process, PUREX (Plutonium–URanium EXtraction) reprocesses spent nuclear fuel to extract plutonium and uranium which can be used to form a mixed U/Pu oxide or "MOX" fuel for reuse in nuclear power reactors. MOX fuel production is also a good mechanism to reduce excessive defense plutonium stockpiles for peaceful purposes, which in effect is forging "swords into plowshares." Plutonium isotopes undergo radioactive decay, which produces decay heat . Different isotopes produce different amounts of heat per mass. Pu-238 with a half-life of 88 years has a relatively high heat production rate which makes it useful as a power source with a long service life. The decay heat is usually listed as watt/kilogram, or milliwatt/gram. Pu-238 is a heat source in radioisotope thermoelectric generators, which are used to power spacecraft and extra-terrestrial rovers. As a power and heat source, Pu-238 has also been used to power instruments left on the Moon by Apollo astronauts, weather satellites and interplanetary probes and powers the Cassini Saturn mission and the Mars rovers.

Plutonium-238 was at one time used successfully to power artificial heart pacemakers but has been replaced by lithium-based primary cells. Plutonium-238 was studied as a way to provide supplemental heat to scuba divers. Pu-238 mixed with beryllium is a convenient method to generate neutrons. Pu-238 powers the Cassini Mission now orbiting Saturn. (NASA) Hazards Plutonium, along with all of the other transuranium elements, is a radiological hazard and must be handled with specialized equipment and precautions. Animal studies have found that a few milligrams of plutonium per kilogram of tissue are lethal. Plutonium is more dangerous when inhaled than when ingested. When inhaled, plutonium can pass into the bloodstream moving throughout the body and into the bones, liver, or other body organs. Plutonium that reaches body organs generally stays in the body for decades and continues to expose the surrounding tissue to radiation and thus may cause cancer. Precautions must also be taken to prevent the unintentional formulation of a critical mass.

Pure Pu-239 metal present in sufficient quantity with an appropriate geometry can form a critical mass and pose a serious and often lethal criticality event. Low levels of plutonium in the environment is a result of world-wide fallout from numerous above-ground nuclear tests during the Cold War (now banned) pose little health hazard. Disposal of plutonium waste from nuclear power plants and dismantled nuclear weapons built during the Cold War is a nuclear-proliferation and environmental concern. Further Reading D. L. Clark, S. S. Hecker, G. D. Jarvinen, and M. P. Neu (2006) "Plutonium", Chapter 7 in The Chemistry of the Actinide and Transactinide Elements, Third Edition, L. R. Morss, J. Fuger, and N. M. Edelstein, Eds, Springer Publishers. Los Alamos Science (2000) Number 26, Volume I and II, Los Alamos National Laboratory report LA-UR-00-4100. Updated by Dr. David Hobart, August 2013.

LA-UR-13-26653 Lithium»the essentials [WebElements Periodic Table] Element News Lithium: the essentials Lithium is a Group 1 (IA) element containing just a single valence electron (1s22s1). Group 1 elements are called "alkali metals". Lithium is a solid only about half as dense as water and lithium metal is the least dense metal. A freshly cut chunk of lithium is silvery, but tarnishes in a minute or so in air to give a grey surface. Its chemistry is dominated by its tendency to lose an electron to form Li+. It is the first element within the second period. Lithium is mixed (alloyed) with aluminium and magnesium for light-weight alloys, and is also used in batteries, some greases, some glasses, and in medicine. Lithium does not occur as the free metal in nature because of its high reactivity. Deposits are known all aroun the world. It is a minor component of nearly all igneous rocks and is a component of many natural brines.

