descriptive inorganic chemistry || hydrogen

CHAPTER 6

Hydrogen

Although hydrogen is the simplest atom, the chemistry of hydrogen is extensive. Part of thereason for this is because the hydrogen atom resembles the members of two groups ofelements. Because it has the electron configuration of 1s1, it resembles to some extent thealkali metals in Group IA, which have configurations of ns1. However, the ionization energyof hydrogen is about 1314 kJ mol−1 and the strength of the H–H bond in the diatomicmolecule is about 435 kJ mol−1. Consequently, the formation of a simple H+ speciesrequires considerable energy. In fact, so much energy is required for the formation of H+

that compounds in which the single proton is present as a cation are not likely to form. Incontrast, the alkali metals have ionization energies varying from about 377 kJ mol−1 for Csto about 519 kJ mol−1 for Li so that many compounds of these elements are essentiallyionic and contain the singly charged ions. There are, however, some solid compounds thatcontain ions such as H5O2

+ or H9O4+. These ions are solvated protons that contain two and

four water molecules of hydration, respectively. The reason for the existence of ions of thistype is the high heat of hydration of H+ (−1100 kJ mol−1), which results from its small sizeand the resulting high charge to size ratio. The heat of hydration of H3O

+ is estimated to beabout −390 kJ mol−1.

Hydrogen was discovered in 1781 by Cavendish who prepared a gas that produces waterduring its combustion. Shortly thereafter, the name hydrogen was given to the gas. Mostsimple hydrogen compounds are covalent as a result of the hydrogen atom sharing anelectron pair. Because the 1s level is singly occupied in the hydrogen atom, it alsoresembles in many ways the halogens, which also require a single electron to completethe valence shell. Accordingly, hydrogen forms a substantial number of compounds inwhich it gains an electron to form a hydride ion. Thus, although hydrogen is the simplestatom, the chemistry of hydrogen is indeed varied and encompasses many types of compoundsand reactions.

6.1 Elemental and Positive Hydrogen

As a result of the nucleus of the hydrogen atom being a proton that has a spin quantumnumber of ½, a hydrogen molecule may have the spins both being aligned or opposed. Theresult is that there are two forms of elemental hydrogen. These are known as ortho H2 if the

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spins are aligned or para hydrogen if the spins are opposed. At room temperature, themixture is composed of 75% ortho and 25% para hydrogen.

The absence of compounds that contain the simple H+ ion has been discussed here. However,the solvated species are not the only species known that contain H+ in some form. A seriesof cations have been identified in mass spectrometry that result from the attachment of ahydrogen molecule to H+. The simplest of these is the H3

+ ion that has a trigonal planarstructure. The electrostatic interaction of a proton with H2 occurs at the shared pair ofelectrons that can be shown as follows:

H�

H : H

After H+ attaches, the electrons (and there are only two) are shared equally in an equilateraltriangular arrangement. Of course, multiple resonance structures are possible, which can beshown as

HH

H�

H� H

H

H�H

H

The H3+ ion represents the simplest example of a two-electron three-center bond in which a

molecular orbital containing two electrons encompasses all three of the atoms. Instead of theapproach described earlier, a more satisfactory description of the bonding is provided byconstructing a molecular orbital from a combination of three hydrogen wave functions,

ψMO ¼ 1ffiffiffi

3p ðfa þ fb þ fcÞ ð6:1Þ

where ψMO is the wave function for the molecular orbital and fa, fb, and fc representthe 1s atomic wave functions for hydrogen atoms a, b, and c, respectively. The resultingmolecular orbital diagram can be represented as shown in Figure 6.1.

In this case, it is assumed that the molecular orbitals in a hydrogen molecule can berepresented in the usual way (see Section 2.3.1) and that the bonding molecular orbital isthen used with the third atomic orbital to give the bonding molecular orbital for the entirespecies. Because H3

+ contains only two electrons, the bonding molecular orbital is occupiedby two electrons, but it encompasses all three atomic centers. This type of three-centerbonding will also be discussed in later chapters.

The series of species containing a proton attached to hydrogen molecules can be representedby the formula Hn

+ with H3+ being the most stable, and others having n being an odd number

are significantly more stable than those in which n is an even number. These species have

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structures that are shown in Figure 6.2. The ions arise from the addition of H2 molecules atthe corners of the triangle formed in H3

+.

