Chemistry ELEMENTS PERIODIC TABLE

1: H: Hydrogen: 1.008 History From the Greek word hydro (water), and genes (forming). Hydrogen was recognized as a distinct substance by Henry Cavendish in 1776.

Electron Configuration: 1s1 Atomic Radius: 110 pm (Van der Waals) Melting Point: -259.34 °C Boiling Point: -252.87 °C Oxidation States: 1, -1 Hydrogen is the most abundant of all elements in the universe. The heavier elements were originally made from hydrogen atoms or from other elements that were originally made from hydrogen atoms. Hydrogen is estimated to make up more than 90% of all the atoms -- three quarters of the mass of the universe! This element is found in the stars, and plays an important part in powering the universe through both the proton-proton reaction and carbon-nitrogen cycle. Stellar hydrogen fusion processes release massive amounts of energy by combining hydrogens to form helium. Isotopes The ordinary isotope of hydrogen, H, is known as Protium, the other two isotopes are Deuterium (a proton and a neutron) and Tritium (a protron and two neutrons). Hydrogen is the only element whose isotopes have been given different names. Deuterium and Tritium are both used as fuel in nuclear fusion reactors. One atom of Deuterium is found in about 6000 ordinary hydrogen atoms. Deuterium is used as a moderator to slow down neutrons. Tritium atoms are also present but in much smaller proportions. Tritium is readily produced in nuclear reactors and is used in the production of the hydrogen (fusion) bomb. It is also used as a radioactive agent in making luminous paints, and as a tracer.

2: He: Helium: 4.0026 From the Greek word helios, the sun. Janssen obtained the first evidence of helium during the solar eclipse of 1868 when he detected a new line in the solar spectrum. Lockyer and Frankland suggested the name helium for the new element. In 1895 Ramsay discovered helium in the uranium mineral cleveite while it was independently discovered in cleveite by the Swedish chemists Cleve and Langlet at about the same time. Rutherford and Royds in 1907 demonstrated that alpha particles are helium nuclei.

Electron Configuration: 1s2 Atomic Radius: 140 pm (Van der Waals) Melting Point: <-272.2 °C Boiling Point: -268.93 °C Properties Helium has the lowest melting point of any element and is widely used in cryogenic research because its boiling point is close to absolute zero. Also, the element is vital in the study of super conductivity. Using liquid helium, Kurti, co-workers and others have succeeded in obtaining temperatures of a few microkelvins by the adiabatic demagnetization of copper nuclei. Helium has other peculiar properties: It is the only liquid that cannot be solidified by lowering the temperature. It remains liquid down to absolute zero at ordinary pressures, but will readily solidify by increasing the pressure. Solid 3He and 4He are unusual in that both can be changed in volume by more than 30% by applying pressure. The specific heat of helium gas is unusually high. The density of helium vapor at the normal boiling point is also very high, with the vapor expanding greatly when heated to room temperature. Containers filled with helium gas at 5 to 10 K should be treated as though they contained liquid helium due to the large increase in pressure resulting from warming the gas to room temperature. Isotopes Seven isotopes of helium are known: Liquid helium (He-4) exists in two forms: He-4I and He-4II, with a sharp transition point at 2.174K. He-4I (above this temperature) is a normal liquid, but He-4II (below it) is unlike any other known substance. It expands on cooling, its conductivity for heat is enormous, and neither its heat conduction nor viscosity obeys normal rules. Helium is extensively used for filling balloons as it is a much safer gas than hydrogen. One of the recent largest uses for helium has been for pressuring liquid fuel rockets. A Saturn booster, like the type used on the Apollo lunar missions, required about 13 million ft3 of helium for a firing, plus more for checkouts. Liquid helium's use in magnetic resonance imaging (MRI) continues to increase as the medical profession accepts and develops new uses for the equipment. This equipment has eliminated some need for exploratory surgery by accurately diagnosing patients. Another medical application uses MRE to determine (by blood analysis) whether a patient has any form of cancer.

4: Be: Beryllium: 9.0122 From the Greek word beryllos; also sweet. Discovered in the oxide form by Vauquelin in both beryl and emeralds in 1798. The metal was isolated in 1828 by Wohler and by Bussy independently by the action of potassium on beryllium chloride.

