|
|- | Triple point || 13.8033 K, 7.042 kPa (Bohr radius) |- | 2H || 0.0115% || colspan="4" | H is stable with 1 neutron Hydrogen (IPA: , , from Ancient Greek (hudor): "water" and Ancient Greek (gennao): "to give birth") is a chemical element that, in the periodic table, has the symbol H and an atomic number of 1. At standard temperature and pressure it is a colorless, odorless, nonmetallic, tasteless, highly flammable diatomic gas (H2). With an atomic mass of 1.00794 g/mol, hydrogen is the lightest element. It is also the most abundant, constituting roughly 75% of the universe's elemental mass. Stars in the main sequence are mainly composed of hydrogen in its plasma state. Elemental hydrogen is relatively rare on Earth, and is industrially produced from hydrocarbons, after which most free hydrogen is used "captively" (meaning locally at the production site), with the largest markets about equally divided between fossil fuel upgrading (e.g., hydrocracking) and in ammonia production (mostly for the fertilizer market). However, hydrogen can easily be produced from water using the process of electrolysis. The most common naturally occurring isotope of hydrogen has a single proton and no neutrons. In ionic compounds it can take on either a positive charge (becoming a cation composed of a bare proton) or a negative charge (becoming an anion known as a hydride). Hydrogen can form compounds with most elements and is present in water and most organic compounds. It plays a particularly important role in acid-base chemistry, in which many reactions involve the exchange of protons between soluble molecules. As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics. Nomenclature The word "hydrogen" has several different meanings: The elemental forms of hydrogen should not be confused with hydrogen as it appears in chemical compounds. Discovery of H Hydrogen gas, H2, was first artificially produced and formally described by T. von Hohenheim (also known as Paracelsus, 1493–1541) via the mixing of metals with strong acids. He was unaware that the flammable gas produced by this chemical reaction was a new chemical element. In 1671, Robert Boyle rediscovered and described the reaction between iron filings and dilute acids, which results in the production of hydrogen gas.• In 1766, Henry Cavendish was the first to recognize hydrogen gas as a discrete substance, by identifying the gas from a metal-acid reaction as "inflammable air", and further finding that the gas produces water when burned. Cavendish had stumbled on hydrogen when experimenting with acids and mercury. Although he wrongly assumed that hydrogen was a liberated component of the mercury rather than the acid, he was still able to accurately describe several key properties of hydrogen. He is usually given credit for its discovery as an element. In 1783, Antoine Lavoisier gave the element the name of hydrogen when he (with Laplace) reproduced Cavendish's finding that water is produced when hydrogen is burned. Lavoisier's name for the gas won out. One of the first uses of H2 was for balloons. The H2 was obtained by reacting sulphuric acid and metallic iron. Infamously, H2 was used in the Hindenburg airship that was destroyed in a midair fire. Role in history of quantum theory Because of its relatively simple atomic structure, consisting only of a proton and an electron, the hydrogen atom, together with the spectrum of light produced from it or absorbed by it, has been central to the development of the theory of atomic structure. Furthermore, the corresponding simplicity of the hydrogen molecule and the corresponding cation H2+ allowed fuller understanding of the nature of the chemical bond, which followed shortly after the quantum mechanical treatment of the hydrogen atom had been developed in the mid-1920s. One of the first quantum effects to be explicitly noticed (but not understood at the time) was Maxwell's observation, half a century before full quantum mechanical theory arrived. He observed that the specific heat capacity of H2 unaccountably departs from that of a diatomic gas below room temperature and begins to increasingly resemble that of a monatomic gas at cryogenic temperatures. According to quantum theory, this behavior arises from the spacing of the (quantized) rotational energy levels, which are particularly wide-spaced in H2 because of its low mass. These widely spaced levels inhibit equal partition of heat energy into rotational motion in hydrogen at low temperatures. Diatomic gases composed of heavier atoms do not have such widely spaced levels and do not exhibit the same effect. Natural occurrence
Electron energy levels