Elemental hydrogen can be prepared conveniently by either the reduction of positivehydrogen, as typified by the reactions of metals with mineral acids, or the oxidation of thehydride ion. The first process may be represented as

2 Mþ 2 HA→H2 þ 2 MA ð6:2Þwhere M is a metal above hydrogen in the electromotive series and A = Cl−, Br−, ½SO4

2−,and so on. The oxidation of a hydride ion by positive hydrogen is the simplest reactionyielding hydrogen from hydrides. The following are typical reactions of this type, but theyalso represent proton transfer (acid-base) processes:

H2Oþ H� →H2 þ OH� ð6:3ÞCH3OHþ H� →H2 þ CH3O

� ð6:4ÞElemental hydrogen may also be prepared by the electrolysis of water:

2 H2O →electricity

2 H2 þ O2 ð6:5Þ

H

HH

H

HH

H H

H H

H

H

H

H

H

H

H

HH

HH

H

HH

H3� H5

� H7� H9

�

Figure 6.2Structures of some Hn

+ ions.

�a �b �c

Ψ *

ψ*

Ψ

ψ

Figure 6.1A molecular orbital diagram for H3

+ in which fa, fb, and fc represent 1s atomic wave functions,ψ and ψ* are wave functions for bonding and antibonding states in H2,

and Ψ and Ψ* are wave functions for the H3+ ion.

Hydrogen 155



Because the reduction of hydrogen at the cathode can be essentially an isolated process, thismethod is useful for the production of high purity H2. Hydrogen is also liberated by theaction of NaOH on amphoteric metals such as aluminum and zinc:

2 Alþ 6 NaOH→ 2 Na3AlO3 þ 3 H2 ð6:6ÞIn aqueous solution, this reaction can be represented more accurately by the equation

2 Alþ 2 NaOHþ 6 H2O→ 2 NaAlðOHÞ4 þ 3 H2 ð6:7ÞTo be commercially feasible, major industrial preparations of hydrogen must use inexpensivematerials and processes. A reducing agent that is as expensive as a metal or electricitysimply will not give an economically feasible process for preparing hydrogen. At very hightemperatures, carbon in the form of coke (obtained by heating coal) reacts with water toproduce H2:

Cþ H2O→COþ H2 ð6:8ÞThe mixture of CO and H2 is known as water gas. It is a good reducing agent (both COand H2 are easily oxidized), and it has been used as a fuel because both gases will readilyburn. Large quantities of hydrogen are also produced by treatment of petroleum products byeither catalytic reforming or dehydrogenation. In catalytic reforming, hydrocarbons havehydrogen removed, and structural changes are produced. For example, hexane is convertedto cyclohexane:

C6H14 →catalyst

C6H12 þ H2 ð6:9ÞThe cyclohexane then can be converted to benzene:

C6H12 →catalyst

C6H6 þ 3 H2 ð6:10ÞThe catalysts used in these types of processes are usually platinum in some form.Dehydrogenation involves removal of hydrogen, and two important processes are conversionof butane to butadiene,

C4H10 →catalyst

CH2¼CH�CH¼CH2 þ 2 H2 ð6:11Þand the conversion of ethyl benzene to styrene,

C6H5CH2CH3 →catalyst

C6H5CH¼CH2 þH2 ð6:12ÞBoth butadiene and styrene are used in very large quantities in the preparation ofpolymers, so these dehydrogenation processes result in the production of large quantitiesof hydrogen.

