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Hydrogen General H-Bonding Hydrides Acids Links General Points Hydrogen forms more compounds than any other element. Hydrogen normally exists as H2 diatomic molecules. Molecular hydrogen is a colourless, odourless gas, virtually insoluble in water. It is easily obtained by the reaction between dilute acids and metals such as Zn or Fe, and by electrolysis of water. Industrially, hydrogen is obtained by the "steam re-forming" of methane over a promoted Nickel catalyst at about 750 oC Hydrogen is not an exceptionally reactive element at low temperatures, because the bond dissociation energy of the molecule is considerably endothermic (434.1 KJ mol-1). Hydrogen burns in air to form water, and it will react explosively with oxygen and with halogens under certain conditions. At high temperatures, hydrogen gas will reduce many oxides to lower oxides, it is also useful for the complete reduction of many metal oxides to the metals. In the presence of iron or ruthenium catalysts at high temperatures and pressure H2 will react with N2 to produce NH3. With electropositive metals and nonmetals, hydrogen forms hydrides. Hydrogen serves as a reducing or hydrogen-transfer agent for a variety of organic and inorganic substances, but a catalyst is required in most cases. The reduction of alkenes to alkanes by hydrogen over a Pt or Ni is a typical example. Three isotopes of hydrogen are known: 1H, 2H (Deuterium orD), and 3H (Tritium or T). Although isotope effects are greatest for hydrogen, the chemical properties of H, D and T are essentially identical, except in matters such as rates and equilibrium constants of reactions. The normal form of the element is the diatomic molecule, with 6 possibilities: H2, D2, T2, HD, HT, DT. Naturally occuring hydrogen contains 0.0156% deuterium, while tritium (formed continuously in the upper atmosphere in nuclear reactions induced by cosmic rays) occurs naturally in only minute amounts that are believed to be of the order of 1 in 1017 and is radioactive. Deuterium, as D2O, is separated from water by fractional distillation or electrolysis and is available in ton quantities for use as a moderator in nuclear reactors. Hydrogen Bonding When Hydrogen is bonded to another atom X, mainly F, O, N, or Cl, such that the X-H bond is quite polar, with H bearing a partial positive charge, it can interact with another negative or electron-rich atom Y, to form what is called a hydrogen bond (H-bond), written as X-H---Y The X-H distance becomes slightly longer, but this bond remains essentially a normal two-electron bond. The H--Y distance is generally longer than that of a normal covalent H-Y bond. In the case of the very strong hydrogen bonds, the X----Y distance becomes quite short and the X-H and Y---H distances come close to being equal. In these cases there are presumably covalent and electrostatic components in both the X-H and Y-H bonds. Experimental evidence for hydrogen bonding came first from comparisons of the physical properties of hydrogen compounds. The apparently abnormal high boiling points of NH3, H2O, and HF are classic examples which imply association of these molecules in the liquid phase. Although physical properties reflecting association are still a useful tool in detecting hydrogen bonding, deeper understanding of H bonds and the determination of their parameters comes from X-ray or neutron diffraction of solids, and from other techniques, i.e. NMR, IR, and Raman spectroscopies. Structural evidence for hydrogen bonding is provided by the X---Y distances which are shorter than the expected van der Waals contact when a hydrogen bond exists. For instance, in crystalline NaHCO3 there are four kinds of O---O distances between HCO3- ions with values of 3.12, 3.15, 3.19, and 2.55 Angstroms. The first three are about equal to twice the van der Waals radius of Oxygen, but the last one indicates a hydrogen bond, O-H--O. When an X-H group enters into hydrogen bonding, the X-H stretching band in the IR spectrum is lowered in frequency, broadened, and increased in integrated intensity. The enthalpies of hydrogen bonds are relatively small in most instances: 20-30 KJ mol-1, as compared with covalent bond enthalpies of 200 KJ mol-1, and up. Nevertheless, these weak bonds can have a profound effect on the properties and chemical reactivity of substances in which they occur. This effect is clearly seen in the case of water, where it would boil at about -100 oC instead of +100 oC if hydrogen bonds did not play their role. Ice and Water The structure of water is very important since it is the medium in which so much chemistry, including the chemistry of life, takes place. The structure of ice is of interest for clues about the structure of water. There are nine known variations of ice, the stability of each depending on temperature and pressure. The ice formed in equilibrium with water at 0 oC and 1 atm is called ice I. The structure is an infinite array of oxygen atoms, each tetrahedrally surrounded by four others with hydrogen bonds linking each pair. The structural nature of liquid water is still controversial. The structure is not random, as found in liquids consisting of more-or-less spherical nonpolar molecules; instead, it is highly structured owing to the persistence of hydrogen bonds. Even at 90 oC only a few percent of the water molecules appear not to be hydrogen bonded. Still, there is considerable disorder, or randomness, as befits a liquid. In one model of liquid water the liquid consists at any instant of an imperfect network, very similar to the network of ice I, but differing in that (a) some interstices contain water molecules that do not belong to the network but, instead, disturb it; (b) the network is patchy and does not extend over long distances without breaks; (c) the short-range ordered regions are constantly disintegrating and re-forming; and (d) the network is slightly expanded compared with ice I. The fact that water has a slightly higher density than ice I may be attributed to the presence of enough interstitial water molecules to offset the expansion and disordering of the ice I network. This model of water receives support from X-ray scattering studies. Hydrides Although all compounds of hydrogen could be termed