CAS Registry ID : 7439-93-2 Lithium: historical information Lithium was discovered by Johan August Arfvedson in 1817 at Stockholm, Sweden. Origin of name : from the Greek word "lithos" meaning "stone", apparently because it was discovered from a mineral source whereas the other two common Group 1 elements, sodium and potassium, were discovered from plant sources.. The mineral petalite (which contains lithium) was discovered by the Brazilian scientist José Bonifácio de Andrada e Silva towards the end of the 18th century while visiting Sweden. Lithium was discovered by Johan August Arfvedson in 1817 during an analysis of petalite ore, an ore now recognised to be LiAl(Si2O5)2, taken from the Swedish island of Utö. Arfvedson subsequently discovered lithium in the minerals spodumene and lepidolite. C.G. Gmelin observed in 1818 that lithium salts colour flames bright red. Neither Gmelin nor Arfvedson were able to isolate the element itself from lithium salts, for example in attempted reductions by heating the oxide with iron or carbon.

The first isolation of elemental lithium was achieved later by W.T. Brande and Sir Humphrey Davy by the electrolysis of lithium oxide. In 1855, Bunsen and Mattiessen isolated larger quantities of the metal by electrolysis of lithium chloride. In 1923 the first commercial production of lithium metal was achieved by Metallgesellschaft AG in Germany using the electrolysis of a molten mixture of lithium chloride and potassium chloride, exploiting a suggestion made by Guntz in 1893. Lithium around us Read more » Lithium seems to have no biological role, but does have an effect on the body if swallowed. Sometimes, lithium-based compounds such as lithium carbonate (Li2CO3) are used as drugs to treat manic-depressive disorders. The dose is around 0.5 g - 2 g daily. Lithium does not occur as the free metal in nature because of its high reactivity. Deposits are known all aroun the world. It is a minor component of nearly all igneous rocks and is a component of many natural brines (see below).

Large deposits are located in California and Nevada (both in the USA) in several rock forms, particularly spodumene. The four main lithium minerals are spodumene, lepidolite, petalite, and amblygonite. spodumene: LiAlSi2O6. This is the most important and abundant of the lithium ores. Deposits are located in North America, Brazil, USSR, Spain, parts of Africa, and Argentina. One method of extraction involves converting α-spodumene (the naturally occurring form) to the β-form (a less dense material) by heating to 1100°C, mixing with sulphuric acid, H2SO4, and heating to 250°C. This is followed by extracting into water to give a lithium sulphate solution, Li2SO4, suitable for further processing. lepidolite: K2Li3Al4Si7O21(OH,F)3. Deposits are located in Canada and parts of Africa. The mineral sometimes contains caesium and rubidium. Lithium can be extracted by similar methods to those of spodumene petalite: LiAlSi4O10.

Deposits are located in parts of Africa and Sweden. amblygonite: LiAl(F,OH)PO)4. Amblygonite occurs in only minor deposits Lithium is also recovered from lakes such as Searles Lake (California, USA) and Clayton Valley (Nevada, USA). Lithium is extracted from the brine by solar evaporation, precipitation of Group 2 elements if necessary, and precipiation of lithium carbonate by addition of sodium carbonate to the hot brine. Abundances for lithium in a number of different environments. More abundance data » Location Second ionisation energy : 7298.1 kJ mol‑1 Isolation Isolation : lithium would not normally be made in the laboratory as it is so readily available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative lithium ion Li+. The ore spodumene, LiAl(SiO3)2, is the most important commercial ore containing lithium. The α form is first converted into the softer β form by heating to around 1100°C.

This is mixed carefully with hot sulphuric acid and extracted into water to form lithium sulphate, Li2SO4, solution. The sulphate is washed with sodium carbonate, Na2CO3, to form a precipitate of the relatively insoluble lithium carbonate, Li2CO3. Li2SO4 + Na2CO3 → Na2SO4 + Li2CO3 (solid) Reaction of lithium carbonate with HCl then provides lithium chloride, LiCl. Li2CO3 + 2HCl → 2LiCl + CO2 +H2O Lithium chloride has a high melting point (> 600°C) meaning that it sould be expensive to melt it in order to carry out the electrolysis. However a mixture of LiCl (55%) and KCl (45%) melts at about 430°C and so much less energy and so expense is required for the electrolysis.

