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Edited by Blonic :
5/29/2014 4:39:37 AM
2882
Don't let this thread die. (EDIT: Time for 10,000.)
Seriously, keep bumping it.
600 goal: completed
800 goal: completed
10,000 goal: ???
I'm raising the goal for you kiddies.
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Blimp
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Oxygen From Wikipedia, the free encyclopedia This article is about the chemical element and its most stable form, O 2 or dioxygen. For other forms of this element, see Allotropes of oxygen. For other uses, see Oxygen (disambiguation) and O2 (disambiguation). Page semi-protected Oxygen 8O ↑ O ↓ S nitrogen ← oxygen → fluorine Oxygen in the periodic table Appearance colorless gas; pale blue liquid. Oxygen bubbles rise in this photo of liquid oxygen. A glass bottle half-filled with a bluish bubbling liquid Spectral lines of oxygen General properties Name, symbol, number oxygen, O, 8 Pronunciation /ˈɒksɨdʒən/ ok-si-jən Element category diatomic nonmetal, chalcogen Group, period, block 16 (chalcogens), 2, p Standard atomic weight 15.999(4) Electron configuration [He] 2s2 2p4 2, 6 Physical properties Phase gas Density (0 °C, 101.325 kPa) 1.429 g/L Liquid density at b.p. 1.141 g·cm−3 Melting point 54.36 K, -218.79 °C, -361.82 °F Boiling point 90.188 K, -182.962 °C, -297.332 °F Triple point 54.361 K, 0.1463 kPa Critical point 154.581 K, 5.043 MPa Heat of fusion (O2) 0.444 kJ·mol−1 Heat of vaporization (O2) 6.82 kJ·mol−1 Molar heat capacity (O2) 29.378 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 61 73 90 Atomic properties Oxidation states 2, 1, −1, −2 Electronegativity 3.44 (Pauling scale) Ionization energies (more) 1st: 1313.9 kJ·mol−1 2nd: 3388.3 kJ·mol−1 3rd: 5300.5 kJ·mol−1 Covalent radius 66±2 pm Van der Waals radius 152 pm Miscellanea Crystal structure cubic Oxygen has a cubic crystal structure Magnetic ordering paramagnetic Thermal conductivity 26.58x10-3 W·m−1·K−1 Speed of sound (gas, 27 °C) 330 m·s−1 CAS registry number 7782-44-7 History Discovery Carl Wilhelm Scheele (1772) Named by Antoine Lavoisier (1777) Most stable isotopes Main article: Isotopes of oxygen iso NA half-life DM DE (MeV) DP 16O 99.76% 16O is stable with 8 neutrons 17O 0.039% 17O is stable with 9 neutrons 18O 0.201% 18O is stable with 10 neutrons v t e · references Blue white glow from an oxygen discharge tube. Oxygen is a chemical element with symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table and is a highly reactive nonmetallic element and oxidizing agent that readily forms compounds (notably oxides) with most elements.[1] By mass, oxygen is the third-most abundant element in the universe, after hydrogen and helium.[2] At STP, two atoms of the element bind to form dioxygen, a diatomic gas that is colorless, odorless, and tasteless, with the formula O 2. Many major classes of organic molecules in living organisms, such as proteins, nucleic acids, carbohydrates, and fats, contain oxygen, as do the major inorganic compounds that are constituents of animal shells, teeth, and bone. Most of the mass of living organisms is oxygen as it is a part of water, the major constituent of lifeforms (for example, about two-thirds of human body mass). Elemental oxygen is produced by cyanobacteria, algae and plants, and is used in cellular respiration for all complex life. Oxygen is toxic to obligately anaerobic organisms, which were the dominant form of early life on Earth until O 2 began to accumulate in the atmosphere. Free elemental O 2 only began to accumulate in the atmosphere about 2.5 billion years ago (see Great oxygenation event), about a billion years after the first appearance of these organisms.[3][4] Diatomic oxygen gas constitutes 20.8% of the volume of air.[5] Oxygen is the most abundant element by mass in the Earth's crust as part of oxide compounds such as silicon dioxide, making up almost half of the crust's mass.[6] Oxygen is an important part of the atmosphere, and is necessary to sustain most terrestrial life as it is used in respiration. However, it is too chemically reactive to remain a free element in Earth's atmosphere without being continuously replenished by the photosynthetic action of living organisms, which use the energy of sunlight to produce elemental oxygen from water. Another form (allotrope) of oxygen, ozone (O 3), strongly absorbs UVB radiation and consequently the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation, but is a pollutant near the surface where it is a by-product of smog. At even higher low earth orbit altitudes, atomic oxygen is a significant presence and a cause of erosion for spacecraft.[7] Oxygen is produced industrially by fractional distillation of liquefied air, use of zeolites with pressure-cycling to concentrate oxygen from air, electrolysis of water and other means. Uses of elemental oxygen include the production of steel, plastics and textiles, brazing, welding and cutting of steels and other metals, rocket propellant, oxygen therapy and life support systems in aircraft, submarines, spaceflight and diving. Oxygen was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph Priestley in Wiltshire, in 1774, but Priestley is often given priority because his work was published first. The name oxygen was coined in 1777 by Antoine Lavoisier,[8] whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. Its name derives from the Greek roots ὀξύς oxys, "acid", literally "sharp", referring to the sour taste of acids and -γενής -genes, "producer", literally "begetter", because at the time of naming, it was mistakenly thought that all acids required oxygen in their composition. Contents [hide] 1 Characteristics 1.1 Structure 1.2 Allotropes 1.3 Physical properties 1.4 Isotopes and stellar origin 1.5 Occurrence 1.6 Analysis 2 Biological role of O2 2.1 Photosynthesis and respiration 2.2 Content in body 2.3 Build-up in the atmosphere 3 History 3.1 Early experiments 3.2 Phlogiston theory 3.3 Discovery 3.4 Lavoisier's contribution 3.5 Later history 4 Industrial production 5 Applications 5.1 Medical 5.2 Life support and recreational use 5.3 Industrial 6 Compounds 6.1 Oxides and other inorganic compounds 6.2 Organic compounds and biomolecules 7 Safety and precautions 7.1 Toxicity 7.2 Combustion and other hazards 8 See also 9 Notes 10 Citations 11 References 12 Further reading 13 External links Characteristics Structure Oxygen O2 molecule. At standard temperature and pressure, oxygen is a colorless, odorless gas with the molecular formula O 2, in which the two oxygen atoms are chemically bonded to each other with a spin triplet electron configuration. This bond has a bond order of two, and is often simplified in description as a double bond[9] or as a combination of one two-electron bond and two three-electron bonds.[10] Triplet oxygen (not to be confused with ozone, O 3) is the ground state of the O 2 molecule.[11] The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals.[a] These orbitals are classified as antibonding (weakening the bond order fr
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Edited by Niedopałek: 5/12/2014 6:56:26 PMwell, I was originally gonna copy and paste the Wikipedia articles for all 118 elements, but I'm getting bored, so I'm stopping at magnesium
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Magnesium From Wikipedia, the free encyclopedia Not to be confused with Manganese. Magnesium 12Mg Be ↑ Mg ↓ Ca sodium ← magnesium → aluminium Magnesium in the periodic table Appearance shiny grey solid Spectral lines of Magnesium General properties Name, symbol, number magnesium, Mg, 12 Pronunciation /mæɡˈniːziəm/ mag-nee-zee-əm Element category alkaline earth metal Group, period, block 2 (alkaline earth metals), 3, s Standard atomic weight 24.305(1) Electron configuration [Ne] 3s2 2, 8, 2 Physical properties Phase solid Density (near r.t.) 1.738 g·cm−3 Liquid density at m.p. 1.584 g·cm−3 Melting point 923 K, 650 °C, 1202 °F Boiling point 1363 K, 1091 °C, 1994 °F Heat of fusion 8.48 kJ·mol−1 Heat of vaporization 128 kJ·mol−1 Molar heat capacity 24.869 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 701 773 861 971 1132 1361 Atomic properties Oxidation states +2, +1[1] (strongly basic oxide) Electronegativity 1.31 (Pauling scale) Ionization energies (more) 1st: 737.7 kJ·mol−1 2nd: 1450.7 kJ·mol−1 3rd: 7732.7 kJ·mol−1 Atomic radius 160 pm Covalent radius 141±7 pm Van der Waals radius 173 pm Miscellanea Crystal structure hexagonal close-packed Magnesium has a hexagonal close packed crystal structure Magnetic ordering paramagnetic Electrical resistivity (20 °C) 43.9 nΩ·m Thermal conductivity 156 W·m−1·K−1 Thermal expansion (25 °C) 24.8 µm·m−1·K−1 Speed of sound (thin rod) (r.t.) (annealed) 4940 m·s−1 Young's modulus 45 GPa Shear modulus 17 GPa Bulk modulus 45 GPa Poisson ratio 0.290 Mohs hardness 2.5 Brinell hardness 260 MPa CAS registry number 7439-95-4 History Naming after Magnesia, Greece Discovery Joseph Black (1755) First isolation Humphry Davy (1808) Most stable isotopes Main article: Isotopes of magnesium iso NA half-life DM DE (MeV) DP 24Mg 78.99% 24Mg is stable with 12 neutrons 25Mg 10.00% 25Mg is stable with 13 neutrons 26Mg 11.01% 26Mg is stable with 14 neutrons v t e · references Magnesium is a chemical element with the symbol Mg and atomic number 12. Its common oxidation number is +2. It is an alkaline earth metal and the eighth-most-abundant element in the Earth's crust[2] and ninth in the known universe as a whole.[3][4] Magnesium is the fourth-most-common element in the Earth as a whole (behind iron, oxygen and silicon), making up 13% of the planet's mass and a large fraction of the planet's mantle. The relative abundance of magnesium is related to the fact that it easily builds up in supernova stars from a sequential addition of three helium nuclei to carbon (which in turn is made from three helium nuclei).[citation needed] Due to magnesium ion's high solubility in water, it is the third-most-abundant element dissolved in seawater.[5] Magnesium is produced in stars larger than 3 solar masses by fusing helium and neon in the alpha process at temperatures above 600 megakelvins.