hydrides, not all hydrogen-containing compounds display Hydridic character. In general, hydridic substances are those that either react as hydride ion (H-) donors or clearly contain anionic hydrogen. Thus it is necessary to distinguish hydridic substances (e.g. NaH) from those that are either neutral (e.g. CH4) or acidic (e.g. HCl). This distinction between hydrogen-containing substances that are hydridic, neutral, or acids runs roughly parallel to the bonding that hydrogen can undergo; that is, hydrogen may be bound in it's compounds essentially as (or serve, on reaction, as a source of) H-, H* (radical), or H+, respectively. It is also, at times, convenient to classify the compounds of hydrogen as being: Either ionic or covalent. Either stoichiometric or nonstoichiometric. Either binary or complex. Covalent Hydrides Briefly, the covalent hydrides include: Neutral, binary XH4 compounds, for example, CH4. Somewhat basic, binary XH3 compounds, for example, NH3 and PH3. Weakly acidic or amphoteric, binary XH2 compounds, for example, H2S and H2O. Strongly acidic, binary HX compounds, for example, HCl and HI. Numerous covalent hydrides of Boron Hydridic, complex compounds of hydrogen, two examples of which are LiAlH4 and NaBH4, which serve as powerful reducing agents despite the fact that the Al-H and B-H bonds in these substances are essentially covalent in nature. The latter two compounds provide an interesting illustration of covalent hydrides that are hydridic. First of all, although the two compounds are ionic (being Li+ and Na+ salts), the tetrahedral anions in these salts contain essentially covalent bonds to hydrogen. Furthermore, the tetrahydroaluminate and tetrahydroborate anions are each hydridic, being formed by the action of LiH and Al2Cl6 in ether, and by the action of NaH on diborane. Also, both of the tetrahydro anions are powerful hydrogen-transfer agents. Saline Hydrides The most electropositive elements, the alkali metals and the larger of the alkaline earth metals, react directly with dihydrogen to form stoichiometric hydrides having considerable ionic character. These compounds are called the saline (saltlike) hydrides. Those of the heavier metals are truly hydridic substances since they are properly considered to contain metal cations and H- ions. However, due to the small size and high charge density of the ions of the smaller metals (Be and Mg in Group 2 and Li in Group 1), their hydrides have more covalent character, and BeH2 is best described as a covalent polymer having Be-H-Be bridges. The saline hydrides are ionic substances, as shown by the facts that (a) they conduct electricty when molten, and (b) when dissolved and electrolyzed in molten halides, the saline hydrides evolve dihydrogen at the positive electrode (anode), where oxidation of H- takes place. The ionic nature of the saline hydrides is further indicated by their structures. The ionic radius of H- lies between that of F- and Cl-, and the alkali metals all adopt the NaCl structure. The saline hydrides are all prepared by the direct interaction of the metals with elemental hydrogen at 300-700 oC. The rates of reaction are in the order Li > Cs > K > Na. When pure the compounds are white, but are usually grey due to traces of the metals from which they were made. All of the saline hydrides decompose thermally to give the metal and hydrogen, although Lithium hydride alone is stable to it's melting point (688 oC). Also, only LiH is unreactive at moderate temperatures towards oxygen or chlorine. Because of it's relative unreactivity, LiH finds practical use only in the synthesis of LiAlH4. Since they are hydridic, the saline hydrides (except LiH) are quite reactive with water and air. The saline hydrides are powerful reducing or hydrogen-transfer agents. Acids and Bases The concepts of acidity and basicity are so important in Chemistry that acids and bases have been defined many times and in a variety of ways. One definition, and possibly the oldest, is so narrow as to take water as the only solvent. According to this definition, acids and bases are sources of H+ and OH- respectively. A definition that spawned from this idea is the Bronsted-Lowry definition, it gives a broader definition which is applicable to all protonic solvents. The Bronsted-Lowry Definition An acid is a substance that supplies protons and a base is a proton acceptor. Thus, in water, any substance that increases the concentration of hydrated protons (H3O+) above that due to autodissociation of the water is an acid, and any substance that lowers it is a base. Any solute that supplies hydroxide ions (OH-) is a base, since these combine with protons to reduce the H3O+ concentration. However, other substances, such as sulphides, oxides, or anions of weak acids (F- or CN-), are also bases. The Lewis Definition This is one of the most general (and useful) of all definitions, and was proposed by G.N.Lewis. He defined an acid as an electron-pair acceptor and a base as an electron-pair donor. This definition includes the Bronsted-Lowry definition as a special case, since the proton can be regarded as an electron-pair acceptor and the base, be it OH-, NH2-, HSO4-, and so on, as an electron-pair donor. Measuring Acidity An important characteristicof hydrogen compounds (HX) is the extent to which they ionise in water or other solvents, that is, the extent to which they act as acids. The strength of an acid depends not only on the nature of the acid itself but very much on t he medium in which it is dissolved. For example CF3CO2H and HClO4 are strong acids in water, but in 100% H2SO4 , the former is non-acidic and the latter only a very weak acid. Similarly H3PO4 is a base in 100% H2SO4. Although acidity can be measured in a wide variety of solvents, the most important is water, for which the pH scale is useful. The pH scale has a range of -1 to 15 and is calculate using the concentration of H+ ions present. pH = - log10[H+] This however, is only useful when using water as a solvent. Other ways are to calculate the pka value by dividing Go by 5.71 . Here we can compare the acidity of a compound in various different solvents. Some Useful Links Proton NMR What is Water? Lithium Aluminium Hydride Got any comments? Mail me!!!! General H-Bonding Hydrides Acids Top of Page