Electron Configuration: [He]2s2 Atomic Radius: 153 pm (Van der Waals) Melting Point: 1287 °C Boiling Point: 2471 °C Oxidation States: 2 Sources Beryllium is found in some 30 mineral species, the most important of which are bertrandite, beryl, chrysoberyl, and phenacite. Aquamarine and emerald are precious forms of beryl. Beryl and bertrandite are the most important commercial sources of the element and its compounds. Most of the metal is now prepared by reducing beryllium fluoride with magnesium metal. Beryllium metal did not become readily available to industry until 1957. Properties The metal, steel gray in color, has many desirable properties. As one of the lightest of all metals, it has one of the highest melting points of the light metals. Its modulus of elasticity is about one third greater than that of steel. It resists attack by concentrated nitric acid, has excellent thermal conductivity, and is nonmagnetic. It has a high permeability to X-rays and when bombarded by alpha particles, as from radium or polonium, neutrons are produced in the amount of about 30 neutrons/million alpha particles. At ordinary temperatures, beryllium resists oxidation in air, although its ability to scratch glass is probably due to the formation of a thin layer of the oxide. Uses Beryllium is used as an alloying agent in producing beryllium copper, which is extensively used for springs, electrical contacts, spot-welding electrodes, and non-sparking tools. It is applied as a structural material for high-speed aircraft, missiles, spacecraft, and communication satellites. Other uses include windshield frame, brake discs, support beams, and other structural components of the space shuttle. Because beryllium is relatively transparent to X-rays, ultra-thin Be-foil is finding use in X-ray lithography for reproduction of micro-miniature integrated circuits. Beryllium is used in nuclear reactors as a reflector or moderator for it has a low thermal neutron absorption cross section. It is used in gyroscopes, computer parts, and instruments where lightness, stiffness, and dimensional stability are required. The oxide has a very high melting point and is also used in nuclear work and ceramic applications.

5: B: Boron: 10.81 From the Arabic word Buraq, Persian Burah. Boron compounds have been known for thousands of years, but the element was not discovered until 1808 by Sir Humphry Davy and by Gay-Lussac and Thenard.

Electron Configuration: [He]2s22p1 Atomic Radius: 192 pm (Van der Waals) Melting Point: 2075 °C Boiling Point: 4000 °C Oxidation States: 3 Properties Optical characteristics include transmitting portions of the infrared. Boron is a poor conductor of electricity at room temperature but a good conductor at high temperature. Uses Amorphous boron is used in pyrotechnic flares to provide a distinctive green color, and in rockets as an igniter. By far the most commercially important boron compound in terms of dollar sales is Na2B4O7 • 5H2O. This pentahydrate is used in very large quantities in the manufacture of insulation fiberglass and sodium perborate bleach. Boric acid is also an important boron compound with major markets in textile products. Use of borax as a mild antiseptic is minor in economical terms. Boron compounds are also extensively used in the manufacture of borosilicate glasses. Other boron compounds show promise in treating arthritis. The isotope boron-10 is used as a control for nuclear reactors, as a shield for nuclear radiation, and in instruments used for detecting neutrons. Boron nitride has remarkable properties and can be used to make a material as hard as diamond. The nitride also behaves like an electrical insulator but conducts heat like a metal. Boron also has lubricating properties similar to graphite. The hydrides are easily oxidized with considerable energy liberation, and have been studied for use as rocket fuels. Demand is increasing for boron filaments, a high-strength, lightweight material chiefly employed for advanced aerospace structures. Boron is similar to carbon in that it has a capacity to form stable covalently bonded molecular networks. Carbonates, metalloboranes, phosphacarboranes, and other families comprise thousands of compounds.

6: C: Carbon: 12.011 From the Latin word carbo: charcoal. Carbon, an element of prehistoric discovery, is very widely distributed in nature. It is found in abundance in the sun, stars, comets, and atmospheres of most planets. Carbon in the form of microscopic diamonds is found in some meteorites. Natural diamonds are found in kimberlite of ancient volcanic "pipes," found in South Africa, Arkansas, and elsewhere. Diamonds are now also being recovered from the ocean floor off the Cape of Good Hope. About 30% of all industrial diamonds used in the U.S. are now made synthetically. The energy of the sun and stars can be attributed at least in part to the well-known carbon-nitrogen cycle.