[citation needed] The free element (metal) is not found naturally on Earth, as it is highly reactive (though once produced, it is coated in a thin layer of oxide (see passivation), which partly masks this reactivity). The free metal burns with a characteristic brilliant-white light, making it a useful ingredient in flares. The metal is now obtained mainly by electrolysis of magnesium salts obtained from brine. In commerce, the chief use for the metal is as an alloying agent to make aluminium-magnesium alloys, sometimes called magnalium or magnelium. Since magnesium is less dense than aluminium, these alloys are prized for their relative lightness and strength. In human biology, magnesium is the eleventh-most-abundant element by mass in the human body. Its ions are essential to all living cells, where they play a major role in manipulating important biological polyphosphate compounds like ATP, DNA, and RNA. Hundreds of enzymes, thus, require magnesium ions to function. Magnesium compounds are used medicinally as common laxatives, antacids (e.g., milk of magnesia), and in a number of situations where stabilization of abnormal nerve excitation and blood vessel spasm is required (e.g., to treat eclampsia). Magnesium ions are sour to the taste, and in low concentrations they help to impart a natural tartness to fresh mineral waters. In vegetation, magnesium is the metallic ion at the center of chlorophyll, and is, thus, a common additive to fertilizers.[6] Contents [hide] 1 Characteristics 1.1 Physical properties 1.2 Chemical properties 1.3 Occurrence 2 Forms 2.1 Alloy 2.2 Compounds 2.3 Isotopes 3 Production 4 History 5 Applications 5.1 As metal 5.2 In compounds 6 Biological roles 6.1 Detection in biological fluids 6.2 Disease 6.3 Magnesium overdose 7 Safety precautions for magnesium metal 8 See also 9 References 10 External links Characteristics[edit] Physical properties[edit] Elemental magnesium is a rather strong, silvery-white, light-weight metal (two-thirds the density of aluminium). It tarnishes slightly when exposed to air, although, unlike the alkali metals, an oxygen-free environment is unnecessary for storage because magnesium is protected by a thin layer of oxide that is fairly impermeable and difficult to remove. Like its lower periodic table group neighbor calcium, magnesium reacts with water at room temperature, though it reacts much more slowly than calcium. When submerged in water, hydrogen bubbles almost unnoticeably begin to form on the surface of the metal—though, if powdered, it reacts much more rapidly. The reaction occurs faster with higher temperatures (see precautions). Magnesium's ability to react with water can be harnessed to produce energy and run a magnesium-based engine. Magnesium also reacts exothermically with most acids, such as hydrochloric acid (HCl). As with aluminium, zinc, and many other metals, the reaction with HCl produces the chloride of the metal and releases hydrogen gas. Chemical properties[edit] Magnesium is a highly flammable metal, but, while it is easy to ignite when powdered or shaved into thin strips, it is difficult to ignite in mass or bulk. Once ignited, it is difficult to extinguish, being able to burn in nitrogen (forming magnesium nitride), carbon dioxide (forming magnesium oxide, and carbon) and water (forming magnesium oxide and hydrogen). This property was used in incendiary weapons used in the firebombing of cities in World War II, the only practical civil defense being to smother a burning flare under dry sand to exclude the atmosphere. On burning in air, magnesium produces a brilliant-white light that includes strong ultraviolet. Thus, magnesium powder (flash powder) was used as a source of illumination in the early days of photography. Later, magnesium ribbon was used in electrically ignited flashbulbs. Magnesium powder is used in the manufacture of fireworks and marine flares where a brilliant white light is required. Flame temperatures of magnesium and magnesium alloys can reach 3,100 °C (3,370 K; 5,610 °F),[7] although flame
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Sodium From Wikipedia, the free encyclopedia This article is about the chemical element. For the PlayStation Home game, see Sodium (PlayStation Home). For the racehorse, see Sodium (horse). "Natrium" redirects here. For the town in West Virginia, see Natrium, West Virginia. Page semi-protected Sodium 11Na ↑ Na ↓ K neon ← sodium → magnesium Sodium in the periodic table Appearance silvery white metallic Spectral lines of sodium General properties Name, symbol, number sodium, Na, 11 Pronunciation /ˈsoʊdiəm/ soh-dee-əm Element category alkali metal Group, period, block 1 (alkali metals), 3, s Standard atomic weight 22.98976928(2) Electron configuration [Ne] 3s1 2,8,1 Physical properties Phase solid Density (near r.t.) 0.968 g·cm−3 Liquid density at m.p. 0.927 g·cm−3 Melting point 370.944 K, 97.794 °C, 208.029 °F Boiling point 1156.090 K, 882.940 °C, 1621.292 °F Critical point (extrapolated) 2573 K, 35 MPa Heat of fusion 2.60 kJ·mol−1 Heat of vaporization 97.42 kJ·mol−1 Molar heat capacity 28.230 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 554 617 697 802 946 1153 Atomic properties Oxidation states 1, −1 (strongly basic oxide) Electronegativity 0.93 (Pauling scale) Ionization energies (more) 1st: 495.8 kJ·mol−1 2nd: 4562 kJ·mol−1 3rd: 6910.3 kJ·mol−1 Atomic radius 186 pm Covalent radius 166±9 pm Van der Waals radius 227 pm Miscellanea Crystal structure body-centered cubic Sodium has a body-centered cubic crystal structure Magnetic ordering paramagnetic[1] Electrical resistivity (20 °C) 47.7 nΩ·m Thermal conductivity 142 W·m−1·K−1 Thermal expansion (25 °C) 71 µm·m−1·K−1 Speed of sound (thin rod) (20 °C) 3200 m·s−1 Young's modulus 10 GPa Shear modulus 3.3 GPa Bulk modulus 6.3 GPa Mohs hardness 0.5 Brinell hardness 0.69 MPa CAS registry number 7440-23-5 History Discovery Humphry Davy (1807) First isolation Humphry Davy (1807) Most stable isotopes Main article: Isotopes of sodium iso NA half-life DM DE (MeV) DP 22Na trace 2.602 y β+→γ 0.5454 22Ne* 1.27453(2)[2] 22Ne ε→γ - 22Ne* 1.27453(2) 22Ne β+ 1.8200 22Ne 23Na 100% 23Na is stable with 12 neutrons * = excited state v t e · references Sodium is a chemical element with the symbol Na (from Latin: natrium) and atomic number 11. It is a soft, silver-white, highly reactive metal and is a member of the alkali metals; its only stable isotope is 23Na. The free metal does not occur in nature, but instead must be prepared from its compounds; it was first isolated by Humphry Davy in 1807 by the electrolysis of sodium hydroxide. Sodium is the sixth most abundant element in the Earth's crust, and exists in numerous minerals such as feldspars, sodalite and rock salt (NaCl). Many salts of sodium are highly water-soluble, and their sodium has been leached by the action of water so that sodium and chlorine (Cl) are the most common dissolved elements by weight in the Earth's bodies of oceanic water. Many sodium compounds are useful, such as sodium hydroxide (lye) for soap-making, and sodium chloride for use as a de-icing agent and a nutrient (edible salt). Sodium is an essential element for all animals and some plants. In animals, sodium ions are used against potassium ions to build up charges on cell membranes, allowing transmission of nerve impulses when the charge is dissipated. The consequent need of animals for sodium causes it to be classified as a dietary inorganic macro-mineral. Contents [hide] 1 Characteristics 1.1 Physical 1.2 Chemical 1.3 Isotopes 1.4 Occurrence 2 Compounds 2.1 Aqueous solutions 2.2 Electrides and sodides 2.3 Organosodium compounds 3 History 4 Commercial production 5 Applications 5.1 Free element 5.1.1 Heat transfer 6 Biological role 7 Precautions 8 See also 9 References 10 External links Characteristics Physical Sodium at standard temperature and pressure is a soft, silvery metal that can be readily cut with a knife, and is a good conductor of electricity. These properties change dramatically at elevated pressures: at 1.5 Mbar, the color changes from silvery metallic to black; at 1.9 Mbar the material becomes transparent, with a red color; and at 3 Mbar sodium is a clear and transparent solid. All of these high-pressure allotropes are insulators and electrides.[3] When sodium or its compounds are introduced into a flame, they turn it yellow,[4] because the excited 3s electrons of sodium emit a photon when they fall from 3p to 3s; the wavelength of this photon corresponds to the D line at 589.3 nm. Spin-orbit interactions involving the electron in the 3p orbital split the D line into two; hyperfine structures involving both orbitals cause many more lines.[5] Chemical Emission spectrum for sodium, showing the D line. A positive flame test for sodium has a bright yellow color. When freshly cut, sodium has a bright, silvery luster. If exposed to air, the surface rapidly tarnishes, darkening at first and then forming a white coating of sodium hydroxide and sodium carbonate. Sodium is generally less reactive than potassium and more reactive than lithium.[6] Like all the alkali metals, it reacts exothermically with water, to the point that sufficiently large pieces melt to a sphere and may explode; this reaction produces caustic sodium hydroxide and flammable hydrogen gas. When burned in dry air, it mainly forms sodium peroxide as well as some sodium oxide. In moist air, sodium hydroxide results.[7] Sodium metal is highly reducing, with the reduction of sodium ions requiring −2.71 volts.[8] Hence, the extraction of sodium metal from its compounds (such as with sodium chloride) uses a significant amount of energy.[7] However, potassium and lithium have even more negative potentials.[9] Isotopes Main article: Isotopes of sodium 20 isotopes of sodium are known, but only 23Na is stable. Two radioactive, cosmogenic isotopes are the byproduct of cosmic ray spallation: 22Na with a half-life of 2.6 years and 24Na with a half-life of 15 hours; all other isotopes have a half-life of less than one minute.[10] Two nuclear isomers have been discovered, the longer-lived one being 24mNa with a half-life of around 20.2 microseconds. Acute neutron radiation, such as from a nuclear criticality accident, converts some of the stable 23Na in human blood to 24Na; by measuring the concentration of 24Na in relation to 23Na, the neutron radiation dosage of the victim can be calculated.[11] Occurrence 23Na is created in the carbon-burning process in stars by fusing two carbon atoms together; this requires temperatures above 600 megakelvins and a star of at least three solar masses.[12] The Earth's crust contains 2.6% sodium by weight, making it the sixth most abundant element on Earth.[13] Because of its high reactivity, it is never found as a pure element. It is found in many different minerals, some very soluble, such as halite and natron, others much less soluble such as amphibole, and zeolite. The insolubility of certain sodium min