cathode: Li+(l) + e- → Li (l) anode: Cl-(l) → 1/2Cl2 (g) + e- Lithium isotopes Read more » Li-7 is used to control the pH level of the coolant in the primary water circuit of pressurized water reactors. Li-7 is also used for the production of the medical research radioisotope Be-7. Li-6 is used in thermonuclear weapons and the export and use of Li-6 is therefore strictly controlled. Li-6 can also be used for the production of the radioisotope H-3, which is used in biochemistry research. Mendeleev's First Periodic Table Elements and Atoms: Chapter 12 Mendeleev's First Periodic Table Dmitrii Mendeleev (1834-1907; see photo in Chemical Achievers at the Chemical Heritage Foundation) was born in Tobolsk, in Western Siberia. His chief contribution to chemistry was the establishment of the periodic system of elements.

Mendeleev was one of a number of independent discoverers of the periodic law in the 1860s--that number ranging from one [Leicester 1948] to six [van Spronsen 1969] depending on the criteria one adopts. Mendeleev's formulation was clearly superior in several respects to the work of contemporary classifiers: it was the clearest, most consistent, and most systematic formulation, and Mendeleev made several testable predictions based on it. It was not, however, free from error. Scientists, even great scientists, trying to see further than others have in the past, do not always see the whole picture clearly. As noted below, Mendeleev himself corrected some of the errors within a few years; others persisted well into the 20th century. This table and the accompanying observations were first presented to the Russian Chemical Society in March 1869. (Actually, Mendeleev was ill, and his colleague Nikolai Menshutkin presented his paper [Menschutkin 1934].) The paper was published in the first volume of the new society's journal.

That same year, a German abstract of the paper, consisting of the table and eight comments, was published in Zeitschrift für Chemie. The German abstract was the vehicle by which Mendeleev's ideas reached chemists working in Western Europe. An English translation of that German abstract is presented here. View a manuscript draft of the table . On the Relationship of the Properties of the Elements to their Atomic Weights D. Mendelejeff, Zeitschrift für Chemie 12, 405-406 (1869); translation by Carmen Giunta By ordering the elements according to increasing atomic weight in vertical rows so that the horizontal rows contain analogous elements, [1] still ordered by increasing atomic weight, one obtains the following arrangement, from which a few general conclusions may be derived. Ti=50 ?In=75,6 Th=118? The elements, if arranged according to their atomic weights, exhibit an evident stepwise variation [12] of properties. Chemically analogous elements have either similar atomic weights [13] (Pt, Ir, Os), or weights which increase by equal increments (K, Rb, Cs).

[14] The arrangement according to atomic weight corresponds to the valence [15] of the element and to a certain extent the difference in chemical behavior, for example Li, Be, B, C, N, O, F. The elements distributed most widely in nature have small atomic weights [16] , and all such elements are marked by the distinctness of their behavior. They are, therefore, the representative elements; and so the lightest element H is rightly chosen as the most representative. The magnitude of the atomic weight determines the properties of the element. Therefore, in the study of compounds, not only the quantities and properties of the elements and their reciprocal behavior is to be taken into consideration, but also the atomic weight of the elements. Thus the compounds of S and Tl [sic--Te was intended], Cl and J, display not only analogies, but also striking differences. [17] One can predict the discovery of many new elements, for example analogues of Si and Al with atomic weights of 65-75. [18] A few atomic weights will probably require correction; for example Te cannot have the atomic weight 128, but rather 123-126.

[19] From the above table, some new analogies between elements are revealed. Thus Bo (?) [sic--apparently Ur was intended] appears as an analogue of Bo and Al, as is well known to have been long established experimentally. (Russian Chemical Society 1, 60) Notes [1] The principle of periodicity is apparent in this first sentence: repetition of chemical properties in a series of elements arranged by atomic weight. Note also the cosmetic difference between Mendeleev's layout and that of modern tables, which order the elements in horizontal rows such that families of elements appear in vertical columns. Mendeleev would use the latter sort of layout before long [ Mendeleev 1871 ]. [2] This table contains several implicit predictions of unknown elements. Mendeleev soon retreated from this prediction of a heavier analogue of titanium and zirconium. His later tables [ Mendeleev 1871 ] erroneously placed lanthanum in this spot. This original prediction was actually borne out in 1923 with the discovery of hafnium. [3] Rhodium (Rh) is misplaced. It belongs between ruthenium (Ru) and palladium (Pd).