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One of the most important processes for the production of hydrogen is the steamreformer process. In that process, a hydrocarbon and steam are passed over a nickelcatalyst at 900 °C. The reactions that take place when the hydrocarbon is methane areshown here:

CH4 þ 2 H2O→CO2 þ 4 H2 ð6:13Þ

CH4 þ H2O→COþ 3 H2 ð6:14ÞThe mixture of gases that results contains H2, CO, and CO2 along with some steam. Thismixture is then passed into a shift converter and at 450 °C, the CO is converted into CO2:

COþ H2O→CO2 þ H2 ð6:15ÞIn years past, another process for making hydrogen was important. The basis for thatprocess was the following reaction:

3 Feþ 4 H2O →700 oC

Fe3O4 þ 4 H2 ð6:16Þ

When all the iron becomes oxidized, it can be regenerated to Fe by the following reactionscarried out at red heat:

Fe3O4 þ 4 H2 → 3 Feþ 4 H2O ð6:17ÞFe3O4 þ 4 CO→ 3 Feþ 4 CO2 ð6:18Þ

Finally, a preparation of hydrogen that is of considerable importance is the electrolysis ofaqueous NaCl solutions:

2 Naþ þ 2 Cl� þ 2 H2O →electricity

Cl2 þ H2 þ 2 Naþ þ 2 OH� ð6:19ÞThis process is of enormous importance because it represents the commercial preparation ofchlorine and sodium hydroxide. Because it is carried out on a large scale, this process alsoproduces large quantities of hydrogen.

Hydrogen is used as the fuel in the oxy-hydrogen torch to produce an extremely hotflame. It is also used in bubble chambers for tracking elementary particles and as arocket fuel. Hydrogen has several attributes that make it an attractive energy source.First, a great deal of energy is released when hydrogen burns. Second, the combustionproduct is water, so there are no environmental issues. Third, the major source ofhydrogen is water, which is abundantly available. However, the difficulty is how toseparate the H2O molecules to obtain large quantities of hydrogen economically. Thereare also difficulties associated with fabricating engines that run for extended periodsbecause of the reaction of hydrogen with metals (see Section 6.3.2) and with storage of

Hydrogen 157



hydrogen in the liquid state. Thus, an economy based on hydrogen as a major fuel maynot be forthcoming soon.

6.2 Occurrence and Properties

Elemental hydrogen does not occur in the earth’s atmosphere to any significant extentowing to the low molecular mass of the molecules. However, it does occur to a large extentin other parts of the universe. Combined hydrogen is present on the earth’s surface in awide variety of compounds, especially water. The principal use of elemental hydrogen is inseveral important hydrogenation reactions, notably in the production of ammonia, methylalcohol, and a large number of organic materials and foodstuffs.

Ordinary hydrogen consists of a mixture of three isotopes. The isotope with mass number 1,H, is about 6400 times as abundant as deuterium, D, the isotope with mass number 2.Tritium, T, the isotope with mass number 3, is many orders of magnitude less abundantthan deuterium. Although six different diatomic molecules are possible from these threeisotopes, H2, D2, HD, and T2 have been more thoroughly studied. Of these, the two mostcommon forms are H2 and D2. Some of the properties of these two forms of hydrogen arelisted in Table 6.1.

Although D2 undergoes almost all the same reactions as H2, the rates of these reactions aregenerally lower. In some cases, there is a large difference between the rate of the reactionwhen deuterium is involved and the rate when hydrogen is involved. The rate is influencedby the kinetic isotope effect, which refers to the fact that the atomic mass of deuteriumbeing twice that of hydrogen causes the rates of most reactions involving deuterium to belower. It is only in the case of hydrogen and deuterium that such a large relative massdifference occurs. The effect arises from the fact that whereas X–H and X–D bonds areabout the same strength, the vibrational frequencies are much different with that for X–Hbeing much higher. Therefore, the X–H bond reacts more rapidly.

By far the most frequently encountered deuterium compound is D2O or “heavy” water.This material is generally obtained by enrichment from natural water by electrolysis.The normal water, H2O, is preferentially electrolyzed, leaving water enriched in D2O.Eventually, almost pure D2O can be obtained. Some of the properties of D2O and H2Oare shown in Table 6.2.

