Чамберс К., Холлидей А.К. Современная неорганическая химия, 1975

176 GROUP IV

in silicon, 226kJmol~1 ) and catenation (the phenomenon of selflinkage between atoms of the same element) is consequently less marked with silicon than with carbon; the higher silanes decompose slowly even at room temperature. Silanes are far more sensitive to oxygen than alkanes and all the silanes are spontaneously inflammable in air, for example

SiH4 + 2O2 -> SiO2 + 2H2O

This greater reactivityof the silanes may be due to several factors, for example, the easier approach of an oxygen molecule (which may attach initially to the silane by use of the vacant silicon d orbitals) and the formation of strong Si—O bonds (stronger than C—O).

Halogen derivativesof silanes can be obtained but direct halogenation often occurs with explosive violence; the halogen derivatives are usually prepared by reacting the silane at low temperature with a carbon compound such as tetrachloromethane, in the presence of the corresponding aluminium halide which acts as a catalyst.

Silanes are very sensitive to attack by alkalis and will even react with water made alkaline by contact with glass; this reaction is in marked contrast to the reactions shown by alkanes. Unlike alkanes, silanes are found to have marked reducing properties and will reduce, for example, potassium manganate(VII) to manganese(IV) oxide, and iron(III) to iron(II).

In addition to the volatile silanes, silicon also forms non-volatile hydrides with formulae(SiH2)x but little is known about their structure. Silicon, however, does not form unsaturated hydrides corresponding to the simple alkenes.

Germanium

Germanium forms a series of hydridesof general formula GenH2n+2 which are quite similar to the corresponding silanes. Only a small number of germanes have so far been prepared. Germanes are not as inflammable as the corresponding silanes (the Ge—O bond is not as strong as the Si—O bond) and they are also less reactive towards alkalis, monogermane being resistant to quite concentrated alkali.

Tin

The greater metallic nature of tin is clearly indicated here for tin forms only one hydride, stannane, SnH4. It is best prepared by the

178 GROUP IV

(p. 398) reacts with carbon monoxide in preference to oxygen so preventing the haemoglobin from acting in its normal capacity as an oxygen carrier.

Carbon monoxide is formed by the incomplete combustion of carbon. It is prepared in the laboratory by dropping methanoic (formic) acid into warm concentrated sulphuric acid; the latter dehydrates the methanoic acid:

HCOOHFil^4COt + H2O

The gas is passed through caustic soda solution to remove any sulphur dioxide or carbon dioxide produced in side reactions. Carbon monoxide is also obtained when an ethanedioate (oxalate) is heated with concentrated sulphuric acid:

c2oj- + H2so4 -> cot + co2| + H2o + sor

The carbon dioxide is removed by passage of the gas through a mixture of sodium and calcium hydroxides. Very pure carbon monoxide is produced by heating nickel tetracarbonyl (see p. 179):

Ni(CO)4 -> Ni + 4COt

The commercial production of carbon monoxide in the form of water gas is now largely obsolete. The production by the reaction between steam and hydrocarbons is considered later (p. 180).

The structure of carbon monoxide can be represented as a resonance hybrid between two structures

C=O and C=O

In structure (a) each atom has a complete octet; in the actual molecule, the carbon-oxygen bond length is greater than would be expected for a triple bond, and the molecule has a much smaller dipole moment than would be expected if the oxygen was donating electrons to the carbon as in (a); hence structure (b) must contribute to the actual structure. A simplified orbital picture of structure is shown at the top of the next page, where nl is formed by sharing electrons from both carbon and oxygen and n2 is formed byelectrons donated from oxygen only.

This structure indicates that carbon monoxide should have donor properties, the carbon atom having a lone pair of electrons. Carbon

GROUP IV 179

Full sp carbon

hybrid orbital \ \ ^ ^ \ / ^ ^ Singly-filled p

orbitaIs on oxygen

Singlyfil led p orbital on carbon

monoxide is in fact found to have donor properties and forms donor compounds, for example with diborane (p. 145)it splits the molecule by donating to the borane, BH3.