Technetium (Tc), the element which belongs between ruthenium and molybdenum (Mo) has no stable isotopes and was not synthesized until 1937. [4] Most of the elements in this column are slightly out of order. After tungsten (W) should come rhenium (Re), which was not yet discovered, followed by osmium (Os), iridium (Ir), platinum (Pt), gold (Au), mercury (Hg), thallium (Tl), lead (Pb), and bismuth (Bi). Bismuth, however, is placed correctly insofar as it completes the row beginning with nitrogen (N). At this time, lead was frequently miscategorized, placed among elements which form compounds with one atom of oxygen (PbO analogous to CaO, for example); however, lead also forms a compound with two atoms of oxygen (PbO2 analogous to CO2) and it belongs in the same group as carbon (C).

Similarly, thallium was often placed among elements which form compounds with one atom of chlorine (TlCl analogous to NaCl, for example); however, thallium also forms a compound with three atoms of chlorine (TlCl3 analogous to BCl3) and it belongs in the same group as boron (B). [5] The classification of hydrogen has been an issue throughout the history of periodic systems. Some tables place hydrogen with the alkali metals (lithium, sodium, etc.), some with the halogens (fluorine, chlorine, etc.), some with both, and some in a box of its own detached from the main body of the table. Mendeleev's original table did none of the above, placing it in the same row as copper, silver, and mercury. [6] The prediction of an unknown analogue of aluminum was borne out by the discovery in 1875 of gallium (atomic weight = 70), the first of Mendeleev's predictions to be so confirmed. [ Lecoq de Boisbaudran 1877 ] [7] Uranium (standard symbol U) is misplaced. Its atomic weight is actually more than double the value given here.

The element which belongs between cadmium (Cd) and tin (Sn) is indium (In), and Mendeleev put indium there in the next version of his table [ Mendeleev 1871 ]. The proper place for uranium, however, would not be found until the 1940s. [8] The prediction of an unknown analogue of silicon was borne out by the discovery in 1886 of germanium (atomic weight = 73). [ Winkler 1886 ] [9] In German publications, J is frequently used instead of I as the chemical symbol for iodine (Jod, in German). Iodine is placed correctly after tellurium (i.e., with the halogens) despite having a lower atomic weight than tellurium. See comment 7 after the table. [10] The prediction of an unknown element following calcium is a weak version of Mendeleev's subsequent prediction of the element we now know as scandium, discovered in 1879 [ Nilson 1879 ].

In Mendeleev's 1871 table [ Mendeleev 1871 ] the missing element is correctly placed between calcium and titanium, and as an analogue of yttrium. That the 1869 prediction is flawed can be seen from the fact that every other entry in the bottom reaches of the table is wrong. (See next note .) Still, the prediction deserves more credit than van Spronsen gave it [van Spronsen 1969, p. 220]: "That the element next to calcium later proved to be scandium, was fortuitous; Mendeleev cannot be said to have already foreseen this element in 1869." [11] The elements placed in the last four rows of the table puzzled Mendeleev, as is apparent from the glut of question marks and the fact that several are out of order according to their assigned atomic weights. Many of these elements were rare and poorly characterized at the time. Didymium appeared in many lists of elements at this time, but it was later proved to consist of two elements, praseodymium and neodymium.

The atomic weights of erbium, yttrium (standard symbol Y), indium, cerium, lanthanum, and thorium are wrong. The interdependence of atomic weights and chemical formulas that plagued determinations of atomic weight since the time of Dalton was still problematic for these elements. Most of them elements (erbium, yttrium, cerium, lanthanum, and the component elements of didymium) belong to the family of rare earths, a group whose classification would present problems for many years to come. (Thorium belongs to the group of elements immediately below most of the rare earths.) Many of the rare earths were not yet discovered, and (as already noted) the atomic weights of the known elements were not well determined. The chemical properties of the rare earths are so similar that they were difficult to distinguish and to separate. Mendeleev made some progress with these elements in the next couple of years. His 1871 table [ Mendeleev 1871 ] has correct weights for yttrium, indium, cerium, and thorium, and correct classification for yttrium and indium.