Table 6.1: Physical Properties of H2 and D2

Property H2 D2

Boiling point, K 20.28 23.59Triple point, K 13.92 18.71

Heat of fusion, J mol−1 117 219

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One of the important features of D2O is its ability to undergo isotope exchange with manycompounds. For example, when NH3 is placed in D2O, all of the H atoms exchangerapidly with deuterium. Presumably, this is because the exchange with NH3 can proceedby interaction of D2O with the unshared pair of electrons on the nitrogen atom in anassociative process that can be illustrated as

H3N: ⋯D�OD

The exchange with NH4+ is much slower because attachment of deuterium depends on

a dissociative mechanism of NH4+. The exchange of D with H in [Co(NH3)6]

3+ in D2Ois very slow for the same reason. Likewise, the exchange of hydrogen by formic acid,HCOOH, in D2O proceeds rapidly for the hydrogen attached to oxygen, but very slowlyfor the hydrogen attached to carbon. For phosphorous acid, H3PO3, two hydrogen atomsexchange rapidly, but the third does not. The reason for this is that in this molecule,one hydrogen atom is bonded directly to the phosphorus atom:

P

O

H

OHHO

To undergo facile exchange between D and H, the bond must have some polarity. Becausethe electronegativities of P and H are almost exactly the same, the hydrogen attached to thephosphorus atom does not exchange readily. As we observed earlier, even if the bond ispolar (as in NH3), there must be an unshared electron pair to provide a site of attack or atleast there must be some low-energy pathway for the reaction.

The major use of D2O is as a moderator for neutrons in nuclear reactors. Because thedeuterium atom is small, neutrons colliding with it cause the deuterium atom to recoil sothat more energy is absorbed than would be if the atom did not recoil. Thus, “heavy” water

Table 6.2: Properties of H2O and D2O

Property H2O D2O

Melting point, °C 0.00 3.8Boiling point, °C 100.00 101.42

Temperature of max. density, °C 3.96 11.6Heat of vaporization, kJ mol−1 40.66 41.67

Heat of fusion, kJ mol−1 6.004 6.368Dielectric constant at 20 °C 82 80.5

Solubility of KCl at 0 °C, molesper 100 moles of solvent

6.81 5.69

Hydrogen 159



is used for this purpose. Most of the other compounds of interest that contain positivehydrogen will be discussed in chapters on the chemistry of the other elements. We will nowdiscuss the chemistry of negative hydrogen.

6.3 Hydrides

In most compounds, hydrogen exhibits the oxidation state of +1. Because of its relativelyhigh electronegativity, there are compounds in which hydrogen assumes a negative oxidationstate either by gaining an electron in the 1s state or by sharing electrons with an element oflower electronegativity. Formally, this is true in any binary compound in which the otherelement has a lower electronegativity than hydrogen. Hydrides are conveniently grouped intothree classes, depending on the mode of binding to the other element. However, as in the caseof other chemical bonds, the bond types are not completely separable because there is agradual transition from ionic to covalent depending on the relative electronegativities of theatoms. As a result, the three classifications of hydrides described next are somewhat artificial.

6.3.1 Ionic Hydrides

It is logical to expect that elements of low electronegativity might produce hydride ions bygiving up electrons to hydrogen. Hydrogen has an electron affinity of −74.5 kJ mol−1, butthis is not sufficient to overcome the ionization energy of the metal if only the pair of ionswere formed. The process becomes energetically favorable when a solid crystal is formed,and solid hydrides are known for the Group IA and IIA elements of the periodic table.Generally, these compounds can be formed by direct union of the elements:

2 Mþ H2 → 2 MH ð6:20ÞThe reaction is carried out at about 400 °C for sodium, potassium, and rubidium and atabout 700 °C with lithium. Metal compounds such as the nitrides may also be used:

Na3Nþ 3 H2 → 3 NaHþ NH3 ð6:21ÞThe ionic hydrides are white solids with high melting points, and all of the alkali metalhydrides have the sodium chloride crystal structure. Because they resemble the salts of thealkali and alkaline earth metals, the ionic hydrides are often referred to as saline or salt-likehydrides. The properties of the alkali metal hydrides are shown in Table 6.3, and those ofthe alkaline earth hydrides are shown in Table 6.4.

The apparent radius of the H− is determined from the M–H distance in the crystal by subtractingthe known radius of M+. It is obvious from the data shown in Table 6.3 that the radiusassigned to H− in Li–H is smaller than in other hydrides of Group IA. The Li–H bond isconsidered to have a substantial amount of covalent character. Probably the greater covalency

160 Chapter 6



in this compound compared to the other alkali metal hydrides is a result of the higherionization potential of Li and the fact that the 1s orbital of hydrogen and the 2s orbital oflithium are of similar size. The vast difference in size between 1s of H and the 3s or 4s of Naor K decreases the effectiveness of overlap. Also, the fact that Na and K have lowerionization potentials than Li tends to make NaH and KH more ionic than LiH. The H− ion is alarge, soft electronic species that is easily polarizable, so electron density would be drawntoward a smaller ion such as Li+ to a greater extent than with other ions of Group IA metals.