Isotopes Hydrogen has three naturally occurring isotopes, denoted 1H, 2H, and 3H. Other, highly unstable nuclei (4H to 7H) have been synthesized in the laboratory but not observed in nature. Small amounts of tritium occur naturally because of the interaction of cosmic rays with atmospheric gases; tritium has also been released during nuclear weapons tests. It is used in nuclear fusion reactions, as a tracer in isotope geochemistry, and specialized in self-powered lighting devices. Tritium was once routinely used in chemical and biological labeling experiments as a radiolabel (this has become less common). Hydrogen is the only element that has different names for its isotopes in common use today. (During the early study of radioactivity, various heavy radioactive isotopes were given names, but such names are no longer used). The symbols D and T (instead of 2H and 3H) are sometimes used for deuterium and tritium, but the corresponding symbol P is already in use for phosphorus and thus is not available for protium). IUPAC states that while this use is common it is not preferred. Elemental molecular forms There are two different types of diatomic hydrogen molecules that differ by the relative spin of their nuclei. In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state; in the parahydrogen form the spins are antiparallel and form a singlet. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form, also known as the "normal form". The equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but since the ortho form is an excited state and has a higher energy than the para form, it is unstable and cannot be purified. At very low temperatures, the equilibrium state is composed almost exclusively of the para form. The physical properties of pure parahydrogen differ slightly from those of the normal form. The ortho/para distinction also occurs in other hydrogen-containing molecules or functional groups, such as water and methylene. The uncatalyzed interconversion between para and ortho H2 increases with increasing temperature; thus rapidly condensed H2 contains large quantities of the high-energy ortho form that convert to the para form very slowly. The ortho/para ratio in condensed H2 is an important consideration in the preparation and storage of liquid hydrogen: the conversion from ortho to para is exothermic and produces enough heat to evaporate the hydrogen liquid, leading to loss of the liquefied material. Catalysts for the ortho-para interconversion, such as iron compounds, are used during hydrogen cooling. Chemical and physical properties The solubility and adsorption characteristics of hydrogen with various metals are very important in metallurgy (as many metals can suffer hydrogen embrittlement) and in developing safe ways to store it for use as a fuel. Hydrogen is highly soluble in many compounds composed of rare earth metals and transition metals and can be dissolved in both crystalline and amorphous metals. Hydrogen solubility in metals is influenced by local distortions or impurities in the metal crystal lattice. Combustion
The Hindenburg Disaster Contrary to common belief, hydrogen may not have been entirely responsible for the infamous Hindenburg tragedy on May 6th, 1937. Though never proven, it may have been a spark from the gasoline powered motors that ignited an aluminium powder in the paint. This caused the Hindenburg to combust very quickly, ultimately killing about one third of its passengers. Although hydrogen, like all combustible gases, is dangerous, its unique ability to rise makes it one of the safest combustible gases. Everyone who perished on the Hindenburg died either from falling to their death or from gasoline burns. The hydrogen had risen away from the scene of the accident within minutes; it was the gasoline, which burned for several hours, that ultimately made the Hindenburg incident such a catastrophe. Compounds Covalent and organic compounds While H2 is not very reactive under standard conditions, it does form compounds with most elements. Millions of hydrocarbons are known, but they are not formed by the direct reaction of elementary hydrogen and carbon. Hydrogen can form compounds with elements that are more electronegative, such as halogens (e.g., F, Cl, Br, I) and chalcogens (O, S, Se); in these compounds hydrogen takes on a partial positive charge. When bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a form of strong noncovalent bonding called hydrogen bonding, which is critical to the stability of many biological molecules. Hydrogen also forms compounds with less electronegative elements, such as the metals and metalloids, in which it takes on a partial negative charge. These compounds are often known as hydrides. Hydrogen forms a vast array of compounds with carbon. Because of their general association with living things, these compounds came to be called organic compounds; the study of their properties is known as organic chemistry and their study in the context of living organisms is known as biochemistry. By some definitions, "organic" compounds are only required to contain carbon (as a classic historical example, urea). However, most of them also contain hydrogen, and since it is the carbon-hydrogen bond which gives this class of compounds most of its particular chemical characteristics, carbon-hydrogen bonds are required in some definitions