2CO + B2H6 ^ 2OC-»BH3

It also forms compounds known as carbonyls with many metals. The best known is nickel tetracarbonyl, Ni(CO)4, a volatile liquid, clearly covalent. Here, donation of two electrons by each carbon atom brings the nickel valencyshell up to that ofkrypton (28+ 4 x 2 ) ; the structure may be written Ni( <- C=O)4. (The actual structure is more accurately represented as a resonance hybrid of Ni( <- C=O)4 and Ni(=C=O)4 with the valency shell of nickel further expanded.) Nickel tetracarbonyl has a tetrahedral configuration,

O

C

C

O

Other examples are iron pentacarbonyl, Fe(CO)5, and chromium hexacarbonyl, Cr(CO)6, which have trigonal bipyramidal and octahedral configurations respectively.

Carbon monoxide burns with a characteristic blue flame in air or oxygen. The reaction

2CO + O2 -> 2CO2 :AH=- 283 kJ mol"i

is very exothermic and as expected, therefore, carbon monoxide reacts with heated oxides of a number of metals, for example copper, lead, iron, reducing them to the metal. For example :

PbO + CO -> Pb + CO2|

Carbon monoxide forms addition compounds. With chlorine in sunlight or in the presence of charcoal in the dark,carbonyl chloride

180 GROUP IV

(phosgene). COCl2. is formed :

CO + C12 -> COCK

^Vith ammoniacal or hydrochloric acid solution of copper(I) chloride, carbon monoxide formsthe addition compound CuCl .CO. 2H2O. This reaction can be used to quantitatively remove carbon monoxide from gaseous mixtures.

Although carbon monoxide appears to be the anhydride of methanoic acid it does not react with water to give the acid; however, it will react with sodium hydroxide solution above 450 K, under pressure, to give sodium methanoate:

CO + NaOH -> HCOO~Na+

Carbon dioxide, CO2. Carbon dioxide is present in air and escapes from fissures in the earth in volcanic regions and where ^mineral springs' occur. It may be prepared by:

(1) the action of dilute acid on any metal carbonate or hydrogencarbonate, for example

CaCO3 + 2HC1 -> CaCU + CO2| + H2O

(2) the action of heat on a hydrogencarbonate,

2HCO3" -» H2O + CO2| + COf-

(3) the action of heat on a metal carbonate, other than those of the alkali metals or barium (see later, p. 185). Industrially, carbon dioxide is obtained in large quantities by heating limestone:

CaCO3'->CaO + CO2f

It is obtained as a by-product in the fermentation of sugars to give alcohols:

C6H1206 ^ 2C2H5OH + 2C02

Appreciable quantities are also obtained as a by-product in the manufacture of hydrogen from naphtha-gaseous hydrocarbons. In this process the gaseous hydrocarbon and superheated steam under a pressure of about 10 atmospheres and at a temperature of 1000 K are passed over a nickel-chromium catalyst. Carbon monoxide and hydrogen are produced:

CnHm + nH20 -* nCO+ ^--^ H2

The hydrocarbons used depend on availability. Natural gas is now

GROUP IV 181

being used by some large industrial organisations but others use petroleum from a refinery. The second stage in the process is the so-called %water-gas shift' reaction; this reaction was originally used with vwater-gas'—a mixture of CO and H2 obtained by passing superheated steam through white hot coke. The gaseous mixture containing an excess of steam still at 10 atmospheres pressure, is passed at 700 K over an iron catalyst when the carbon monoxide reacts with the steam to form carbon dioxide and hydrogen:

CO + H2O -> CO2 + H2

In one process the carbon dioxide is removed using potassium carbonate solution, potassium hydrogencarbonate being produced:

K2CO3 + H2O + CO2 -» 2KHCO3

This reaction can be reversed by heat and the potassium carbonate and carbon dioxide recovered. (Other compounds which absorb carbon dioxide and evolve it again at a lower temperature are also in common usage*).

STRUCTURE

Carbon dioxide has a linear structure. The simple double-bonded formula, however, does not fully explain the structure since the measured carbon-oxygen bond lengths are equal but intermediate between those expected for a double and a triple bond. A more accurate representation is, therefore,obtained by considering carbon dioxide as a resonance hybrid of the three structures given below:

O =C^O <-> O=C=O <-> O—C =O

PROPERTIES

Carbon dioxide is a colourless gas which is virtually odourless and tasteless. Its density, relative to air, is 1.53; hence it accumulates at

* Some of the carbon monoxide and hydrogen produced in the steam-naphtha reforming process react to form methane:

CO + 3H2 <±CH4 4- H2O

This reaction is an undesirable side reaction in the manufacture of hydrogen but utilised as a means of removing traces of carbon monoxide left at the end of the second stage reaction. The gases are passed over a nickel catalyst at 450 K when traces of carbon monoxide form methane. (Methane does not poison the catalyst in the Haber process -carbon monoxide docs.)