Molten hydrides of Groups IA and IIA are good electrical conductors, and hydrogen isliberated at the anode as a result of the oxidation of H−:

2 H� →H2 þ 2 e� ð6:22ÞThe alkali metal or alkaline earth metal is reduced at the cathode.

The dominant feature of the chemistry of the ionic hydrides is the strongly basic characterof the H− ion. All the ionic hydrides react readily with protonic solvents to producehydrogen gas and a base that is weaker than H−:

H� þ H2O→OH� þ H2 ð6:23ÞH� þ ROH→RO� þ H2 ð6:24ÞH� þ NH3 →NH2

� þ H2 ð6:25ÞBecause the hydride ion has a pair of electrons, it can function as a Lewis base as well. Forexample, many coordination compounds are known in which the hydride ion is present as a

Table 6.4: Some Properties of Alkaline Earth Hydrides

CompoundΔHf

o

(kJ mol−1)Apparent Charge on

H (e units)Density, Hydride

(g cm−3)Density, Metal

(g cm−3)

CaH2 −195 −0.27 1.90 1.55SrH2 −177 −0.31 3.27 2.60BaH2 −172 −0.36 4.15 3.59

Table 6.3: Selected Properties of the Group I Hydrides

CompoundΔHf

o

(kJ mol−1)U

(kJ mol−1)H− Radius

(pm)

ApparentCharge on H(e units)

Density, Hydride(g cm−3)

Density, Metal(g cm−3)

LiH −89.1 916 136 −0.49 0.77 0.534NaH −59.6 808 146 −0.50 1.36 0.972KH −57.7 720 152 −0.60 1.43 0.859RbH −47.7 678 153 −0.63 2.59 1.525CsH −42.6 644 154 −0.65 3.41 1.903

Hydrogen 161



ligand. Some of the simplest of these are the tetrahydridoaluminate(III), AlH4−, and the

tetrahydridoborate(III) ions. Although these species are largely covalent, they can beconsidered as arising from the coordination of four H− ions to Al3+ and B3+ ions. The saltsof these complex ions, LiAlH4 and NaBH4, are widely used as hydrogenating (reducing)agents in organic synthesis reactions. They may also be used for preparing other hydridesby reactions such as

MR2 þ 2 LiAlH4 →MH2 þ 2 LiAlH3R ð6:26Þwhere M = Zn, Cd, Be, or Mg and R = CH3 or C2H5.

Because of its large size, the electron cloud of H− is easily polarizable, and therefore, H− isa soft base. Consequently, complexes with transition metals are usually formed in which themetals are soft Lewis acids. The metals are usually in low oxidation states, which makesthem softer acids than the same metals in higher oxidation states. Typical among thesecompounds are Fe(CO)4H2, Re(CO)(P(C6H5)3)3H, and Mn(CO)5H.

6.3.2 Interstitial Hydrides

Unlike the ionic hydrides, interstitial hydrides may have positive heats of formation, andthey always have densities that are lower than the parent metal. Interstitial hydrides aregenerally described as the larger spherical atoms of the metal forced slightly apart with asmaller hydrogen atom occupying (interstitial) holes between them. Metallic palladium canabsorb up to 700 times its own volume of hydrogen gas. Frequently, these compounds arecharacterized by formulas such as CuH0.96, LaH2.78, TiH1.21, TiH1.7, or PdH0.62.

In the formation of interstitial hydrides, the metal lattice is expanded to accommodate thehydrogen atoms, but the metal atoms retain their original crystal structure. There is, however,a significant change in physical and chemical properties of the metal. These properties arepredictable in a very straightforward manner.