Electron Configuration: [He]2s22p2 Atomic Radius: 170 pm (Van der Waals) Melting Point: 3550 °C (diamond) Boiling Point: 3800°C (sublimation) Oxidation States: 2, 4, -4 Forms Carbon is found free in nature in three allotropic forms: graphite, diamond, and fullerines. A fourth form, known as "white" carbon, is now thought to exist. Ceraphite is one of the softest known materials while diamond is one of the hardest. Graphite exists in two forms: alpha and beta. These have identical physical properties, except for their crystal structure. Naturally occurring graphites are reported to contain as much as 30% of the rhombohedral (beta) form, whereas synthetic materials contain only the alpha form. The hexagonal alpha type can be converted to the beta by mechanical treatment, and the beta form reverts to the alpha on heating it above 1000°C. In 1969 a new allotropic form of carbon was produced during the sublimation of pyrolytic graphite at low pressures. Under free-vaporization conditions above ~2550°K, "white" carbon forms as small transparent crystals on the edges of the planes of graphite. The interplanar spacings of "white" carbon are identical to those of carbon form noted in the graphite gneiss from the Ries (meteroritic) Crater of Germany. "White" carbon is a transparent birefringent material. Little information is presently available about this allotrope. Compounds In combination, carbon is found as carbon dioxide in the atmosphere of the earth and dissolved in all natural waters. It is a component of great rock masses in the form of carbonates of calcium (limestone), magnesium, and iron. Coal, petroleum, and natural gas are chiefly hydrocarbons. Carbon is unique among the elements in the vast number and variety of compounds it can form. With hydrogen, oxygen, nitrogen, and other elements, it forms a very large number of compounds, carbon atom often being linked to another carbon atom. There are close to ten million known carbon compounds, many thousands of which are vital to organic and life processes. Without carbon, the basis for life would be impossible. While it has been thought that silicon might take the place of carbon in forming a host of similar compounds, it is now not possible to form stable compounds with very long chains of silicon atoms. The atmosphere of Mars contains 96.2% CO2. Some of the most important compounds of carbon are carbon dioxide (CO2), carbon monoxide (CO), carbon disulfide (CS2), chloroform (CHCl3), carbon tetrachloride (CCl4), methane (CH4), ethylene (C2H4), acetylene (C2H2), benzene (C6H6), acetic acid (CH3COOH), and their derivatives. Isotopes Carbon has seven isotopes. In 1961 the International Union of Pure and Applied Chemistry adopted the isotope carbon-12 as the basis for atomic weights. Carbon-14, an isotope with a half-life of 5715 years, has been widely used to date such materials as wood, archaeological specimens, etc.

7: N: Nitrogen: 14.007 From the Latin word nitrum, Greek Nitron, native soda; and genes, forming. Nitrogen was discovered by chemist and physician Daniel Rutherford in 1772. He removed oxygen and carbon dioxide from air and showed that the residual gas would not support combustion or living organisms. At the same time there were other noted scientists working on the problem of nitrogen. These included Scheele, Cavendish, Priestley, and others. They called it "burnt" or" dephlogisticated air," which meant air without oxygen.

Electron Configuration: [He]2s22p3 Atomic Radius: 155 pm (Van der Waals) Melting Point: -210.0 °C Boiling Point: -195.79 °C Oxidation States: -3, 5 Nitrogen gas (N2) makes up 78.1% of the Earth's air, by volume. The atmosphere of Mars, by comparison, is only 2.6% nitrogen. From an exhaustible source in our atmosphere, nitrogen gas can be obtained by liquefaction and fractional distillation. Nitrogen is found in all living systems as part of the makeup of biological compounds. The Element The French chemist Antoine Laurent Lavoisier mistakenly named nitrogen azote, meaning without life. However, nitrogen compounds are found in foods, organic materials, fertilizers, poisons, and explosives. Nitrogen, as a gas is colorless, odorless, and generally considered an inert element. As a liquid (boiling point = minus 195.8°C), it is also colorless and odorless, and is similar in appearance to water. Nitrogen gas can be prepared by heating a water solution of ammonium nitrite (NH4NO3). Nitrogen Compounds and Nitrogen in Nature Sodium nitrate (NaNO3) and potassium nitrate (KNO3) are formed by the decomposition of organic matter with compounds of these metals present. In certain dry areas of the world these saltpeters are found in quantity and are used as fertilizers. Other inorganic nitrogen compounds are nitric acid (HNO3), ammonia (NH3), the oxides (NO, NO2, N2O4, N2O), cyanides (CN-), etc. The nitrogen cycle is one of the most important processes in nature for living organisms. Although nitrogen gas is relatively inert, bacteria in the soil are capable of "fixing" the nitrogen into a usable form (as a fertilizer) for plants. In other words, Nature has provided a method to produce nitrogen for plants to grow. Animals eat the plant material where the nitrogen has been incorporated into their system, primarily as protein. The cycle is completed when other bacteria convert the waste nitrogen compounds back to nitrogen gas. Nitrogen is crucial to life, as it is a component of all proteins. Ammonia Ammonia (NH3) is the most important commercial compound of nitrogen. It is produced by the Haber Process. Natural gas (methane, CH4) is reacted with steam to produce carbon dioxide and hydrogen gas (H2) in a two step process. Hydrogen gas and nitrogen gas reacted via the Haber Process to produce ammonia. This colorless gas with a pungent odor is easily liquefied (in fact, the liquid is used as a nitrogen fertilizer). Ammonia is also used in the production of urea, NH2CONH2, which is used as a fertilizer, used in the plastic industry, and used in the livestock industry as a feed supplement. Ammonia is often the starting compound for many other nitrogen compounds.

