R. T. Sanderson
Arizona State
University
Tempe
J
Principles of Hydrogen Chemistry
compounds of hydrogen possihly outnumber even those of carbon and are as varied as chemistry itself. Underneath this vast and seemingly heterogeneous array lies the constancy of hydrogen itself, a powerful unifying influence to lend stability and organization to hydrogen chemistry. This constant quality of hydrogen permits the recognition of a set of basic facts and principles, relatively few in number, from which practically all the varied chemistry of this element can be understood. This paper is an attempt to recognize and to summarize these factsand principles, illustrating their meaning, where appropriate, by specific examples. The hydrogen atom consists of one proton, as a nucleus, and one electron in the first principal quantum level. The capacity of the jirst principal ouantum level being two, the hydrogen nucleus i s n,ot completely shielded by the one electron, and therefore can attract another. This can be accommodated in the single (s) orbital of the $rst shell. A hydrogen atom can thus form one covalent bond, Since it possesses the prerequisite of one outer-shell half-filled orbital. Formation of one single covalent bond by a hydrogen atom completely uses all of the atom's cmventional bonding capacity. It has no additional electrons and no orbital vacancies, after a single covalent bond has been formed. A hydrogen molecule, Hz, in which two atoms have united through such a bond, is therefore inert, being unable to donate, share, or accept electrons. The molecule must he dissociated or, a t the least, its bond must first become highly deformed before the hydrogen can react with any other substance. Association among hydrogen molecules is limited to van der Waals interactions, which with only two electrons per molecule are extremely weak, resulting in extremely low melting and boiling temperatures for hydrogen. Hydrogen i s the m l y element whose atoms have no electrons other than that used far valence. Interactions among all other kinds of atoms are influenced by repulsions among the electron clouds that underly the valence shell, and by shielding of the nuclei by such clouds. Only hydrogen atoms lack these influences. On the basis that its outer shell i s half-jilled, hydrogen belongs in the periodic table near 0 t h elements (Group I V - A ) whose outer shell i s half-$lied. Hydrogendoes not logically belong with the alkali metals, by virtue of their having also one outermost electron per atom, hecause they exceed it by having six more vacancies. Likewise, it does not belong with the halogens because, unlike them, it does not have three pairs of electrons per atom in the outermost shell, besides the valence electron. Since a hydrogen atom can form but one covalent bond, whereas IV-A elements can form four covalent bonds per atom, hydrogen should actually occupy a separate position in the periodic table. Most
logically, however, this should be near to Group IV-A which hydrogen resembles in electronegativity. Hydrogen has a n electronegativity only slightly above the median. I t can therefore act as oxidizing agent toward less electronegative elements, and as reducing agent toward more electronegative elements. The order of increasing electronegativity is: Cs, Rb, K, Na, Li, Ba, Sr, Ca, Mg, Be, Al, Cd, Si, B, Zn, In, Hg, T1, Pb, Sn, Bi, P, Sb, H, Ge, Te, C, I, As, S, Se, N, Br, C1,0, and F. Toward any of the elements to the left of hydrogen, it acts as oxidizing agent. Toward any of the elements to the right - of hydrogen, . - . it acts as reducing -agent. . Hydrogen forms conventional binary compounds m l y with the rare-gas-shell and 18-shell type elements. It f o m compounds with each of these. These compounds have the composition expected from the capacity of each atom to form conventional single covalent bonds. Their type formulas for each major group of the periodic table are: I-A, E H ; 11-A, EH,; 111-A, EHa. IV-A, EHs V-A EHs; VI-A EHz; and VII-A, EH. The strength of covalent bonds to hydrogen appears to be gveatest for atoms that are relatively small and to decrease with increasing size (principal quantum number of the valence shell) of the othe~atom. The most stable binary hydrogen compounds are those of carbon, nitrogen, oxygen, and fluorine. Within each major group of the periodic table, the stability tends to diminish with increasing size of the atom joined to hydrogen. Hydrogen compounds of mercury, thallium, lead, and bismuth are so unstable that knowledge of them is scant. Hydrogen telluride is least stable in its group (possibly excepting polonium hydride), and hydrogen iodide is least stable in its group (possibly excepting hydrogen astatide). When hydrogen acts as a n oxidizing agent, it acquires partial negative charge. I n proportion to the amount of this charge, hydrogen compounds tend to become associated, with consequent reduction of volatility and increase in melting print. The table lists the binary compounds of hydrogen in order of decreasing negative and increasing positive partial charge. The hydrides of the alkali and alkaline earth metals, for example, form ionic crystalline lattices, most of which remain solid and nonvolatile to fairly high temperatures, when they decompose before melting. Hydrides of other elements less electronegative than hydrogen, in which the bonds are not as polar as in the salt-like hydrides, appear to become associated through hydridic bridging, in which hydrogen atoms bridge two molecular units. This may extend to highly polymeric solids, as in BeH2, MgH1, AIHp, In&, or it may result merely in the formation of volatile dimers, such as BzHBand possibly Ga2He. Volume 41, Number 6, June 1964
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Estimated Partial Charge an Combined Hydrogen (in electrons)
Negative CsH RbH KH NaH LiH BaH* SrH* CnHn MgH1 BeH* AIHa CdHl ZnH2 RnRs InHa SiH,