[12] Translator's note: In his 1889 Faraday lecture [ Mendeleev 1889 ], Mendeleev used the word "periodicity" rather than the phrase "stepwise variation" in translating this sentence from his 1869 paper. "Periodicity" is certainly an appropriate term to describe the cyclic repetition in properties evident in this arrangement. It is worth noting, however, that the German words read in 1869 by Western European scientists (stufenweise Abänderung) lack the implication of repetition inherent in the term periodicity. --CJG [13] Groups of similar elements with consecutive atomic weights are a little-emphasized part of classification systems from Mendeleev's time and before (cf. Newlands 1864 ) to the present. [14] The existence of a very regular progression in atomic weight among elements with similar chemical behavior had attracted the attention of chemists almost from the time they began to measure atomic weights [ Döbereiner 1829 ]. The triad of elements Mendeleev cites here includes two (rubidium and cesium) discovered in the early 1860s.

Mendeleev's table, however, goes beyond strictly regular isolated triads of elements to a systematic classification (albeit not always correct) of all known elements. [15] The valence of an element is essentially the number of bonds that element can make when it forms compounds with other elements. An atom of hydrogen, for example, can make just one bond, so its valence is one; we call it monovalent. An atom of oxygen can bond with two atoms of hydrogen, so its valence is two. Some elements, particularly heavier elements, have more than one characteristic valence. (For example, lead has valence 2 and 4; thallium has valence 1 and 3. See note 4 above.) The elements in the cited series have valences 1, 2, 3, 4, 3, 2, and 1 respectively. [16] Mendeleev is correct in this observation. The two lightest elements, hydrogen and helium (the latter as yet unknown) are the most common elements in the universe, making up the bulk of stars. Oxygen and silicon are the most common elements in the earth's crust.

Iron is the heaviest element among the most abundant elements in the stars and the earth's crust. [17] Although the chemical behavior of elements in the same family is similar, it is not identical: there are differences due to the difference in atomic weight. For example, both chlorine and iodine form compounds with one atom of hydrogen: HCl and HI. These compounds are similar, in that they are both corrosive gases which dissolve readily in water. But they differ in that HI has, for example, a higher boiling point and melting point than HCl (typical of the heavier of a pair of related compounds). [18] In later publications [ Mendeleev 1871 ] Mendeleev went into considerable detail regarding the properties of predicted elements. The success of these predictions played a part in establishing the periodic system, although apparently not the primary part. [Brush 1996] See Scerri & Worrall 2001 for a discussion of prediction and accommodation in the periodic table.

[19] Mendeleev went on to incorporate this "correction" in his 1871 table [ Mendeleev 1871 ], listing the atomic weight of tellurium as 125. But the "correction" is erroneous. Mendeleev was right to put tellurium in the same group with sulfur and oxygen; however, strict order of atomic weights according to the best information he had available would have required iodine (127) to come before tellurium (128). He was suspicious of this apparent inversion of atomic weight order; as it happens, the atomic weights Mendeleev had available to him agree with the currently accepted values. While his suggestion to change that of tellurium was wrong, his classification was correct and his faith in the regularity of the periodic system was only slightly misplaced. The natural order of the elements is not quite one of increasing atomic weight, but one of increasing atomic number. In 1913, a discovery by Henry Moseley made the atomic number more than simply a rank order for the elements [ Moseley 1913, 1914 ].

The atomic number is the same as the quantity of positive charge in the nucleus of an atom. The periodic system contains a few "inversions" of atomic weight, but no inversions of atomic number. References

Question: With the Atomic Number 3, which is the first metal in the Periodic Table of elements?
Answer: LITHIUM