In a metal, certainly the transition metals, the electrons are more or less free to move inconduction bands. This fact is responsible for the high electrical conductivity of metals.When hydrogen atoms are present in the holes between the atoms, the movement of theelectrons is somewhat impaired. As a result, the metal hydrides of this class are poorerconductors than the pure metals. The presence of hydrogen atoms makes the metal atomsless mobile and more restricted to particular lattice sites. Accordingly, the interstitial metalhydrides are more brittle than the parent metal. Also, the inclusion of the hydrogen atomscauses a small degree of lattice expansion so that the interstitial hydrides are less dense thanthe parent metal alone.

The nature of the process of forming the interstitial metal hydrides explains why some ofthese compounds are accompanied by a positive heat of formation. The bond energy in H2

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is about 432.6 kJ mol−1, and the metal lattice must be expanded to allow hydrogen atomsto enter the lattice. The interaction between the hydrogen atoms and metal atoms is notexothermic enough to compensate for the large energy requirement. The transition metalshave electronegativities that are higher than the alkali or alkaline earth metals (generallyin the range 1.4 to 1.7), which makes it unlikely that the hydrides of these metals wouldbe ionic. Although the transition metals, lanthanides, and actinides have holes in theirlattices that can accommodate hydrogen atoms, these hydrides probably should also not beconsidered as covalent. In some ways, the interstitial hydrides are probably best consideredas solutions of atomic hydrogen in the metals. Because these metals interact with molecularhydrogen causing it to be separated into atoms, these metals can function as effectivehydrogenation catalysts. These metals essentially function to split the hydrogen moleculesinto atoms, which, for the most part, remain in the metal lattice. Some escape of hydrogendoes occur, and the hydrogen is evolved as highly reactive atomic hydrogen at the surfaceof the metal.

The prior treatment that the metal surface has received greatly affects the ease of hydrideformation because the hydrogen must first be adsorbed on the metal surface before dissolutionoccurs. As a result, it is possible that not all of the available interstitial positions will becomeoccupied by hydrogen, resulting in compositions that are variable depending on the temperatureand pressure used in preparing the metal hydride. As a result, the composition may not beexactly stoichiometric, and hydrides of this type are sometimes referred to as nonstoichiometrichydrides.

6.3.3 Covalent Hydrides

Strictly speaking, a binary compound of hydrogen is not a hydride unless hydrogen hasthe higher electronegativity of the two elements. Obviously, this is the case in the metalcompounds such as the ionic and interstitial hydrides just discussed. There are, however,some covalent compounds in which this is also the case depending on the electronegativityof the other element. To illustrate this point, Table 6.5 shows the electronegativities of somemain group elements.

Table 6.5: Electronegativities of Some Main Group Elements

Element Electronegativity Element Electronegativity

H 2.1 P 2.1B 2.0 As 2.0C 2.5 Sb 1.8Si 1.8 F 4.0Ge 1.8 Cl 3.0N 3.0 Br 2.8

Hydrogen 163



On the basis of electronegativities, we would say that SbH3 is a hydride, whereas NH3 isnot. Although any binary compound of hydrogen is sometimes referred to as a “hydride,”we will not use that classification here. Thus, H2O is an oxide of hydrogen rather than ahydride of oxygen. Similarly, HF, HCl, HBr, and HI are considered to be hydrogen halides,not halogen hydrides. Typically, the binary compounds of B, Si, Ge, Sn, P, As, and Sb withhydrogen are considered to be hydrides.

Beryllium and magnesium hydrides, BeH2 and MgH2, appear to be polymeric covalenthydrides rather than ionic hydrides as are those formed by the other Group II metals.The structure of these compounds consists of chains in which the H atoms form bridgesas illustrated here for BeH2:

Be Be Be

H H HH

H H HH

The four H atoms around each Be are arranged in an approximately tetrahedral manner.

When magnesium is heated with boron, the product is magnesium boride, MgB2 (sometimeswritten as Mg3B2). The hydrolysis of the compound produces Mg(OH)2 and a boronhydride. Instead of the expected product of BH3, borane, the product is B2H6, diborane.This compound has the structure

B

H

H

B

H H

H H

in which the boron atoms are located in the centers of somewhat distorted tetrahedra ofhydrogen atoms (see Chapter 8). Diborane can also be prepared by the reaction of BF3 withNaBH4:

3 NaBH4 þ BF3 → 3 NaFþ 2 B2H6 ð6:27Þ

Many other boron hydrides are known, and the chemistry of these interesting compoundswill be discussed in Chapter 8.