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Neon From Wikipedia, the free encyclopedia This article is about the noble gas. For other uses, see Neon (disambiguation). Neon 10Ne ↑ Ne ↓ Ar fluorine ← neon → sodium Neon in the periodic table Appearance colorless gas exhibiting an orange-red glow when placed in a high voltage electric field Spectral lines of neon in the visible region General properties Name, symbol, number neon, Ne, 10 Pronunciation /ˈniːɒn/ Element category noble gases Group, period, block 18 (noble gases), 2, p Standard atomic weight 20.1797(6) Electron configuration [He] 2s2 2p6 2, 8 Physical properties Phase gas Density (0 °C, 101.325 kPa) 0.9002 g/L Liquid density at b.p. 1.207[1] g·cm−3 Melting point 24.56 K, -248.59 °C, -415.46 °F Boiling point 27.104 K, -246.046 °C, -410.883 °F Triple point 24.556 K, 43.37[2][3] kPa Critical point 44.4918 K, 2.7686[3] MPa Heat of fusion 0.335 kJ·mol−1 Heat of vaporization 1.71 kJ·mol−1 Molar heat capacity 5R/2 = 20.786 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 12 13 15 18 21 27 Atomic properties Oxidation states 1[4], 0 Ionization energies (more) 1st: 2080.7 kJ·mol−1 2nd: 3952.3 kJ·mol−1 3rd: 6122 kJ·mol−1 Covalent radius 58 pm Van der Waals radius 154 pm Miscellanea Crystal structure face-centered cubic Neon has a face-centered cubic crystal structure Magnetic ordering diamagnetic[5] Thermal conductivity 49.1×10−3 W·m−1·K−1 Speed of sound (gas, 0 °C) 435 m·s−1 Bulk modulus 654 GPa CAS registry number 7440-01-9 History Prediction William Ramsay (1897) Discovery William Ramsay & Morris Travers[6] (1898) First isolation William Ramsay & Morris Travers[7] (1898) Most stable isotopes Main article: Isotopes of neon iso NA half-life DM DE (MeV) DP 20Ne 90.48% 20Ne is stable with 10 neutrons 21Ne 0.27% 21Ne is stable with 11 neutrons 22Ne 9.25% 22Ne is stable with 12 neutrons v t e · references Neon is a chemical element with symbol Ne and atomic number 10. It is in group 18 (noble gases) of the periodic table. Neon is a colorless, odorless, inert monatomic gas under standard conditions, with about two-thirds the density of air. It was discovered (along with krypton and xenon) in 1898 as one of the three residual rare inert elements remaining in dry air, after nitrogen, oxygen, argon and carbon dioxide are removed. Neon was the second of these three rare gases to be discovered, and was immediately recognized as a new element from its bright red emission spectrum. The name neon is derived from the Greek word, νέον, neuter singular form of νέος [neos], meaning new. Neon is chemically inert and forms no uncharged chemical compounds. During cosmic nucleogenesis of the elements, large amounts of neon are built up from the alpha-capture fusion process in stars. Although neon is a very common element in the universe and solar system (it is fifth in cosmic abundance after hydrogen, helium, oxygen and carbon), it is very rare on Earth. It composes about 18.2 ppm of air by volume (this is about the same as the molecular or mole fraction), and a smaller fraction in Earth's crust. The reason for neon's relative scarcity on Earth and the inner (terrestrial) planets, is that neon forms no compounds to fix it to solids, and is highly volatile, therefore escaping from the planetesimals under the warmth of the newly ignited Sun in the early Solar System. Even the atmosphere of Jupiter is somewhat depleted of neon, presumably for this reason. Neon gives a distinct reddish-orange glow when used in either low-voltage neon glow lamps or in high-voltage discharge tubes or neon advertising signs.[8][9] The red emission line from neon is also responsible for the well known red light of helium–neon lasers. Neon is used in a few plasma tube and refrigerant applications but has few other commercial uses. It is commercially extracted by the fractional distillation of liquid air. It is considerably more expensive than helium, since air is its only source. Contents [hide] 1 History 2 Isotopes 3 Characteristics 4 Occurrence 5 Applications 6 Compounds 7 See also 8 References 9 External links History[edit] Neon gas-discharge lamps forming the symbol for neon "Ne" Neon (Greek νέον (neon), neuter singular form of νέος meaning "new"), was discovered in 1898 by the British chemists Sir William Ramsay (1852–1916) and Morris W. Travers (1872–1961) in London, England.[10] Neon was discovered when Ramsay chilled a sample of air until it became a liquid, then warmed the liquid and captured the gases as they boiled off. The gases nitrogen, oxygen, and argon had been identified, but the remaining gases were isolated in roughly their order of abundance, in a six-week period beginning at the end of May 1898. First to be identified was krypton. The next, after krypton had been removed, was a gas which gave a brilliant red light under spectroscopic discharge. This gas, identified in June, was named neon, the Greek analogue of "novum", (new), the name Ramsay's son suggested.[11] The characteristic brilliant red-orange color that is emitted by gaseous neon when excited electrically was noted immediately; Travers later wrote, "the blaze of crimson light from the tube told its own story and was a sight to dwell upon and never forget."[12] Finally, the same team discovered xenon by the same process, in June. Neon's scarcity precluded its prompt application for lighting along the lines of Moore tubes, which used nitrogen and which were commercialized in the early 1900s. After 1902, Georges Claude's company, Air Liquide, was producing industrial quantities of neon as a byproduct of his air liquefaction business. In December 1910 Claude demonstrated modern neon lighting based on a sealed tube of neon. Claude tried briefly to get neon tubes to be used for indoor lighting, due to their intensity, but failed, as homeowners rejected neon light sources due to their color. Finally in 1912, Claude's associate began selling neon discharge tubes as advertising signs, where they were instantly more successful as eye catchers. They were introduced to the U.S. in 1923, when two large neon signs were bought by a Los Angeles Packard car dealership. The glow and arresting red color made neon advertising completely different from the competition.[13] Neon played a role in the basic understanding of the nature of atoms in 1913, when J. J. Thomson, as part of his exploration into the composition of canal rays, channeled streams of neon ions through a magnetic and an electric field and measured their deflection by placing a photographic plate in their path. Thomson observed two separate patches of light on the photographic plate (see image), which suggested two different parabolas of deflection. Thomson eventually concluded that some of the atoms in the neon gas were of higher mass than the rest. Though not understood at the time by Thomson, this was the first discovery of isotopes of stable atoms. It was made by using a crud
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Fluorine From Wikipedia, the free encyclopedia page is in the middle of an expansion or major revamping This article or section is in the process of an expansion or major restructuring. You are welcome to assist in its construction by editing it as well. If this article or section has not been edited in several days, please remove this template. This article was last edited by ClueBot NG (talk | contribs) 26 seconds ago. (Purge) Fluorine 9F ↑ F ↓ Cl oxygen ← fluorine → neon Fluorine in the periodic table Appearance gas: very pale yellow liquid: bright yellow solid: transparent (beta), opaque (alpha) Small sample of pale yellow liquid fluorine condensed in liquid nitrogen Liquid fluorine at cryogenic temperatures General properties Name, symbol, number fluorine, F, 9 Pronunciation /ˈflʊəriːn/ fluu-reen, /ˈflʊərɪn/, /ˈflɔəriːn/ Element category diatomic nonmetal Group, period, block 17 (halogens), 2, p Standard atomic weight 18.998403163(6) Electron configuration [He] 2s2 2p5[1] 2, 7 Physical properties Phase gas Density (0 °C, 101.325 kPa) 1.696[2] g/L Liquid density at b.p. 1.505[3] g·cm−3 Melting point 53.48 K, −219.67 °C, −363.41[4] °F Boiling point 85.03 K, −188.11 °C, −306.60[4] °F Triple point 53.48 K, 90[4] kPa Critical point 144.41 K, 5.1724[4] MPa Heat of vaporization 6.51[2] kJ·mol−1 Molar heat capacity (Cp) (21.1 °C) 31[3] J·mol−1·K−1 (Cv) (21.1 °C) 23[3] J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 38 44 50 58 69 85 Atomic properties Oxidation states −1 (oxidizes oxygen) Electronegativity 3.98[1] (Pauling scale) Ionization energies (more) 1st: 1681[5] kJ·mol−1 2nd: 3374[5] kJ·mol−1 3rd: 6147[5] kJ·mol−1 Covalent radius 64[6] pm Van der Waals radius 135[7] pm Miscellanea Crystal structure monoclinic Fluorine has a monoclinic base-centered crystal structure alpha state (low-temperature)[8] Magnetic ordering diamagnetic (−1.2×10−4 (SI)[9][10]) Thermal conductivity 0.02591[11] W·m−1·K−1 CAS registry number 7782-41-4[1] History Naming after the mineral fluorite, itself named after Latin fluo (to flow, in smelting) Discovery André-Marie Ampère (1810) First isolation Henri Moissan[1] (June 26, 1886) Named by Humphry Davy Most stable isotopes Main article: Isotopes of fluorine iso NA half-life DM DE (MeV) DP 18F trace 109.77 min β+ (96.9%) 0.634 18O ε (3.1%) 1.656 18O 19F 100% 19F is stable with 10 neutrons reference[12] v t e · references Fluorine is an extremely reactive and poisonous chemical element with atomic number 9. The lightest halogen and most electronegative element, it exists as a pale yellow diatomic gas at standard conditions. Almost all other elements, including some noble gases, form compounds with fluorine. Fluorite (calcium fluoride, CaF 2), the primary mineral source of fluorine, was first described in 1529; trace amounts of F2 lie embedded within it. Contemporarily the Latin verb fluo, meaning "flow", became associated with fluorite rocks as an additive to metal ores which lowered melting points for smelting. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, with several early experimenters dying or sustaining injuries from their attempts. In 1886, French chemist Henri Moissan succeeded in isolating elemental fluorine using low-temperature electrolysis, a process still employed for modern production. The element is 24th in universal abundance and 13th in terrestrial abundance. Due to the expense of refining pure fluorine, nearly all commercial applications handle it bound to compounds. About half of mined fluorite goes into steelmaking, the rest converted into corrosive hydrogen fluoride, a precursor to various organic fluorides and the critical aluminium refining material cryolite. Organic fluorides have very high chemical and thermal stability, their major uses being refrigerants and – as PTFE – cookware, as well as electrical insulation. Pharmaceuticals such as atorvastatin and fluoxetine also contain fluorine, while the fluoride ion inhibits dental cavities, thus finding use in toothpaste and water fluoridation. Uranium enrichment, the largest application for free fluorine, began in World War II during the Manhattan Project. Global fluorochemical sales amount to over US$15 billion a year. Fluorocarbon gases are generally greenhouse gases, with global-warming potentials 100 to 20,000 (for sulfur hexafluoride) times that of carbon dioxide. Organofluorines persist in the environment due to the carbon–fluorine bond's strength, but the potential health impact of the most persistent such compounds is unclear. While a few plants and bacteria synthesise organofluorine poisons for defence against herbivores, fluorine has no metabolic role in mammals. Contents [hide] 1 Characteristics 1.1 Chemical reactivity 1.2 Toxicity 1.3 Phases 1.4 Electron arrangement 1.5 Isotopes 2 Occurrence 2.1 Universe 2.2 Earth 3 Compounds 3.1 Metal fluorides 3.2 Hydrogen fluoride 3.3 Nonmetal fluorides 3.4 Noble gas compounds 3.5 Organic compounds 4 History 4.1 Early discoveries and etymology 4.2 Isolation 4.3 Application development 5 Industry and applications 5.1 Inorganic fluorides 5.2 Organic fluorochemicals 5.3 Fluorine gas 6 Production of fluorine gas 6.1 Industrial 6.2 Chemical 7 Environmental concerns 7.1 Atmosphere 7.2 Biopersistance 8 Biological aspects 8.1 Natural biochemistry 8.2 Medicine 8.3 Agrichemicals and poisons 9 Fluorine-related hazards 9.1 Hydrofluoric acid 9.2 Fluoride ion 10 See also 11 Notes 12 Sources Characteristics[edit] Chemical reactivity[edit] Main article: Chemical characteristics of fluorine Fluorine's superlative reactivity stems from two reasons.[13] Compared with Cl 2 and Br 2, difluorine's bond energy is much lower, similar to those of weak peroxide bonds;[14][13] elemental fluorine thus dissociates easily in reactions. Conversely, bonds to other atoms are very strong because of fluorine's high electronegativity. Otherwise inert substances like powdered steel, glass fragments and asbestos fibres react quickly with cold fluorine gas, while wood and water spontaneously combust under a fluorine jet.[2][15] External video Bright flames during fluorine reactions Fluorine reacting with caesium Reactions of elemental fluorine with metals require varying conditions: alkali metals cause explosions and alkaline earth metals display vigorous activity in bulk, but most other metals such as aluminium and iron must be powdered to prevent protective metal fluoride layers from passivating,[13] and noble metals require pure fluorine gas at 300–450 °C.[16] Metalloids and some solid nonmetals (sulfur, phosphorus, selenium) burn with a flame in room temperature fluorine.[17][18] Hydrogen sulfide and sulfur dioxide combine readily with fluorine, the latter with chances of exploding, but sulfuric acid exhibits much less activity.[17] Hydrogen, analogous to alkali metals, reacts explosively with fluori
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Nitrogen From Wikipedia, the free encyclopedia Nitrogen 7N ↑ N ↓ P carbon ← nitrogen → oxygen Nitrogen in the periodic table Appearance colorless gas, liquid or solid Liquid nitrogen Spectral lines of nitrogen General properties Name, symbol, number nitrogen, N, 7 Pronunciation /ˈnaɪtrədʒən/ ny-trə-jən Element category diatomic nonmetal Group, period, block 15 (pnictogens), 2, p Standard atomic weight 14.007(1) Electron configuration [He] 2s2 2p3 2, 5 Physical properties Phase gas Density (0 °C, 101.325 kPa) 1.251 g/L Liquid density at b.p. 0.808 g·cm−3 Melting point 63.15 K, −210.00 °C, −346.00 °F Boiling point 77.355 K, −195.795 °C, −320.431 °F Triple point 63.151 K, 12.52 kPa Critical point 126.192 K, 3.3958 MPa Heat of fusion (N2) 0.72 kJ·mol−1 Heat of vaporization (N2) 5.56 kJ·mol−1 Molar heat capacity (N2) 29.124 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 37 41 46 53 62 77 Atomic properties Oxidation states 5, 4, 3, 2, 1, −1, −2, −3 (strongly acidic oxide) Electronegativity 3.04 (Pauling scale) Ionization energies (more) 1st: 1402.3 kJ·mol−1 2nd: 2856 kJ·mol−1 3rd: 4578.1 kJ·mol−1 Covalent radius 71±1 pm Van der Waals radius 155 pm Miscellanea Crystal structure hexagonal Nitrogen has a hexagonal crystal structure Magnetic ordering diamagnetic Thermal conductivity 25.83 × 10−3 W·m−1·K−1 Speed of sound (gas, 27 °C) 353 m·s−1 CAS registry number 7727-37-9 History Discovery Daniel Rutherford (1772) Named by Jean-Antoine Chaptal (1790) Most stable isotopes Main article: Isotopes of nitrogen iso NA half-life DM DE (MeV) DP 13N syn 9.965 min ε 2.220 13C 14N 99.634% 14N is stable with 7 neutrons 15N 0.366% 15N is stable with 8 neutrons v t e · references Nitrogen, symbol N, is the chemical element of atomic number seven. At room temperature, it is a gas of diatomic molecules and is colorless and odorless. Nitrogen is a common element in the universe, estimated at about seventh in total abundance in our galaxy and the Solar System. On Earth, the element is primarily found as the gas molecule; it forms about 78% of Earth's atmosphere. The element nitrogen was discovered as a separable component of air, by Scottish physician Daniel Rutherford, in 1772. Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in converting the N 2 into useful compounds, but at the same time causing release of large amounts of often useful energy when the compounds burn, explode, or decay back into nitrogen gas. Synthetically-produced ammonia and nitrates are key industrial fertilizers and fertilizer nitrates are key pollutants in causing the eutrophication of water systems. Outside their major uses as fertilizers and energy-stores, nitrogen compounds are versatile organics. Nitrogen is part of materials as diverse as Kevlar fabric and cyanoacrylate "super" glue. Nitrogen is a constituent of molecules in every major pharmacological drug class, including the antibiotics. Many drugs are mimics or prodrugs of natural nitrogen-containing signal molecules: for example, the organic nitrates nitroglycerin and nitroprusside control blood pressure by being metabolized to natural nitric oxide. Plant alkaloids (often defense chemicals) contain nitrogen by definition, and thus many notable nitrogen-containing drugs, such as caffeine and morphine are either alkaloids or synthetic mimics that act (as many plant alkaloids do) upon receptors of animal neurotransmitters (for example, synthetic amphetamines). Nitrogen occurs in all organisms, primarily in amino acids (and thus proteins) and also in the nucleic acids (DNA and RNA). The human body contains about 3% by mass of nitrogen, the fourth most abundant element in the body after oxygen, carbon, and hydrogen. The nitrogen cycle describes movement of the element from the air, into the biosphere and organic compounds, then back into the atmosphere. Contents [hide] 1 History and etymology 2 Production 3 Properties 3.1 Isotopes 3.2 Electromagnetic spectrum 3.3 Reactions 4 Occurrence 5 Compounds 6 Applications 6.1 Nitrogen gas 6.2 Liquid nitrogen 6.3 Nitrogen compounds 7 Biological role 8 Safety 9 See also 10 Notes 11 References 12 Bibliography 13 Further reading 14 External links History and etymology[edit] Nitrogen is formally considered to have been discovered by Scottish physician Daniel Rutherford in 1772, who called it noxious air or fixed air.[1][2] The fact that there was a component of air that does not support combustion was clear to Rutherford. Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as "mephitic air" or azote, from the Greek word ἄζωτος azotos, "lifeless".[3] In it, animals died and flames were extinguished. This "mephitic air" consisted mostly of N2, but might also have included more than 1% argon. Lavoisier's name for nitrogen is used in many languages (French, Italian, Polish, Russian, Albanian, etc.) and still remains in English in the common names of many compounds, such as hydrazine and compounds of the azide ion. The English word nitrogen (1794) entered the language from the French nitrogène, coined in 1790 by French chemist Jean-Antoine Chaptal (1756–1832), from the Greek νίτρον nitron, "sodium carbonate" and the French -gène, "producing" from Greek -γενής -genes, "producer, begetter". The gas had been found in nitric acid. Chaptal's meaning was that nitrogen gas is the essential part of nitric acid, in turn formed from saltpetre (potassium nitrate), then known as nitre.[4] This word in the more ancient world originally described sodium salts that did not contain nitrate, and is a cognate of natron.[citation needed] Nitrogen compounds were well known by the Middle Ages. Alchemists knew nitric acid as aqua fortis (strong water). The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold (the king of metals). The earliest military, industrial, and agricultural applications of nitrogen compounds used saltpetre (sodium nitrate or potassium nitrate), most notably in gunpowder, and later as fertilizer. In 1910, Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", an allotrope considered to be monoatomic. The "whirling cloud of brilliant yellow light" produced by his apparatus reacted with quicksilver to produce explosive mercury nitride.[5] For a long time sources of nitrogen compounds were limited. Natural sources originated either from biology or deposits of nitrates produced by atmospheric reactions. Nitrogen fixati
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Carbon From Wikipedia, the free encyclopedia This article is about the chemical element. For other uses, see Carbon (disambiguation). This is a good article. Click here for more information.Page semi-protected Carbon 6C - ↑ C ↓ Si boron ← carbon → nitrogen Carbon in the periodic table Appearance diamond: clear graphite: black Spectral lines of Carbon General properties Name, symbol, number carbon, C, 6 Pronunciation /ˈkɑrbən/ Element category polyatomic nonmetal sometimes considered a metalloid Group, period, block 14, 2, p Standard atomic weight 12.011(1) Electron configuration [He] 2s2 2p2 2, 4 Physical properties Phase solid Density (near r.t.) amorphous:[1] 1.8–2.1 g·cm−3 Density (near r.t.) diamond: 3.515 g·cm−3 Density (near r.t.) graphite: 2.267 g·cm−3 Sublimation point 3915 K, 3642 °C, 6588 °F Triple point 4600 K, 10800[2][3] kPa Heat of fusion 117 (graphite) kJ·mol−1 Molar heat capacity diamond: 6.155 J·mol−1·K−1 Molar heat capacity graphite: 8.517 J·mol−1·K−1 Atomic properties Oxidation states 4, 3[4], 2, 1[5], 0, −1, −2, −3, −4[6] Electronegativity 2.55 (Pauling scale) Ionization energies (more) 1st: 1086.5 kJ·mol−1 2nd: 2352.6 kJ·mol−1 3rd: 4620.5 kJ·mol−1 Covalent radius sp3: 77 pm sp2: 73 pm sp: 69 pm Van der Waals radius 170 pm Miscellanea Crystal structure diamond Carbon has a diamond crystal structure (diamond, clear) simple hexagonal Carbon has a Simple Hexagonal crystal structure (graphite, black) Magnetic ordering diamagnetic[7] Thermal conductivity diamond: 900-2300 W·m−1·K−1 Thermal conductivity graphite: 119-165 W·m−1·K−1 Thermal expansion (25 °C) (diamond) 0.8[8] µm·m−1·K−1 Speed of sound (thin rod) (20 °C) (diamond) 18350 m·s−1 Young's modulus diamond: 1050[8] GPa Shear modulus diamond: 478[8] GPa Bulk modulus diamond: 442[8] GPa Poisson ratio diamond: 0.1[8] Mohs hardness diamond: 10 graphite: 1-2 CAS registry number 7440-44-0 History Discovery Egyptians and Sumerians[9] (3750 BC) Recognized as an element by Antoine Lavoisier[10] (1789) Most stable isotopes Main article: Isotopes of carbon iso NA half-life DM DE (MeV) DP 11C syn 20 min β+ 0.96 11B 12C 98.9% 12C is stable with 6 neutrons 13C 1.1% 13C is stable with 7 neutrons 14C trace 5730 y β− 0.15 0 14N v t e · references Carbon (from Latin: carbo "coal") is the chemical element with symbol C and atomic number 6. As a member of group 14 on the periodic table, it is nonmetallic and tetravalent—making four electrons available to form covalent chemical bonds. There are three naturally occurring isotopes, with 12C and 13C being stable, while 14C is radioactive, decaying with a half-life of about 5,730 years.