Oxides of Carbon
Two important oxides of carbon are carbon monoxide, CO and carbon dioxide, CO2.
11.8.1 Carbon Monoxide
Direct oxidation of C in limited supply of oxygen or air yields carbon monoxide.
2C (s) + O2 (g) → 2CO(g)
On small scale pure CO is prepared by dehydration of formic acid with concentrated H2SO4 at 373 K
HCOOH → H2O + CO
On commercial scale it is prepared by the passage of steam over hot coke. The mixture of CO and H2 thus produced is known as water gas or synthesis gas.
C (s) + H2O (g) → CO(g) + H2 (g)
When air is used instead of steam, a mixture of CO and N2 is produced, which is called producer gas.
2C (s) + O2 (g) + 4N2 (g) → 2CO (g) + 4N2 (g)
Water gas and producer gas are very important industrial fuels. Carbon monoxide in water gas or producer gas can undergo further combustion forming carbon dioxide with the liberation of heat.
Carbon monoxide is a colourless, odourless and almost water insoluble gas. It is a powerful reducing agent and reduces almost all metal oxides other than those of the alkali and alkaline earth metals, aluminium and a few transition metals.
This property of CO is used in the extraction of many metals from their oxides ores.
Fe2O3 (s) + 3CO (g) → 2Fe(s) + 3CO2 (g)
ZnO (s) + CO(g) → Zn (s) + CO2 (g)
In CO molecule, there are one sigma and two bonds between carbon and oxygen, :C ≡ O: .
Because of the presence of a lone pair on carbon, CO molecule acts as a donor and reacts with certain metals when heated to form metal carbonyls.
The highly poisonous nature of CO arises because of its ability to form a complex with haemoglobin, which is about 300 times more stable than the oxygen-haemoglobin complex.
This prevents haemoglobin in the red blood corpuscles from carrying oxygen round the body and ultimately resulting in death.
11.8.2 Carbon Dioxide
It is prepared by complete combustion of carbon and carbon containing fuels in excess of air.
C (s) + O2 (g) → CO2 (g)
CH4(g) + 2O2 (g) → CO2 (g) + 2H2O (g)
In the laboratory it is conveniently prepared by the action of dilute HCl on calcium carbonate.
CaCO3 (s) + 2HCl → CaCl2 (aq) + CO2 (g) + H2O (l)
On commercial scale it is obtained by heating limestone. It is a colourless and odourless gas. Its low solubility in water makes it of immense biochemical and geo-chemical importance.
With water, it forms carbonic acid, H2CO3 which is a weak dibasic acid and dissociates in two steps:
H2CO3 (aq) + H2O (l) → HCO3-(aq) + H3O+(aq)
HCO3-(aq) + H2O (l) → CO32- (aq) + H3O+ (aq)
H2CO3/HCO3 – buffer system helps to maintain pH of blood between 7.26 to 7.42. Being acidic in nature, it combines with alkalies to form metal carbonates.
Carbon dioxide, which is normally present to the extent of ~ 0.03 % by volume in the atmosphere, is removed from it by the process known as photosynthesis.
It is the process by which green plants convert atmospheric CO2 into carbohydrates such as glucose. The overall chemical change can be expressed as:
6CO2 + 12H2O → C6H12O6 + 6O2 + 6H2O
By this process plants make food for themselves as well as for animals and human beings. Unlike CO, it is not poisonous.
But the increase in combustion of fossil fuels and decomposition of limestone for cement manufacture in recent years seem to increase the CO2 content of the atmosphere.
This may lead to increase in green house effect and thus, raise the temperature of the atmosphere which might have serious consequences. Carbon dioxide can be obtained as a solid in the form of dry ice by allowing the liquified CO2 to expand rapidly.
Dry ice is used as a refrigerant for ice-cream and frozen food. Gaseous CO2 is extensively used to carbonate soft drinks. Being heavy and non-supporter of combustion it is used as fire extinguisher. A substantial amount of CO2 is used to manufacture urea.