-0.59 -0.58 -0.58 -0.50 -0.50 -0.33 -0.31 -0.29 -0.18 -0.13 -0.12 -0.09 -0.06 -0.05 -0.05 -0.05
Nearly neutral SnH, -0.02 GeH4 -0.01 PA, -0.01 S ~ H I -0.01 CH4 0.01 H5e 0.01 AsHz 0.0'2
Positive HI HrS HzSe SHS HBr H1O HC1 HF
0.04 0.05 0.06 0.06 0.12 0.12 0.16 0.25
When conventional covalence would result i n binary hydrogen compound molecules having one or more outershell vacant orbitals and the hydrogen bears partial negative charge insumient to form a n ionic lattice, these orbitals tend to share the electrons of the hydrogen atoms of other molecules, resulting in an association of the molecules through hydridic bridging. The most familiar example of hydridic bridging occurs in dihoraue, BEHB. Each boron atom has four outer orbitals but only three outer electrons. When it has joined to three hydrogen atoms by conventional covalent bonds, then it still possesses an outer vacant orbital. I n the presence of some electron donor, this orbital quickly becomes occupied through formation of a molecular addition compound, such as BH3:N(CH3)> When only BHI groups would be present, however, they pair completely. One hydrogen atom of each BHI appears to lend the pair of bonding electrons to the other boron in its otherwise unoccupied orbital. Thus is formed a three-center bond, in which one pair of electrons holds two boron atoms to the same hydrogen through one orbital from each boron and one on the hydrogen. Two such three-center bonds appear to comprise the hydridic bridging in diborane. Such hydrogen bridging is called "hydridic" because negative charge on hydrogen seems to be required. Similar situations in which the hydrogen is not partially negative; such as BC1H2 (partial charge on H, 0.03), do not exhihit the dimerization characteristic of diborane. I n proportion to the amount of its negative charge, hydrogen loses oxidizing power and acquires reducing power i n the compound and tends to gain electron donating ability. All compounds in which hydrogen is negatively charged tend to be active reducing agents, although the relative potential reducing power may be obscured by differences in reaction mechanism. Even water is an effective oxidizing agent, with its positive hydrogen; and most compounds of negative hydrogen are readily hydrolyzed, with the liberation of molecular hydrogen formed by uniting negative hydrogen with a proton. Highly negative hydrogen can become a hydride ion ligand, contributing a pair of electrons to an available orbital of another atom. For example, it can join with a BHa group to form the horohydride ion, BH4-, or with an AIHa group to form the aluminohydride ion, AlH4-. To act as ligand, the hydrogen apparently must have a higher negative charge than the hydrogen atoms already attached to the boron or aluminum. 332
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When hydrogen acts as a reducing agent, it acquires partial positive charge. I n proportion to the amount of this positive charge, hydrogen gains oxidizing power and loses reducing power in the compound and, as a releasable proton, tends to gain electron accepting ability or, in other words, becomes acidic. All compounds of hydrogen in which the hy&ogen has partial positive charge are at least potentially oxidizing and acidic; they tend to be more so, the more positive the charge. However, the most important factor, in any reaction involving the release of a proton from a hydrogen compound, is the electron-donating capacity of the negative ion that would remain. The transfer of a proton is a competitive process, essentially between an electron pair on each of two atoms. The relative availability of each electron pair to the proton determines the extent of transfer, if any. Therefore, although the hydrogen tends to be more protonic, the higher its positive charge, a high charge is not necessarily a criterion for high acidity under all conditions. For example, in the hydrogen halides, the acidity of aqueous solutions increases from hydrofluoric acid, which is weak, to hydriodic acid which is very strong. (Aqueous hydrochloric and hydrohromic acids are also strong, but in more acidic solvents than water the distinction can he made that HI is stronger than HBr which in turn is stronger than HC1.) Bond polarity is in the opposite order, the highest positive charge on hydrogen occurring in HF, and in HI the hydrogen is only slightly positive. The difference here lies in the capacity of the halide ion in water to donate an electron pair to a proton. The fluoride ion does this so effectively that H F is a weak acid. The iodide ion does this poorly, as do also chloride and bromide ions, making HC1, HBr, and HI strong acids. An explanation is suggested below. Positive hydrogen, of course, is important not just in water solutions for it possesses oxidizing power in the compounds separate from water. Such compounds can, for example, remove electrons from metals and react with compounds of negative hydrogen, releasing molecular hydrogen gas, or with compounds of negative alkyl, releasing the corresponding alkane. Positively charged hydrogen can become attracted to an electron pair on another molecule (or on another atom of the same molecule) electrostatically, with s u f k e n t force to form a weak bond, called a "protonic bridge," holding the molecul~stogether. This is what is commonly called a "hydrogen bond." The term "protonic bridge" seems preferable to indicate that positive charge on the hydrogen is necessary and to distinguish it from the different type of hydrogen bonding called "hydridic" bridging. Protonic bridging is a property exclusively of hydrogen, because of the absence of electrons underlying the valence shell, in hydrogen alone. Whenthe valence electrons are partially withdrawn from the hydrogen, the nucleus becomes somewhat exposed on the far side of the bond and can exert an electrostatic attraction for an electron pair of another atom if that pair is sufficiently available. I n order t o be available, the atom on which the pair resides must have partial negative charge, and it must be small. The requirement of negative charge means that the electrons are more easily available, being farther removed from their own atomic nucleus and that any repulsive effect of that nucleus toward the positive hydrogen is effectively shielded. The require-