[11] Carbon is one of the few elements known since antiquity.[12] There are several allotropes of carbon of which the best known are graphite, diamond, and amorphous carbon.[13] The physical properties of carbon vary widely with the allotropic form. For example, diamond is highly transparent, while graphite is opaque and black. Diamond is the hardest naturally-occurring material known, while graphite is soft enough to form a streak on paper (hence its name, from the Greek word "γράφω" which means "to write"). Diamond has a very low electrical conductivity, while graphite is a very good conductor. Under normal conditions, diamond, carbon nanotube and graphene have the highest thermal conductivities of all known materials. All carbon allotropes are solids under normal conditions, with graphite being the most thermodynamically stable form. They are chemically resistant and require high temperature to react even with oxygen. The most common oxidation state of carbon in inorganic compounds is +4, while +2 is found in carbon monoxide and other transition metal carbonyl complexes. The largest sources of inorganic carbon are limestones, dolomites and carbon dioxide, but significant quantities occur in organic deposits of coal, peat, oil and methane clathrates. Carbon forms a vast number of compounds, more than any other element, with almost ten million compounds described to date,[14] which in turn are a tiny fraction of such compounds that are theoretically possible under standard conditions. Carbon is the 15th most abundant element in the Earth's crust, and the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen. It is present in all known life forms, and in the human body carbon is the second most abundant element by mass (about 18.5%) after oxygen.[15] This abundance, together with the unique diversity of organic compounds and their unusual polymer-forming ability at the temperatures commonly encountered on Earth, make this element the chemical basis of all known life. On 21 February 2014, NASA announced a greatly upgraded database for tracking polycyclic aromatic hydrocarbons (PAHs) in the universe. According to scientists, more than 20% of the carbon in the universe may be associated with PAHs, possible starting materials for the formation of life. PAHs seem to have been formed shortly after the Big Bang, are widespread throughout the universe, and are associated with new stars and exoplanets.[16] Contents [hide] 1 Characteristics 1.1 Allotropes 1.2 Occurrence 1.3 Isotopes 1.4 Formation in stars 1.5 Carbon cycle 2 Compounds 2.1 Organic compounds 2.2 Inorganic compounds 2.3 Organometallic compounds 3 History and etymology 4 Production 4.1 Graphite 4.2 Diamond 5 Applications 5.1 Diamonds 6 Precautions 7 Bonding to carbon 8 See also 9 References 10 External links Characteristics Theoretically predicted phase diagram of carbon The different forms or allotropes of carbon (see below) include the hardest naturally occurring substance, diamond, and also one of the softest known substances, graphite. Moreover, it has an affinity for bonding with other small atoms, including other carbon atoms, and is capable of forming multiple stable covalent bonds with such atoms. As a result, carbon is known to form almost ten million different compounds; the large majority of all chemical compounds.[14] Carbon also has the highest sublimation point of all elements. At atmospheric pressure it has no melting point as its triple point is at 10.8 ± 0.2 MPa and 4,600 ± 300 K (~4,330 °C or 7,820 °F),[2][3] so it sublimes at about 3,900 K.[17][18] Carbon sublimes in a carbon arc which has a temperature of about 5,800 K (5,530 °C; 9,980 °F). Thus, irrespective of its allotropic form, carbon remains solid at higher temperatures than the highest melting point metals such as tungsten or rhenium. Although thermodynamically prone to oxidation, carbon resists oxidation more effectively than elements such as iron and copper that are weaker reducing agents at room temperature. Carbon compounds form the basis of all known life on Earth, and the carbon-nitrogen cycle provides some of the energy produced by the Sun and other stars. Although it forms an extraordinary variety of compounds, most forms of carbon are c
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Boron From Wikipedia, the free encyclopedia This article is about the chemical element. For other uses, see Boron (disambiguation). Not to be confused with borium, a tungsten carbide product. Boron 5B - ↑ B ↓ Al beryllium ← boron → carbon Boron in the periodic table Appearance black-brown Boron, shown here in the form of its β-rhombohedral phase (its most thermodynamically stable allotrope)[1] General properties Name, symbol, number boron, B, 5 Pronunciation /ˈbɔərɒn/ Element category metalloid Group, period, block 13, 2, p Standard atomic weight 10.81(1) Electron configuration [He] 2s2 2p1 2, 3 Physical properties Phase solid Liquid density at m.p. 2.08 g·cm−3 Melting point 2349 K, 2076 °C, 3769 °F Boiling point 4200 K, 3927 °C, 7101 °F Heat of fusion 50.2 kJ·mol−1 Heat of vaporization 508 kJ·mol−1 Molar heat capacity 11.087 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 2348 2562 2822 3141 3545 4072 Atomic properties Oxidation states 3, 2, 1[2] (mildly acidic oxide) Electronegativity 2.04 (Pauling scale) Ionization energies (more) 1st: 800.6 kJ·mol−1 2nd: 2427.1 kJ·mol−1 3rd: 3659.7 kJ·mol−1 Atomic radius 90 pm Covalent radius 84±3 pm Van der Waals radius 192 pm Miscellanea Crystal structure rhombohedral Boron has a rhombohedral crystal structure Magnetic ordering diamagnetic[3] Electrical resistivity (20 °C) ~106 Ω·m Thermal conductivity 27.4 W·m−1·K−1 Thermal expansion (25 °C) (β form) 5–7[4] µm·m−1·K−1 Speed of sound (thin rod) (20 °C) 16,200 m·s−1 Mohs hardness ~9.5 CAS registry number 7440-42-8 History Discovery Joseph Louis Gay-Lussac and Louis Jacques Thénard[5] (30 June 1808) First isolation Humphry Davy[6] (9 July 1808) Most stable isotopes Main article: Isotopes of boron iso NA half-life DM DE (MeV) DP 10B 19.9(7)% 10B is stable with 5 neutrons[7] 11B 80.1(7)% 11B is stable with 6 neutrons[7] 10B content may be as low as 19.1% and as high as 20.3% in natural samples. 11B is the remainder in such cases.[8] v t e · references Boron is a chemical element with symbol B and atomic number 5. Because boron is produced entirely by cosmic ray spallation and not by stellar nucleosynthesis,[9] it is a low-abundance element in both the solar system and the Earth's crust. Boron is concentrated on Earth by the water-solubility of its more common naturally occurring compounds, the borate minerals. These are mined industrially as evaporites, such as borax and kernite. Chemically uncombined boron, which is classed as a metalloid, is found in small amounts in meteoroids, but is not found naturally on Earth. Industrially, very pure boron is produced with difficulty, as boron tends to form refractory materials containing small amounts of carbon or other elements. Several allotropes of boron exist: amorphous boron is a brown powder, and crystalline boron is black, extremely hard (about 9.5 on the Mohs scale), and a poor conductor at room temperature. Elemental boron is used as a dopant in the semiconductor industry. The major industrial-scale uses of boron compounds are in sodium perborate bleaches, and (Owens-Corning) Borosilicate glass which it trademarked as Pyrex, with greater strength and breakage resistance (thermal shock resistance) than ordinary soda lime glass. Boron polymers and ceramics play specialized roles as high-strength lightweight structural and refractory materials. Boron compounds are used in silica-based glasses and ceramics to give them resistance to thermal shock. Boron-containing reagents are used as intermediates in the synthesis of organic fine chemicals. A few boron-containing organic pharmaceuticals are used, or are in study. Natural boron is composed of two stable isotopes, one of which (boron-10) has a number of uses as a neutron-capturing agent. In biology, borates have low toxicity in mammals (similar to table salt), but are more toxic to arthropods and are used as insecticides. Boric acid is mildly antimicrobial, and a natural boron-containing organic antibiotic is known.[10] Boron is essential to life. Small amounts of boron compounds play a strengthening role in the cell walls of all plants, making boron necessary in soils. Experiments indicate a role for boron as an ultratrace element in animals, but its role in animal physiology is unknown. Contents [hide] 1 History 2 Characteristics 2.1 Allotropes 2.2 Chemistry of the element 2.2.1 Chemical compounds 2.2.1.1 Organoboron chemistry 2.2.1.2 Compounds of B(I) and B(II) 2.3 Isotopes 2.3.1 Commercial isotope enrichment 2.3.2 Enriched boron (boron-10) 2.3.3 Depleted boron (boron-11) 2.3.4 NMR spectroscopy 2.4 Occurrence 3 Production 3.1 Market trend 4 Applications 4.1 Glass and ceramics 4.2 Detergent formulations and bleaching agents 4.3 Insecticides 4.4 Semiconductors 4.5 Magnets 4.6 High-hardness and abrasive compounds 4.6.1 Boron carbide 4.6.2 Other superhard boron compounds 4.7 Shielding in nuclear reactors 4.8 Other nonmedical uses 4.9 Pharmaceutical and biological applications 4.10 Research areas 5 Natural biological role 5.1 Analytical quantification 5.2 Health issues and toxicity 6 See also 7 References 8 External links History[edit] The word boron was coined from borax, the mineral from which it was isolated, by analogy with carbon, which it resembles chemically. For the etymology of borax, see that article. Sassolite Borax glazes were used in China from AD 300, and some tincal (crude borax) reached the West, where the Persian alchemist Jābir ibn Hayyān seems to mention it in AD 700. Marco Polo brought some glazes back to Italy in the 13th century. Agricola, around 1600, reports the use of borax as a flux in metallurgy. In 1777, boric acid was recognized in the hot springs (soffioni) near Florence, Italy, and became known as sal sedativum, with mainly medical uses. The rare mineral is called sassolite, which is found at Sasso, Italy. Sasso was the main source of European borax from 1827 to 1872, at which date American sources replaced it.[11][12] Boron compounds were relatively rarely used chemicals until the late 1800s when Francis Marion Smith's Pacific Coast Borax Company first popularized these compounds and made them in volume and hence cheap.[13] Boron was not recognized as an element until it was isolated by Sir Humphry Davy[6] and by Joseph Louis Gay-Lussac and Louis Jacques Thénard.[5] In 1808 Davy observed that electric current sent through a solution of borates produced a brown precipitate on one of the electrodes. In his subsequent experiments he used potassium to reduce boric acid instead of electrolysis. He produced enough boron to confirm a new element and named the element boracium.[6] Gay-Lussac and Thénard used iron to reduce boric acid at high temperatures. They showed by oxidizing boron with air that boric acid is an oxidation product of boron.[5][14] Jöns Jakob Berzelius identified boron as an element in 1824.[15] Pure boron was arguably first produced by the Ameri