In CO2 molecule carbon atom undergoes sp hybridisation. Two sp hybridised orbitals of carbon atom overlap with two p orbitals of oxygen atoms to make two sigma bonds while other two electrons of carbon atom are involved in p– p bonding with oxygen atom.
This results in its linear shape [with both C–O bonds of equal length (115 pm)] with no dipole moment. The resonance structures are shown below:
Resonance structures of carbon dioxide is shown
11.8.3 Silicon Dioxide, SiO2
95% of the earth’s crust is made up of silica and silicates. Silicon dioxide, commonly known as silica, occurs in several crystallographic forms.
Quartz, cristobalite and tridymite are some of the crystalline forms of silica, and they are interconvertable at suitable temperature.
Silicon dioxide is a covalent, three-dimensional network solid in which each silicon atom is covalently bonded in a tetrahedral manner to four oxygen atoms.
Each oxygen atom in turn covalently bonded to another silicon atoms as shown in diagram (Figure). Each corner is shared with another tetrahedron.
The entire crystal may be considered as giant molecule in which eight membered rings are formed with alternate silicon and oxygen atoms.
Silica in its normal form is almost nonreactive because of very high Si—O bond enthalpy. It resists the attack by halogens, dihydrogen and most of the acids and metals even at elevated temperatures.
However, it is attacked by HF and NaOH.
SiO2 + 2NaOH → Na2SiO3 + H2O
SiO2 + 4HF → SiF4 + 2H2O
Quartz is extensively used as a piezoelectric material; it has made possible to develop extremely accurate clocks, modern radio and television broadcasting and mobile radio communications.
Silica gel is used as a drying agent and as a support for chromatographic materials and catalysts. Kieselghur, an amorphous form of silica is used in filtration plants.
They are a group of organosilicon polymers, which have (R2SiO) as a repeating unit. The starting materials for the manufacture of silicones are alkyl or aryl substituted silicon chlorides, RnSiCl(4–n), where R is alkyl or aryl group.
When methyl chloride reacts with silicon in the presence of copper as a catalyst at a temperature 573K various types of methyl substituted chlorosilane of formula MeSiCl3, Me2SiCl2, Me3SiCl with small amount of Me4Si are formed.
Hydrolysis of dimethyldichlorosilane, (CH3)2SiCl2 followed by condensation polymerisation yields straight chain polymers.
The chain length of the polymer can be controlled by adding (CH3)3SiCl which blocks the ends as shown below:
Silicones being surrounded by non-polar alkyl groups are water repelling in nature. They have in general high thermal stability, high dielectric strength and resistance to oxidation and chemicals.
They have wide applications. They are used as sealant, greases, electrical insulators and for water proofing of fabrics. Being biocompatible they are also used in surgical and cosmetic plants.
A large number of silicates minerals exist in nature. Some of the examples are feldspar, zeolites, mica and asbestos. The basic structural unit of silicates is SiO4 4– (Figure) in which silicon atom is bonded to four oxygen atoms in tetrahedron fashion.
In silicates either the discrete unit is present or a number of such units are joined together via corners by sharing 1,2,3 or 4 oxygen atoms per silicate units.
When silicate units are linked together, they form chain, ring, sheet or three-dimensional structures. Negative charge on silicate structure is neutralised by positively charged metal ions.
If all the four corners are shared with other tetrahedral units, three-dimensional network is formed. Two important man-made silicates are glass and cement.
If aluminium atoms replace few silicon atoms in three-dimensional network of silicon dioxide, overall structure known as aluminosilicate, acquires a negative charge.
Cations such as Na+, K+ or Ca2+ balance the negative charge. Examples are feldspar and zeolites. Zeolites are widely used as a catalyst in petrochemical industries for cracking of hydrocarbons and isomerisation, e.g., ZSM-5 (A type of zeolite) used to convert alcohols directly into gasoline.
Hydrated zeolites are used as ion exchangers in softening of “hard” water.