Because most of the covalent hydrides consist of molecular units with only weakintermolecular forces between them, they are volatile compounds. Accordingly, thecovalent hydrides are sometimes referred to as the volatile hydrides. The nomenclature,melting points, and boiling points of several covalent hydrides are shown in Table 6.6.

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Procedures similar to those described for the preparation of diborane can be employed forthe preparation of the hydrides of silicon and germanium. For example, Mg2Si results whenMg and Si react. This compound reacts with HCl to produce silane, SiH4:

Mg2Siþ 4 HCl→ 2 MgCl2 þ SiH4 ð6:28ÞLithium aluminum hydride reacts with SiCl4 as follows to form silane:

SiCl4 þ LiAlH4 →SiH4 þ LiClþ AlCl3 ð6:29ÞThe reaction is not actually this simple, and a mixture of silanes is produced.

Silanes are more reactive than the corresponding carbon compounds, and silane is flammablein air. Although ethane is unreactive toward water, silanes such as Si2H6 undergo hydrolysisin basic solution:

Si2H6 þ 2 H2Oþ 4 NaOH→ 2 Na2SiO3 þ 7 H2 ð6:30ÞAcid hydrolysis of a metal phosphide produces phosphine:

2 AlPþ 3 H2SO4 → 2 PH3 þ Al2ðSO4Þ3 ð6:31Þ

The action of sodium hydroxide on elemental white phosphorus also produces phosphine:

P4 þ 3 NaOHþ 3 H2O→PH3 þ 3 NaH2PO2 ð6:32Þ

Table 6.6: Properties of Some Covalent Hydrides

Name Formula m.p.,°C b.p.,°C

Diborane B2H6 −165.5 −92.5Tetraborane B4H10 −120 18

Pentaborane-9 B5H9 −46.6 48Pentaborane-11 B5H11 −123 63Hexaborane B6H10 −65 110Ennaborane B9H15 2.6 —

Decaborane B10H14 99.7 213Silane SiH4 −185 −119.9Disilane Si2H6 −132.5 −14.5Trisilane Si3H8 −117 54Germane GeH4 −165 −90Digermane Ge2H6 −109 29Trigermane Ge3H8 −106 110Phosphine PH3 −133 −87.7Diphosphine P2H4 −99 −51.7

Arsine AsH3 −116.3 −62.4Stibine SbH3 −88 −18

Hydrogen 165



This reaction also produces small amounts of diphosphine, P2H4, which is spontaneouslyflammable in air. Phosphine also burns readily in air,

4 PH3 þ 8 O2 → 6 H2Oþ P4O10 ð6:33ÞAlthough PH3 is somewhat similar to ammonia, it is a much weaker base toward protons inaccord with the hard-soft interaction principle (see Chapter 5). Phosphonium salts areusually stable only when the anion is also large (e.g., PH4I) and the parent acid is a strongone (e.g., HBr or HI). However, PH3 and substituted phosphines are very good ligandstoward second- and third-row transition metal ions. Many coordination compounds containingsuch ligands are known, and some of them have important catalytic properties.

The bonding in PH3 apparently involves much more nearly s-p overlap than is present inNH3. In NH3, the H–N–H bond angles are about 107°, whereas in PH3 the H–P–H anglesare about 93°. Although increasing the extent of hybridization reduces repulsion by givingbond angles closer to the tetrahedral angle, it decreases the effectiveness of the overlap ofthe small hydrogen 1s orbital with the sp3 orbitals on the central atom. The bond angles inAsH3 and SbH3 are even slightly smaller than those in PH3. A similar trend is seen forthe series H2O, H2S, H2Se, and H2Te. In the hydrogen compounds of the heavier membersof each group, there is little tendency for the central atom to form sp3 hybrids. Thus, theH–X–H bond angles indicate that the central atom uses essentially pure p orbitals in bondingto the hydrogen atoms, and the bond angles in these compounds are approximately 90°.