of the word "organic" in chemistry. (This latter definition is not perfect, however, as in this definition urea would not be included as an organic compound). In inorganic chemistry, hydrides can also serve as bridging ligands that link two metal centers in a coordination complex. This function is particularly common in group 13 elements, especially in boranes (boron hydrides) and aluminum complexes, as well as in clustered carboranes. Hydrides Compounds of hydrogen are often called hydrides, a term that is used fairly loosely. To chemists, the term "hydride" usually implies that the H atom has acquired a negative or anionic character, denoted H−. The hydride anion is theoretically convenient but does not actually exist in practice - alkali metal hydrides, e.g. sodium hydride (NaH), are polymeric and have no solution chemistry. Electrolysis of molten lithium hydride (LiH) produced a stoichiometric quantity of hydrogen at the anode. In lithium aluminum hydride, the AlH4− anion carries hydridic centers firmly attached to the Al(III). Although hydrides can be formed with almost all main-group elements, the number and combination of possible compounds varies widely; for example, there are over 100 binary borane hydrides known, but only one binary aluminum hydride. Binary indium hydride has not yet been identified, although larger complexes exist. "Protons" and acids Oxidation of H2 formally gives the proton, H+. This species is central to discussion of acids, though the term proton is used loosely to refer to positively charged or cationic hydrogen, denoted H+. A bare proton H+ cannot exist in solution because of its strong tendency to attach itself to atoms or molecules with electrons. To avoid the convenient fiction of the naked "solvated proton" in solution, acidic aqueous solutions are sometimes considered to contain the hydronium ion (H3O+) organized into clusters to form H9O4+. Other oxonium ions are found when water is in solution with other solvents. Although exotic on earth, one of the most common ions in the universe is the H3+ ion, known as protonated molecular hydrogen or the triatomic hydrogen cation. Production H2 is produced in chemistry and biology laboratories, often as a by-product of other reactions; in industry for the hydrogenation of unsaturated substrates; and in nature as a means of expelling reducing equivalents in biochemical reactions. Laboratory syntheses In the laboratory, H2 is usually prepared by the reaction of acids on metals such as zinc. Zn + 2 H+ → Zn2+ + H2 Aluminum produces H2 upon treatment with acids but also with base: 2 Al + 6 H2O → 2 Al(OH)3 + 3 H2 The electrolysis of water is a simple method of producing hydrogen, although the resulting hydrogen necessarily has less energy content than was required to produce it. A low voltage current is run through the water, and gaseous oxygen forms at the anode while gaseous hydrogen forms at the cathode. Typically the cathode is made from platinum or another inert metal when producing hydrogen for storage. If, however, the gas is to be burnt on site, oxygen is desirable to assist the combustion, and so both electrodes would be made from inert metals. (Iron, for instance, would oxidize, and thus decrease the amount of oxygen given off.) The theoretical maximum efficiency (electricity used vs. energetic value of hydrogen produced) is between 80–94%. Bellona Report on Hydrogen 2H2O(aq) → 2H2(g) + O2(g) Industrial syntheses Hydrogen can be prepared in several different ways but the economically most important processes involve removal of hydrogen from hydrocarbons. Commercial bulk hydrogen is usually produced by the steam reforming of natural gas.Oxtoby DW, Gillis HP, Nachtrieb NH. (2002). Principles of Modern Chemistry 5th ed. Thomson Brooks/Cole At high temperatures (700–1100 °C; 1,300–2,000 °F), steam (water vapor) reacts with methane to yield carbon monoxide and H2. CH4 + H2O → CO + 3 H2 This reaction is favored at low pressures but is nonetheless conducted at high pressures (20 atm; 600 inHg) since high pressure H2 is the most marketable product. The product mixture is known as "synthesis gas" because it is often used directly for the production of methanol and related compounds. Hydrocarbons other than methane can be used to produce synthesis gas with varying product ratios. One of the many complications to this highly optimized technology is the formation of coke or carbon: CH4 → C + 2 H2 Consequently, steam reforming typically employs an excess of H2O. Additional hydrogen from steam reforming can be recovered from the carbon monoxide through the water gas shift reaction, especially with an iron oxide catalyst. This reaction is also a common industrial source of carbon dioxide: CO + H2O → CO2 + H2 Other important methods for H2 production include partial oxidation of hydrocarbons: CH4 + 0.5 O2 → CO + 2 H2 and the coal reaction, which can serve as a prelude to the shift reaction above: C + H2O → CO + H2 NB. Hydrogen is sometimes produced and consumed in the