8: O: Oxygen: 15.999 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.

Electron Configuration: [He]2s22p4 Atomic Radius: 152 pm (Van der Waals) Melting Point: -218.79 °C Boiling Point: -182.95 °C Oxidation States: -2 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%.

9: F: Fluorine: 18.998 From the Latin and French fluere: flow or flux. In 1529, Georigius Agricola described the use of fluorspar as a flux, and as early as 1670 Schwandhard found that glass was etched when exposed to fluorspar treated with acid. Scheele and many later investigators, including Davy, Gay-Lussac, Lavoisier, and Thenard, experimented with hydrofluoric acid, some experiments ending tragically. The element was finally isolated in 1866 by Moissan after nearly 74 years of continuous effort.

Electron Configuration: [He]2s22p5 Atomic Radius: 147 pm (Van der Waals) Melting Point: -219.67 °C Boiling Point: -188.12 °C Oxidation States: -1 Properties Fluorine is the most electronegative and reactive of all elements. It is a pale yellow, corrosive gas, which reacts with most organic and inorganic substances. Finely divided metals, glass, ceramics, carbon, and even water burn in fluorine with a bright flame. Until World War II, there was no commercial production of elemental fluorine. The nuclear bomb project and nuclear energy applications, however, made it necessary to produce large quantities. Uses Fluorine and its compounds are used in producing uranium (from the hexafluoride) and more than 100 commercial fluorochemicals, including many high-temperature plastics. Hydrofluoric acid etches glass of light bulbs. Fluorochlorohydrocarbons are extensively used in air conditioning and refrigeration. The presence of fluorine as a soluble fluoride in drinking water to the extent of 2 ppm may cause mottled enamel in teeth when used by children acquiring permanent teeth; in smaller amounts, however, fluoride helps prevent dental cavities. Elemental fluorine has been studied as a rocket propellant as it has an exceptionally high specific impulse value. Compounds One hypothesis says that fluorine can be substituted for hydrogen wherever it occurs in organic compounds, which could lead to an astronomical number of new fluorine compounds. Compounds of fluorine with rare gases have now been confirmed in fluorides of xenon, radon, and krypton.

10: Ne: Neon: 20.180 From the Greek word neos, new. Discovered by Ramsay and Travers in 1898. Neon is a rare gaseous element present in the atmosphere to the extent of 1 part in 65,000 of air. It is obtained by liquefaction of air and separated from the other gases by fractional distillation.

Electron Configuration: [He]2s22p6 Atomic Radius: 154 pm (Van der Waals) Melting Point: -258.59 °C Boiling Point: -246.08 °C Isotopes Natural neon is a mixture of three isotopes. Six other unstable isotopes are known. Compounds Neon is a very inert element, however, it has been reported to form a compound with fluorine. It is still questionable if true compounds of neon exist, but evidence is mounting in favor of their existence. The ions, Ne+, (NeAr)+, (NeH)+, and (HeNe+) are known from optical and mass spectrometric studies. Neon also forms an unstable hydrate. Properties In a vacuum discharge tube, neon glows reddish orange. It has over 40 times more refrigerating capacity per unit volume than liquid helium and more than three times that of liquid hydrogen. It is compact, inert, and is less expensive than helium when it meets refrigeration requirements. Of all the rare gases, the discharge of neon is the most intense at ordinary voltages and currents. Uses Although neon advertising signs account for the bulk of its use, neon also functions in high-voltage indicators, lightning arrestors, wave meter tubes, and TV tubes. Neon and helium are used in making gas lasers. Liquid neon is now commercially available and is finding important application as an economical cryogenic refrigerant.

3: Li: Lithium: 6.94 From the Greek word lithos, stone. Discovered by Arfvedson in 1817. Lithium is the lightest of all metals, with a density only about half that of water.

Lithium is presently being recovered from brines of Searles Lake, in California, and from those in Nevada. Large deposits of quadramene are found in North Carolina. The metal is produced electrolytically from the fused chloride. Lithium is silvery in appearance, much like Na, K, and other members of the alkali metal series. It reacts with water, but not as vigorously as sodium. Lithium imparts a beautiful crimson color to a flame, but when the metal burns strongly, the flame is a dazzling white. Uses It ranks as a leading contender as a battery anode material as it has a high electrochemical potential. Lithium is used in special glasses and ceramics. The glass for the 200-inch telescope at Mt. Palomar contains lithium as a minor ingredient. Lithium chloride is one of the most hygroscopic materials known, and it, as well as lithium bromide, is used in air conditioning and industrial drying systems. Lithium stearate is used as an all-purpose and high-temperature lubricant. Other lithium compounds are used in dry cells and storage batteries. Lithium carbonate is used for the treatment of bipolar disease and other mental illness conditions.