The discussion of the covalent hydrides given here is somewhat brief in keeping with theintended purpose of this book. Some of the types of hydrides will be discussed in greaterdetail in chapters dealing with the other elements. The discussion presented here shouldserve to introduce the breadth, scope, and importance of hydride chemistry.

References for Further Reading

Bailar, J. C., Emeleus, H. J., Nyholm, R., & Trotman-Dickinson, A. F. (1973). Comprehensive Inorganic Chemistry(Vol. 1). Oxford: Pergamon Press. This five-volume set is a standard reference work in inorganic chemistry.Hydrogen chemistry is covered in Vol. 1.

Cotton, F. A., Wilkinson, G., Murillo, C. A., & Bochmann, M. (1999). Advanced Inorganic Chemistry (6th ed.).New York: John Wiley. A 1300-page book that covers an incredible amount of inorganic chemistry.Chapter 2 deals with hydrogen. Highly recommended.

Greenwood, N. N., & Earnshaw, A. (1997). Chemistry of the Elements (2nd ed., Chap. 3). Oxford: Butterworth-Heinemann. A comprehensive general reference work.

King, R. B. (1995). Inorganic Chemistry of the Main Group Elements. New York: VCH Publishers. An excellentintroduction to the descriptive chemistry of many elements. Chapter 1 deals with the chemistry ofhydrogen.

Mueller, W. M., Blackledge, J. P., & Libowitz, G. G. (1968). Metal Hydrides. New York: Academic Press. Anadvanced treatise on metal hydride chemistry and engineering.

Muetteries, E. F. (Ed.) (1971). Transition Metal Hydrides. New York: Marcel Dekker. A collection of chaptersby noted researchers in the field. Covers stereochemistry, catalysis, and reactions.

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Problems

1. Explain why there are more covalent than ionic binary compounds of hydrogen.

2. Write balanced equations for the following processes.(a) Dehydrogenation of pentane(b) The reaction of carbon with high temperature steam(c) The reaction of sodium hydride with water(d) The electrolysis of water(e) The reaction of zinc with aqueous HCl

3. Complete and balance the following.(a) Li + H2 →

(b) Ca3N2 + H2 →

(c) CaH2 + CH3OH →

(d) Zn + NaOH + H2O →

(e) Cd(C2H5) + LiAlH4 →

4. Describe the steam reformer process for producing hydrogen.

5. Describe some of the major difficulties associated with a hydrogen gas–based energyprogram.

6. Write the equations to describe the preparation of the following compounds. Giveconditions where possible.(a) H2 (commercial production)(b) Diborane starting with magnesium metal and boron(c) NaOH (commercial production)(d) Styrene, C6H5CH=CH2

7. Given the stability of the water molecule, speculate on the type of processes that couldbe used to produce the large quantities of hydrogen necessary for the gas to serve as themajor energy source. Keep in mind that the processes must produce the gas cheaplyenough for it to be economically competitive with other energy sources.

8. Complete and balance the following.(a) LiAlH4 + SnCl4 →(b) Ca + H2O →

(c) AlCl3 + LiH →

(d) CaH2 + H2O →

(e) Al + NaOH + H2O →

9. Using the hard-soft interaction principle, explain why the apparent ionic radius of H− inLiH is smaller than it is in KH.

Hydrogen 167



10. Which of the following reactions would take place more rapidly if all other conditionsare the same? Provide an explanation for your answer.

CH3COOHþ C2H5OH→CH3COOC2H5 þ H2O

CH3COODþ C2H5OD→CH3COOC2H5 þ D2O

11. How does the electrolysis of water lead to the production of D2O?

12. Explain why acetylene, C2H2, behaves as an acid and reacts with active metals toliberate hydrogen.

13. Explain why a linear structure for H3+ would be less stable than a triangular one.

14. From the standpoint of structure and bonding, explain why transition metals such aspalladium and nickel are often used as catalysts for hydrogenation reactions.

15. Explain the following observations.(a) Interstitial hydrides are less dense than the parent metal.(b) Interstitial hydrides have variable composition.(c) Interstitial hydrides are poorer conductors of electricity than the parent metal.(d) Ionic hydrides are more dense than the parent metal.

168 Chapter 6



descriptive inorganic chemistry || hydrogen

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