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Beryllium From Wikipedia, the free encyclopedia This is a good article. Click here for more information. Beryllium 4Be - ↑ Be ↓ Mg lithium ← beryllium → boron Beryllium in the periodic table Appearance white-gray metallic General properties Name, symbol, number beryllium, Be, 4 Pronunciation /bəˈrɪliəm/ bə-ril-ee-əm Element category alkaline earth metal Group, period, block 2 (alkaline earth metals), 2, s Standard atomic weight 9.0121831(5) Electron configuration [He] 2s2 2, 2 Physical properties Phase solid Density (near r.t.) 1.85 g·cm−3 Liquid density at m.p. 1.690 g·cm−3 Melting point 1560 K, 1287 °C, 2349 °F Boiling point 3243 K, 2970 °C, 5338 °F Critical point (extrapolated) 5205 K, MPa Heat of fusion 12.2 kJ·mol−1 Heat of vaporization 292 kJ·mol−1 Molar heat capacity 16.443 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 1462 1608 1791 2023 2327 2742 Atomic properties Oxidation states 2, 1[1] (amphoteric oxide) Electronegativity 1.57 (Pauling scale) Ionization energies (more) 1st: 899.5 kJ·mol−1 2nd: 1757.1 kJ·mol−1 3rd: 14848.7 kJ·mol−1 Atomic radius 112 pm Covalent radius 96±3 pm Van der Waals radius 153 pm Miscellanea Crystal structure hexagonal close-packed Beryllium has a hexagonal close packed crystal structure Magnetic ordering diamagnetic Electrical resistivity (20 °C) 36 nΩ·m Thermal conductivity 200 W·m−1·K−1 Thermal expansion (25 °C) 11.3 µm·m−1·K−1 Speed of sound (thin rod) (r.t.) 12890[2] m·s−1 Young's modulus 287 GPa Shear modulus 132 GPa Bulk modulus 130 GPa Poisson ratio 0.032 Mohs hardness 5.5 Vickers hardness 1670 MPa Brinell hardness 600 MPa CAS registry number 7440-41-7 History Discovery Louis Nicolas Vauquelin (1797) First isolation Friedrich Wöhler & Antoine Bussy (1828) Most stable isotopes Main article: Isotopes of beryllium iso NA half-life DM DE (MeV) DP 7Be trace 53.12 d ε 0.862 7Li γ 0.477 - 9Be 100% 9Be is stable with 5 neutrons 10Be trace 1.36×106 y β− 0.556 10B v t e · references Beryllium is the chemical element with the symbol Be and atomic number 4. Because any beryllium synthesized in stars is short-lived, it is a relatively rare element in the universe. It is a divalent element which occurs naturally only in combination with other elements in minerals. Notable gemstones which contain beryllium include beryl (aquamarine, emerald) and chrysoberyl. As a free element it is a steel-gray, strong, lightweight and brittle alkaline earth metal. Beryllium improves many physical properties when added as an alloying element to aluminium, copper (notably the alloy beryllium copper), iron and nickel.[3] Tools made of beryllium copper alloys are strong and hard and do not create sparks when they strike a steel surface. In structural applications, the combination of high flexural rigidity, thermal stability, thermal conductivity and low density (1.85 times that of water) make beryllium metal a desirable aerospace material for aircraft components, missiles, spacecraft, and satellites.[3] Because of its low density and atomic mass, beryllium is relatively transparent to X-rays and other forms of ionizing radiation; therefore, it is the most common window material for X-ray equipment and components of particle physics experiments.[3] The high thermal conductivities of beryllium and beryllium oxide have led to their use in thermal management applications. The commercial use of beryllium requires the use of appropriate dust control equipment and industrial controls at all times because of the toxicity of inhaled beryllium-containing dusts that can cause a chronic life-threatening allergic disease called berylliosis in some people.[4] Contents [hide] 1 Characteristics 1.1 Physical properties 1.2 Nuclear properties 1.3 Isotopes and nucleosynthesis 1.4 Occurrence 2 Production 3 Chemical properties 4 History 4.1 Etymology 5 Applications 5.1 Radiation windows 5.2 Mechanical applications 5.3 Mirrors 5.4 Magnetic applications 5.5 Nuclear applications 5.6 Acoustics 5.7 Electronic 6 Precautions 7 See also 8 Notes 9 References 10 Further reading 11 External links Characteristics[edit] Physical properties[edit] Beryllium is a steel gray and hard metal that is brittle at room temperature and has a close-packed hexagonal crystal structure.[3] It has exceptional stiffness (Young's modulus 287 GPa) and a reasonably high melting point. The modulus of elasticity of beryllium is approximately 50% greater than that of steel. The combination of this modulus and a relatively low density results in an unusually fast sound conduction speed in beryllium – about 12.9 km/s at ambient conditions. Other significant properties are high specific heat (1925 J·kg−1·K−1) and thermal conductivity (216 W·m−1·K−1), which make beryllium the metal with the best heat dissipation characteristics per unit weight. In combination with the relatively low coefficient of linear thermal expansion (11.4×10−6 K−1), these characteristics result in a unique stability under conditions of thermal loading.[5] Nuclear properties[edit] Natural beryllium, save for slight contamination by cosmogenic radioisotopes, is essentially beryllium-9, which has a nuclear spin of 3/2-. Beryllium has a large scattering cross section for high-energy neutrons, about 6 barns for energies above approximately 10 KeV. Therefore, it works as a neutron reflector and neutron moderator, effectively slowing the neutrons to the thermal energy range of below 0.03 eV, where the total cross section is at least an order of magnitude lower – exact value strongly depends on the purity and size of the crystallites in the material. The single primordial beryllium isotope 9Be also undergoes a (n,2n) neutron reaction with neutron energies over about 1.9 MeV, to produce 8Be, which almost immediately breaks into two alpha particles. Thus, for high-energy neutrons beryllium is a neutron multiplier, releasing more neutrons than it absorbs. This nuclear reaction is:[6] 9 4Be + n → 2(4 2He) + 2n Neutrons are liberated when beryllium nuclei are struck by energetic alpha particles[5] producing the nuclear reaction 9 4Be + 4 2He → 12 6C + n , where 4 2He is an alpha particle and 12 6C is a carbon-12 nucleus.[6] Beryllium also releases neutrons under bombardment by gamma rays. Thus, natural beryllium bombarded either by alphas or gammas from a suitable radioisotope is a key component of most radioisotope-powered nuclear reaction neutron sources for the laboratory production of free neutrons. As a metal, beryllium is transparent to most wavelengths of X-rays and gamma rays, making it useful for the output windows of X-ray tubes and other such apparatus. Isotopes and nucleosynthesis[edit] Main articles: Isotopes of beryllium and beryllium-10 Both stable and unstable isotopes of beryllium are created in stars, but these do not last long. It is believed that most of the stable beryllium in the universe was originally created in the interstellar med
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Lithium From Wikipedia, the free encyclopedia This article is about the chemical element. For the use of lithium as a medication, see Lithium (medication). For other uses, see Lithium (disambiguation). This is a good article. Click here for more information. Lithium H ↑ Li ↓ Na helium ← lithium → beryllium Lithium in the periodic table Appearance silvery-white (shown floating in oil) Spectral lines of lithium General properties Name, symbol, number lithium, Li, 3 Pronunciation /ˈlɪθiəm/ li-thee-əm Element category alkali metal Group, period, block 1 (alkali metals), 2, s Standard atomic weight 6.94(1) Electron configuration [He] 2s1 2, 1 Physical properties Phase solid Density (near r.t.) 0.534 g·cm−3 Liquid density at m.p. 0.512 g·cm−3 Melting point 453.65 K, 180.50 °C, 356.90 °F Boiling point 1603 K, 1330 °C, 2426 °F Critical point (extrapolated) 3220 K, 67 MPa Heat of fusion 3.00 kJ·mol−1 Heat of vaporization 136 kJ·mol−1 Molar heat capacity 24.860 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 797 885 995 1144 1337 1610 Atomic properties Oxidation states +1 (strongly basic oxide) Electronegativity 0.98 (Pauling scale) Ionization energies 1st: 520.2 kJ·mol−1 2nd: 7298.1 kJ·mol−1 3rd: 11815.0 kJ·mol−1 Atomic radius 152 pm Covalent radius 128±7 pm Van der Waals radius 182 pm Miscellanea Crystal structure body-centered cubic Lithium has a body-centered cubic crystal structure Magnetic ordering paramagnetic Electrical resistivity (20 °C) 92.8 nΩ·m Thermal conductivity 84.8 W·m−1·K−1 Thermal expansion (25 °C) 46 µm·m−1·K−1 Speed of sound (thin rod) (20 °C) 6000 m·s−1 Young's modulus 4.9 GPa Shear modulus 4.2 GPa Bulk modulus 11 GPa Mohs hardness 0.6 CAS registry number 7439-93-2 History Discovery Johan August Arfwedson (1817) First isolation William Thomas Brande (1821) Most stable isotopes Main article: Isotopes of lithium iso NA half-life DM DE (MeV) DP 6Li 7.5% 6Li is stable with 3 neutrons 7Li 92.5% 7Li is stable with 4 neutrons 6Li content may be as low as 3.75% in natural samples. 