same industrial process, without being separated. In the Haber process for the production of ammonia (the world's fifth most produced industrial compound), hydrogen is generated from natural gas. Biological syntheses H2 is a product of some types of anaerobic metabolism and is produced by several microorganisms, usually via reactions catalyzed by iron- or nickel-containing enzymes called hydrogenases. These enzymes catalyze the reversible redox reaction between H2 and its component two protons and two electrons. Evolution of hydrogen gas occurs in the transfer of reducing equivalents produced during pyruvate fermentation to water. Water splitting, in which water is decomposed into its component protons, electrons, and oxygen, occurs in the light reactions in all photosynthetic organisms. Some such organisms — including the alga Chlamydomonas reinhardtii and cyanobacteria — have evolved a second step in the dark reactions in which protons and electrons are reduced to form H2 gas by specialized hydrogenases in the chloroplast. Efforts have been undertaken to genetically modify cyanobacterial hydrogenases to efficiently synthesize H2 gas even in the presence of oxygen. Other rarer but mechanistically interesting routes to H2 production also exist in nature. Nitrogenase produces approximately one equivalent of H2 for each equivalent of N2 reduced to ammonia. Some phosphatases reduce phosphite to H2. Applications Large quantities of H2 are needed in the petroleum and chemical industries. The largest application of H2 is for the processing ("upgrading") of fossil fuels, and in the production of ammonia. The key consumers of H2 in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking.• H2 has several other important uses. H2 is used as a hydrogenating agent, particularly in increasing the level of saturation of unsaturated fats and oils (found in items such as margarine), and in the production of methanol. It is similarly the source of hydrogen in the manufacture of hydrochloric acid. H2 is also used as a reducing agent of metallic ores. Apart from its use as a reactant, H2 has wide applications in physics and engineering. It is used as a shielding gas in welding methods such as atomic hydrogen welding. H2 is used as the rotor coolant in electrical generators at power stations, because it has the highest thermal conductivity of any gas. Liquid H2 is used in cryogenic research, including superconductivity studies. Since H2 is lighter than air, having a little more than 1/15th of the density of air, it was once widely used as a lifting agent in balloons and airships. However, this use was curtailed after the Hindenburg disaster convinced the public that the gas was too dangerous for this purpose. Hydrogen's rarer isotopes also each have specific applications. Deuterium (hydrogen-2) is used in nuclear fission applications as a moderator to slow neutrons, and in nuclear fusion reactions. Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects. Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs, as an isotopic label in the biosciences, and as a radiation source in luminous paints. The triple point temperature of equilibrium hydrogen is a defining fixed point on the ITS-90 temperature scale. Hydrogen as an energy carrier Having been used as an ingredient in some rocket fuels for several decades, hydrogen, or more specifically H2, is now widely discussed in the context of energy. Hydrogen is not an energy source, since it is not an abundant natural resource and more energy is used to produce it than can be ultimately extracted from it. However, it could become useful as a carrier of energy, as elucidated in the United States Department of Energy's 2003 report, "Among the various alternative energy strategies, building an energy infrastructure that uses hydrogen — the third most abundant element on the earth's surface — as the primary carrier that connects a host of energy sources to diverse end uses may enable a secure and clean energy future for the Nation." The hydrogen would then locally be converted into usable energy either via combustion of fossil fuels or by electrochemical conversion into electricity in a fuel cell. One theoretical advantage of using H2 as a carrier is the localization and concentration of environmentally unwelcome aspects of hydrogen manufacture. For example, CO2 sequestration could be conducted at the point of H2 production from methane. Hydrogen could also be produced using the electrolysis of water method; however, this is currently three to six times as expensive as production from natural gas. High-temperature electrolysis, which promises greater efficiency, is being investigated. Currently, however hydrogen production is expensive relative to other energy storage chemicals, and the bulk of hydrogen is now produced by the least expensive method, which (as noted) employs methane and which, as currently practiced, creates greenhouse gas emissions. • See also Further reading | |||||||||||||
|
| ||||||||||||||
 |
|
| |