ment of small size seems to suggest that in larger atoms, even when they bear partial negative charge, the electron pairs are not concentrated within so small a space as in the small atoms and hence not so available to a proton of another molecule. These requirements practically restrict the effectiveness of protonic bridging to compounds in which nitrogen, oxygen, and fluorine bear partial negative charge and are available to positively charged hydrogen. The hydrogen is usually attached to one of these elements, too, but may be attached to some other element such as carbon as long as the hydrogen bears substantial positive charge. The physical properties of ammonia, water, and hydrogen fluoride that reflect the state of aggregation of the molecules, such as melting and boiling points, give ample evidence of protonic bridging in these compounds. The essential absence of such bridging in their congeners is similarly shown by their much lower boiling and melting points, despite their greater molecular weight. Compounds in which the hydrogen is neither appreciably positive nor appreciably negative show no evidence of association through either protonic or hydridic bridging and tend, therefore, to be low-melting and volatile. Only a few elements are close to hydrogen in electronegat i d y , and therefore compounds of relatively uncharged hydrogen are restricted to these. They include all hydrocarbons as well as methane, germanium hydrogen compounds, phosphine PHI, arsine AsH8, and stibine SbHa, and hydrogen telluride also. Compounds in which hydrogen i s close to neutral also are relatively inert chemically at ordinary temperatures, being not easily oxidized or reduced and neither acidic nor basic. Methane and germane, for example, although they will burn in air when ignited, are not spontaneously inflammable as are many compounds of negative hydrogen, and they are inert toward hydrolysis. They are neither acidic nor basic. Hydrogen and carbon are the only two elaenta whose atoms, when they have formed all the covalent bonds of which they are capable, have neither outermost vacancies nor outermost electron pairs available for further reaction. This is highly significant, accounting for the relative unreactivity of the alkanes and of alkyl groups throughout organic chemistry. Such molecules and groups are relatively invulnerable to chemical attack, which commonly requires the initial formation of unstable intermediates through the use of nnshared pairs of orbital vacancies. Outer d orbitals, not present in carbon, are available in all other elements of this group, imparting higher reactivity to their compounds. The small size of carbon permits strong bonds to hydrogen, and the similarity in electronegatitity prevents ap-
preciable bond polarity in hydrocarbons. These facts contribute greatly to the importance of hydrocarbon chemistry, for they permit indefinitely extensive joining together of carbon atoms when the remaining bonds are to hydrogen. The "uniqueness" of organic chemistry is not the result of some highly special quality of carbon, as is often implied, but rather, the resultant of a particular combination of qualities and circumstances, none unusual by itself. In summary, these include: the small radius of carbon so that its bonds to hydrogen are stable; the electmnegativity similarity between carbon and hydrogen; the unavailability of orbital vacancies or unshared electron pairs on either carbon or hydrogen when attached by single covalent bonds; the ability of carbon to form multiple bonds using its p orbitals; and the tetravalence of carbon. Transition metals absorb hydrogen to a widely variable degree, forming interstitial compounds of stoichiomet~y depending on conditions and having properties sometimes resembling but generally unlike those of the major group hydrogen compounds. For example, in scandium, yttrium, and some of the inner transition elements including plutonium, MHz phases appear to occur in which, however, the composition tends to be more nearly MH,.,. Additional hydrogen can be dissolved in these phases. An MHa phase occurs in a few. The MHz phases are hydrolyzed by water and acids; the hydride of Ti, however, appears unaffected by water. In general the most tightly associated of the transition metals (those having d orbitals nearly half full) appear least reactive toward hydrogen. Metal-hydrogen bonds are found in a number of complex molecules or ions. Metals that may be so involved include Ti, V, Cr, Mn, Mn, Fe, Co, Ni, Mo, Tc, Ru, Rh, Pd, Ta, W, Re, Os, Ir, Pt. Although these complexes might be regarded as examples of hydride ion acting as ligand, the hydridic character if any is not necessarily evident. For example, although such complexes appear to be quite susceptible to oxidation, they resist hydrolysis and some tend to act as acids in water solution. Others, however, are basic. The nature of these complexes remains to be clarified. In summary, the chemistry of hydrogen depends largely on the fundamental nature of hydrogen atoms. Application of the principle of electronegativity equalization to the estimation of the condition of combined hydrogen with respect to charge leads to recognitionlof a consistent pattern of behavior of hydrogen compounds, from which useful principles such as those listed here may be derived. In this sense it is possible for teachers to present the chemistry of hydrogen as a self-consistent, logical unit upon which a practical understanding can be based.
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