7Li would therefore have a content of up to 96.25%. v t e · references Lithium (from Greek: λίθος lithos, "stone") is a chemical element with symbol Li and atomic number 3. It is a soft, silver-white metal belonging to the alkali metal group of chemical elements. Under standard conditions it is the lightest metal and the least dense solid element. Like all alkali metals, lithium is highly reactive and flammable. For this reason, it is typically stored in mineral oil. When cut open, lithium exhibits a metallic luster, but contact with moist air corrodes the surface quickly to a dull silvery gray, then black tarnish. Because of its high reactivity, lithium never occurs freely in nature, and instead, only appears in compounds, which are usually ionic. Lithium occurs in a number of pegmatitic minerals, but due to its solubility as an ion is present in ocean water and is commonly obtained from brines and clays. On a commercial scale, lithium is isolated electrolytically from a mixture of lithium chloride and potassium chloride. The nuclei of lithium verge on instability, since the two stable lithium isotopes found in nature have among the lowest binding energies per nucleon of all stable nuclides. Because of its relative nuclear instability, lithium is less common in the solar system than 25 of the first 32 chemical elements even though the nuclei are very light in atomic weight.[1] For related reasons, lithium has important links to nuclear physics. The transmutation of lithium atoms to helium in 1932 was the first fully man-made nuclear reaction, and lithium-6 deuteride serves as a fusion fuel in staged thermonuclear weapons.[2] Lithium and its compounds have several industrial applications, including heat-resistant glass and ceramics, high strength-to-weight alloys used in aircraft, lithium batteries and lithium-ion batteries. These uses consume more than half of lithium production. Trace amounts of lithium are present in all organisms. The element serves no apparent vital biological function, since animals and plants survive in good health without it. Non-vital functions have not been ruled out. The lithium ion Li+ administered as any of several lithium salts has proved to be useful as a mood-stabilizing drug in the treatment of bipolar disorder, due to neurological effects of the ion in the human body. Contents [hide] 1 Properties 1.1 Atomic and physical 1.2 Chemistry and compounds 1.3 Isotopes 2 Occurrence 2.1 Astronomical 2.2 Terrestrial 2.3 Biological 3 History 4 Production 5 Uses 5.1 Ceramics and glass 5.2 Electrical and electronics 5.3 Lubricating greases 5.4 Metallurgy 5.5 Other chemical and industrial uses 5.6 Nuclear 5.7 Medicine 6 Precautions 6.1 Regulation 7 See also 8 Notes 9 References 10 External links Properties[edit] Main article: Alkali metal Atomic and physical[edit] alt1 alt2 Lithium pellets covered in white lithium hydroxide (left) and ingots with a thin layer of black nitride tarnish (right) Like the other alkali metals, lithium has a single valence electron that is easily given up to form a cation.[3] Because of this, it is a good conductor of heat and electricity as well as a highly reactive element, though the least reactive of the alkali metals. Lithium's low reactivity compared to other alkali metals is due to the proximity of its valence electron to its nucleus (the remaining two electrons are in lithium's 1s orbital and are much lower in energy, and therefore they do not participate in chemical bonds).[3] Lithium metal is soft enough to be cut with a knife. When cut, it possesses a silvery-white color that quickly changes to gray due to oxidation.[3] While it has one of the lowest melting points among all metals (180 °C), it has the highest melting and boiling points of the alkali metals.[4] Lithium has a very low density of 0.534 g/cm3, comparable with that of pine wood. It is the least dense of all elements that are solids at room temperature, the next lightest solid element (potassium, at 0.862 g/cm3) being more than 60% denser. Furthermore, apart from helium and hydrogen, it is less dense than any liquid element, being only 2/3 as dense as liquid nitrogen (0.808 g/cm3).[note 1][5] Lithium can float on the lightest hydrocarbon oils and is one of only three metals that can float on water, the other two being sodium and potassium. Lithium floating in oil Lithium's coefficient of thermal expansion is twice that of aluminium and almost four times that of iron.[6] It has the highest specific heat capacity of any solid element. Lithium is superconductive below 400 μK at standard pressure[7] and at higher temperatures (more than 9 K) at very high pressures (>20 GPa)[8] At temperatures below 70 K, lithium, like sodium, undergoes diffusionless phase change transformations. At 4.2 K it has a rhombohedral crystal system (with a nine-layer repeat spacing); at hi
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1 ReplyHydrogen From Wikipedia, the free encyclopedia This article is about the chemistry of hydrogen. For the physics of atomic hydrogen, see Hydrogen atom. For other meanings, see Hydrogen (disambiguation). Page semi-protected Hydrogen 1H Hydrogen (diatomic nonmetal) Helium (noble gas) Lithium (alkali metal) Beryllium (alkaline earth metal) Boron (metalloid) Carbon (polyatomic nonmetal) Nitrogen (diatomic nonmetal) Oxygen (diatomic nonmetal) Fluorine (diatomic nonmetal) Neon (noble gas) Sodium (alkali metal) Magnesium (alkaline earth metal) Aluminium (other metals) Silicon (metalloid) Phosphorus (polyatomic nonmetal) Sulfur (polyatomic nonmetal) Chlorine (diatomic nonmetal) Argon (noble gas) Potassium (alkali metal) Calcium (alkaline earth metal) Scandium (transition metal) Titanium (transition metal) Vanadium (transition metal) Chromium (transition metal) Manganese (transition metal) Iron (transition metal) Cobalt (transition metal) Nickel (transition metal) Copper (transition metal) Zinc (transition metal) Gallium (other metals) Germanium (metalloid) Arsenic (metalloid) Selenium (polyatomic nonmetal) Bromine (diatomic nonmetal) Krypton (noble gas) Rubidium (alkali metal) Strontium (alkaline earth metal) Yttrium (transition metal) Zirconium (transition metal) Niobium (transition metal) Molybdenum (transition metal) Technetium (transition metal) Ruthenium (transition metal) Rhodium (transition metal) Palladium (transition metal) Silver (transition metal) Cadmium (transition metal) Indium (other metals) Tin (other metals) Antimony (metalloid) Tellurium (metalloid) Iodine (diatomic nonmetal) Xenon (noble gas) Caesium (alkali metal) Barium (alkaline earth metal) Lanthanum (lanthanide) Cerium (lanthanide) Praseodymium (lanthanide) Neodymium (lanthanide) Promethium (lanthanide) Samarium (lanthanide) Europium (lanthanide) Gadolinium (lanthanide) Terbium (lanthanide) Dysprosium (lanthanide) Holmium (lanthanide) Erbium (lanthanide) Thulium (lanthanide) Ytterbium (lanthanide) Lutetium (lanthanide) Hafnium (transition metal) Tantalum (transition metal) Tungsten (transition metal) Rhenium (transition metal) Osmium (transition metal) Iridium (transition metal) Platinum (transition metal) Gold (transition metal) Mercury (transition metal) Thallium (other metals) Lead (other metals) Bismuth (other metals) Polonium (other metals) Astatine (metalloid) Radon (noble gas) Francium (alkali metal) Radium (alkaline earth metal) Actinium (actinide) Thorium (actinide) Protactinium (actinide) Uranium (actinide) Neptunium (actinide) Plutonium (actinide) Americium (actinide) Curium (actinide) Berkelium (actinide) Californium (actinide) Einsteinium (actinide) Fermium (actinide) Mendelevium (actinide) Nobelium (actinide) Lawrencium (actinide) Rutherfordium (transition metal) Dubnium (transition metal) Seaborgium (transition metal) Bohrium (transition metal) Hassium (transition metal) Meitnerium (unknown chemical properties) Darmstadtium (unknown chemical properties) Roentgenium (unknown chemical properties) Copernicium (transition metal) Ununtrium (unknown chemical properties) Flerovium (unknown chemical properties) Ununpentium (unknown chemical properties) Livermorium (unknown chemical properties) Ununseptium (unknown chemical properties) Ununoctium (unknown chemical properties) - ↑ H ↓ Li - ← hydrogen → helium Hydrogen in the periodic table Appearance colorless gas Purple glow in its plasma state Spectral lines of hydrogen General properties Name, symbol, number hydrogen, H, 1 Pronunciation /ˈhaɪdrədʒən/ hy-drə-jən[1] Element category diatomic nonmetal Group, period, block 1, 1, s Standard atomic weight 1.008(1) Electron configuration 1s1 1 Physical properties Color colorless Phase gas Density (0 °C, 101.325 kPa) 0.08988 g/L Liquid density at m.p. 0.07 (0.0763 solid)[2] g·cm−3 Liquid density at b.p. 0.07099 g·cm−3 Melting point 13.99 K, −259.16 °C, −434.49 °F Boiling point 20.271 K, −252.879 °C, −423.182 °F Triple point 13.8033 K, 7.041 kPa Critical point 32.938 K, 1.2858 MPa Heat of fusion (H2) 0.117 kJ·mol−1 Heat of vaporization (H2) 0.904 kJ·mol−1 Molar heat capacity (H2) 28.836 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 15 20 Atomic properties Oxidation states 1, −1 (amphoteric oxide) Electronegativity 2.20 (Pauling scale) Ionization energies 1st: 1312.0 kJ·mol−1 Covalent radius 31±5 pm Van der Waals radius 120 pm Miscellanea Crystal structure hexagonal Hydrogen has a hexagonal crystal structure Magnetic ordering diamagnetic[3] Thermal conductivity 0.1805 W·m−1·K−1 Speed of sound (gas, 27 °C) 1310 m·s−1 CAS registry number 1333-74-0 History Discovery Henry Cavendish[4][5] (1766) Named by Antoine Lavoisier[6] (1783) Most stable isotopes Main article: Isotopes of hydrogen iso NA half-life DM DE (MeV) DP 1H 99.985% 1H is stable with 0 neutrons 2H 0.015% 2H is stable with 1 neutron 3H trace 12.32 y β− 0.01861 3He v t e · references Hydrogen is a chemical element with chemical symbol H and atomic number 1. With an atomic weight of 1.00794 u, hydrogen is the lightest element on the periodic table. Its monatomic form (H) is the most abundant chemical substance in the universe, constituting roughly 75% of all baryonic mass.[7][note 1] Non-remnant stars are mainly composed of hydrogen in its plasma state. The most common isotope of hydrogen, termed protium (name rarely used, symbol 1H), has a single proton and zero neutrons. The universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, odorless, tasteless, non-toxic, nonmetallic, highly combustible diatomic gas with the molecular formula H2. Since hydrogen readily forms covalent compounds with most non-metallic elements, most of the hydrogen on Earth exists in molecular forms such as in the form of water or organic compounds. Hydrogen plays a particularly important role in acid–base reactions. In ionic compounds, hydrogen can take the form of a negative charge (i.e., anion) known as a hydride, or as a positively charged (i.e., cation) species denoted by the symbol H+. The hydrogen cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds are always more complex species than that would suggest. As the simplest atom known, the hydrogen atom has had considerable theoretical application. For example, the hydrogen atom is the only neutral atom with an analytic solution to the Schrödinger equation. Hydrogen gas was first artificially produced in the early 16th century, via the mixing of metals with acids. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance,[8] and that it produces water when burned, a property which later gave it its name: in Greek, hydrogen means "water-former". Industrial production is mainly from the steam reforming of natural gas, and less often from more energy-intensive hydrogen production methods like the electrolysis of water.[9] Most hydrogen is employed near its production site, with the two largest uses being fossil fuel processing (e.g., hydrocracking) and ammonia production, mostly for the fertilizer market. Hydrogen is a concern in metallurgy as it can embrittle many metals,[10] complicating the design of pipelines and storage tanks.[11] Contents [hide] 1 Properties 1.1 Combustion 1.2 Electron energy levels 1.3 Elemental molecular forms 1.4 Phases 1.5 Compounds 1.5.1 Covalent and organic compounds 1.5.2 Hydrides 1.5.3 Protons and acids 1.6 Isotopes 2 History 2.1 Discovery and use 2.2 Role in quantum theory 3 Natural occurrence 4 Production 4.1 Metal-acid 4.2 Steam reforming 4.3 Thermochemical 4.4 Anaerobic corrosion 4.5 Geological occurrence: the serpentinization reaction 4.6 Formation in transformers 4.7 Xylose 5 Applications 5.1 Consumption in processes 5.2 Coolant 5.3 Energy carrier 5.4 Semiconductor industry 6 Biological reactions 7 Safety and precautions 8 See also 9 Notes 10 References 11 Further reading 12 External links Properties Combustion A black cup-like object hanging by its bottom with blue glow coming out of its opening. The Space Shuttle Main Engine burnt hydrogen with oxygen, producing a nearly invisible flame at full thrust. Hydrogen gas (dihydrogen or molecular hydrogen)[12] is highly flammable and will burn in air at a very wide range of concentrations between 4% and 75% by volume.[13] The enthalpy of combustion for hydrogen is −286 kJ/mol:[14] 2 H2(g) + O2(g) → 2 H2O(l) + 572 kJ (286 kJ/mol)[note 2] Hydrogen gas forms explosive mixtures with air if it is 4–74% concentrated and with chlorine if it is 5–95% concentrated. The mixtures may be ignited by spark, heat or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C (932 °F).[15] Pure hydrogen-oxygen flames emit ultraviolet light and with high oxygen mix are nearly invisible to the naked eye, as illustrated by the faint plume of the Space Shuttle Main Engine compared to the highly visible plume of a Space Shuttle Solid Rocket Booster. The detection of a burning hydrogen leak may require a flame detector; such leaks can be very dangerous. Hydrogen flames in other conditions are blue, resembling blue natural gas flames.[16] The destruction of the Hindenburg airship was an infamous example of hydrogen combustion; the cause is debated, but the visible orange flames were the result of a rich mixture of hydrogen to oxygen combined with carbon compounds from the airship skin. H2 reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding hydrogen halides, hydrogen chloride and hydrogen fluoride, which are also potentially dangerous acids.[17]
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Helium From Wikipedia, the free encyclopedia This article is about the chemical element. For other uses, see Helium (disambiguation). Page semi-protected - ↑ He ↓ Ne hydrogen ← helium → lithium Helium in the periodic table Appearance colorless gas, exhibiting a red-orange glow when placed in a high-voltage electric field Spectral lines of helium General properties Name, symbol, number helium, He, 2 Pronunciation /ˈhiːliəm/ hee-lee-əm Element category noble gases Group, period, block 18 (noble gases), 1, s Standard atomic weight 4.002602(2) Electron configuration 1s2 2 Physical properties Phase gas Density (0 °C, 101.325 kPa) 0.1786 g/L Liquid density at m.p. 0.145 g·cm−3 Liquid density at b.p. 0.125 g·cm−3 Melting point (at 2.5 MPa) 0.95 K, −272.20 °C, −457.96 °F Boiling point 4.222 K, −268.928 °C, −452.070 °F Triple point 2.177 K, 5.043 kPa Critical point 5.1953 K, 0.22746 MPa Heat of fusion 0.0138 kJ·mol−1 Heat of vaporization 0.0829 kJ·mol−1 Molar heat capacity 5R/2 = 20.786 J·mol−1·K−1 Vapor pressure (defined by ITS-90) P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 1.23 1.67 2.48 4.21 Atomic properties Oxidation states 0 Electronegativity no data (Pauling scale) Ionization energies 1st: 2372.3 kJ·mol−1 2nd: 5250.5 kJ·mol−1 Covalent radius 28 pm Van der Waals radius 140 pm Miscellanea Crystal structure hexagonal close-packed Helium has a hexagonal close-packed crystal structure Magnetic ordering diamagnetic[1] Thermal conductivity 0.1513 W·m−1·K−1 Speed of sound 972 m·s−1 CAS registry number 7440-59-7 History Naming after Helios, Greek god of the Sun Discovery Pierre Janssen, Norman Lockyer (1868) First isolation William Ramsay, Per Teodor Cleve, Abraham Langlet (1895) Most stable isotopes Main article: Isotopes of helium iso NA half-life DM DE (MeV) DP 3He 0.000137%* 3He is stable with 1 neutron 4He 99.999863%* 4He is stable with 2 neutrons *Atmospheric value, abundance may differ elsewhere. v t e · references Helium is a chemical element with symbol He and atomic number 2. It is a colorless, odorless, tasteless, non-toxic, inert, monatomic gas that heads the noble gas group in the periodic table. Its boiling and melting points are the lowest among the elements and it exists only as a gas except in extreme conditions. Helium is the second lightest element and is the second most abundant element in the observable universe, being present at about 24% of the total elemental mass, which is more than 12 times the mass of all the heavier elements combined. Its abundance is similar to this figure in the Sun and in Jupiter. This is due to the very high nuclear binding energy (per nucleon) of helium-4 with respect to the next three elements after helium. This helium-4 binding energy also accounts for why it is a product of both nuclear fusion and radioactive decay. Most helium in the universe is helium-4, and is believed to have been formed during the Big Bang. Large amounts of new helium are being created by nuclear fusion of hydrogen in stars. Helium is named for the Greek god of the Sun, Helios. It was first detected as an unknown yellow spectral line signature in sunlight during a solar eclipse in 1868 by French astronomer Jules Janssen. Janssen is jointly credited with detecting the element along with Norman Lockyer. Jannsen observed during the solar eclipse of 1868 while Lockyer observed from Britain. Lockyer was the first to propose that the line was due to a new element, which he named. The formal discovery of the element was made in 1895 by two Swedish chemists, Per Teodor Cleve and Nils Abraham Langlet, who found helium emanating from the uranium ore cleveite. In 1903, large reserves of helium were found in natural gas fields in parts of the United States, which is by far the largest supplier of the gas today. Helium is used in cryogenics (its largest single use, absorbing about a quarter of production), particularly in the cooling of superconducting magnets, with the main commercial application being in MRI scanners. Helium's other industrial uses—as a pressurizing and purge gas, as a protective atmosphere for arc welding and in processes such as growing crystals to make silicon wafers—account for half of the gas produced. A well-known but minor use is as a lifting gas in balloons and airships.[2] As with any gas whose density differs from that of air, inhaling a small volume of helium temporarily changes the timbre and quality of the human voice. In scientific research, the behavior of the two fluid phases of helium-4 (helium I and helium II) is important to researchers studying quantum mechanics (in particular the property of superfluidity) and to those looking at the phenomena, such as superconductivity, produced in matter near absolute zero. On Earth it is relatively rare — 5.2 ppm by volume in the atmosphere. Most terrestrial helium present today is created by the natural radioactive decay of heavy radioactive elements (thorium and uranium, although there are other examples), as the alpha particles emitted by such decays consist of helium-4 nuclei. This radiogenic helium is trapped with natural gas in concentrations up to 7% by volume, from which it is extracted commercially by a low-temperature separation process called fractional distillation. Helium is a finite resource and is one of the few elements with escape velocity, meaning that once released into the atmosphere, it escapes into space.[3][4][5] Contents [hide] 1 History 1.1 Scientific discoveries 1.2 Extraction and use 2 Characteristics 2.1 The helium atom 2.1.1 Helium in quantum mechanics 2.1.2 The related stability of the helium-4 nucleus and electron shell 2.2 Gas and plasma phases 2.3 Solid, liquid, and superfluid phases 2.3.1 Helium I state 2.3.2 Helium II state 3 Isotopes 4 Compounds 5 Occurrence and production 5.1 Natural abundance 5.2 Modern extraction and distribution 5.3 Conservation advocates 6 Applications 6.1 Controlled atmospheres 6.2 Gas tungsten arc welding 6.3 Minor uses 6.3.1 Industrial leak detection 6.3.2 Flight 6.3.3 Minor commercial and recreational uses 6.3.4 Scientific uses 7 Inhalation and safety 7.1 Effects 7.2 Hazards 8 See also 9 References 10 Bibliography 11 External links History Scientific discoveries The first evidence of helium was observed on August 18, 1868 as a bright yellow line with a wavelength of 587.49 nanometers in the spectrum of the chromosphere of the Sun. The line was detected by French astronomer Jules Janssen during a total solar eclipse in Guntur, India.[6][7] This line was initially assumed to be sodium. On October 20 of the same year, English astronomer Norman Lockyer observed a yellow line in the solar spectrum, which he named the D3 Fraunhofer line because it was near the known D1 and D2 lines of sodium.[8] He concluded that it was caused by an element in the Sun unknown on Earth. Lockyer and English chemist Edward Frankland named the element with the Greek
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And I'm here too
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*waves*
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BAMMBI!!
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SORRY I SHOT YOUR MOM BAMBI!
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BAMBI!
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))())&($?5$
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3 RepliesAre we ever going to reach 10,000?
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I have 10 pieces of bacon You grab 3 . Now I still have 10 pieces and you have a broken arm. :3