182 GROUP IV

the bottom of towers or wells in which it is being prepared, and may reach dangerous concentrations there. (Carbon dioxide does not support respiration, but it is not toxic.) Its critical point is 304 K. i.e. it may be compressed to a liquid below this temperature. However, if carbon dioxide is cooled rapidly (for example by allowing compressed gas to escape through a valve) solid carbon dioxide is formed. This sublimes at 195 K and atmospheric pressure; it is a white solid, now much used as a refrigerant (known as kdry ice' or ^DrikokT), since it leaves no residue after sublimation.

Chemically, carbon dioxide is not very reactive, and it is often used as an inactive gas to replace air when the latter might interact with a substance, for example in the preparation of chromium(II) salts (p. 383). Very reactive metals, for example the alkali metals and magnesium can, however, continue to burn in carbon dioxide if heated sufficiently, for example

4K 4- 3CO2 -> 2K2CO3 + C

Carbon dioxide reacts with a solution of a metal hydroxidegiving the carbonate, which may be precipitated, for example

Ca2+ + 2OH~ + CO2 -> CaCO3i + H2O

This reaction is used as a test for carbon dioxide. Passage of an excess of carbon dioxide produces the soluble hydrogencarbonate :

CaCO3 + CO2 + H2O -» Ca2+ + 2HCOJ

The hydrogencarbonate ion, produced in nature by this reaction, is one of the main causes of temporary hardness in water. Carbon dioxide is fairly soluble in water, 1 cm3 dissolving 1.7cm3 of the gas at stp. The variation of solubility with pressure does not obey Henry's law, since the reaction

CO2 + H2O^=^~

takes place to a small extent, forming carbonic acid (see below).

USES

Carbon dioxide is used in the manufactureof sodium carbonate by the ammonia-soda process, urea, salicyclic acid (for aspirin), fire extinguishers and aerated water. Lesser amounts are used to transfer heat generated by an atomic reactor to water and so produce steam and electric power, whilst solid carbon dioxide is used as a refrigerant, a mixture of solid carbon dioxide and alcohol providing a good low-temperature bath (195 K) in which reactions can be carried out in the laboratory.

G R O U P IV 183

CARBONIC ACID AND CARBONATES

The following equilibria apply to a solution of carbon dioxide in water :

CO2 + H2O ^H2CO3 ^ H+ +HCO

carbonic acid

The amount of carbonic acid present, undissociated or dissociated, is only about 1% of the total concentration of dissolved carbon dioxide. Carbonic acid, in respect of its dissociation into hydrogen and hydrogencarbonate ions, is actually a stronger acid than acetic acid ; the dissociation constant is:

(cf. Ka - 1.8 x 1CT5 mol T 1 for acetic acid)

But a solution of carbon dioxide in water behaves as a very weak acid since the effective dissociation constant K' is given by :

Since carbonic acid is a weak acid, its salts are hydrolysed in aqueous solution :

CO23 + H2O

HCOJ + H2O - OH~ + H2CO

Although both these reactions lie largely to the left, soluble carbonates (i.e. those of the alkali metals) are alkaline in aqueous solution, and the hydrogencarbonates are very feebly alkaline. The equilibria are displaced to the right on addition of an acid and soluble carbonates can therefore be titrated with acids and indeed sodium carbonate is used as a standard base. The titration curve is given below for 0.1 M hydrochloric acid being added to 100cm3 of 0.1 M alkali metal carbonate {Figure 8.4). At A all the CO23^ has been converted to HCO^ and at B all the HCO^ has been converted to CO^". Phenolphthalein changes colour between pH 8.3 and pH 10.0 and can be used to indicate point A whilst methyl orange. changing colour between pH 3.1 and pH 4.4, indicates point B.

Most metal carbonates are insoluble and they are precipitated either as the simple carbonate or as the basic carbonate when