Carbon Compounds — CO and CO₂

Why carbon forms two oxides

The group 14 elements have four electrons in their outermost shell and characteristically exhibit the +4 and +2 oxidation states. For carbon, both states are realised in its oxides: the dioxide $\ce{CO2}$ in which carbon is formally $+4$, and the monoxide $\ce{CO}$ in which carbon is formally $+2$. This single difference in oxidation state propagates into almost every other property — structure, acid–base nature, reducing power and biological behaviour.

NCERT also frames a wider periodic generalisation that is worth holding onto: among the dioxides of group 14, $\ce{CO2}$, $\ce{SiO2}$ and $\ce{GeO2}$ are acidic while $\ce{SnO2}$ and $\ce{PbO2}$ are amphoteric; among the monoxides, $\ce{CO}$ is neutral, $\ce{GeO}$ is distinctly acidic, and $\ce{SnO}$ and $\ce{PbO}$ are amphoteric. The neutrality of $\ce{CO}$ is therefore not an isolated quirk but the starting point of a trend.

Carbon monoxide — preparation

Carbon monoxide is produced whenever carbon or a carbon-containing fuel burns in a limited supply of oxygen or air. The direct oxidation of carbon in a deficiency of oxygen is the simplest route:

$\ce{2C(s) + O2(g) ->[\Delta] 2CO(g)}$

On a small scale, pure $\ce{CO}$ is obtained in the laboratory by dehydrating formic acid with concentrated sulphuric acid at 373 K, where the acid acts as a dehydrating agent:

$\ce{HCOOH ->[\text{conc. } H2SO4][373\ K] H2O + CO}$

On a commercial scale, $\ce{CO}$ is generated by passing steam over hot coke, producing a mixture of carbon monoxide and dihydrogen known as water gas (also called synthesis gas). If air is used instead of steam, a mixture of $\ce{CO}$ and $\ce{N2}$ called producer gas is obtained.

$\ce{C(s) + H2O(g) ->[473-1273\ K] CO(g) + H2(g)}$   (water gas)

$\ce{2C(s) + O2(g) + 4N2(g) ->[1273\ K] 2CO(g) + 4N2(g)}$   (producer gas)

Both water gas and producer gas are important industrial fuels, because the $\ce{CO}$ they contain can be burned further to $\ce{CO2}$ with liberation of heat.

Structure of CO

In the carbon monoxide molecule there are one sigma and two π bonds between carbon and oxygen, giving a carbon–oxygen triple bond, written $\ce{:C#O:}$. Crucially, carbon retains a lone pair. This lone pair on carbon is responsible for the molecule acting as an electron-pair donor (a ligand), which is why $\ce{CO}$ reacts with certain metals when heated to form metal carbonyls — and it is also the structural feature behind its toxicity, discussed below.

Properties of CO — neutral and reducing

Carbon monoxide is a colourless, odourless and almost water-insoluble gas. Two properties dominate its NEET profile, and both are easy to confuse with the dioxide if not learned deliberately.

First, $\ce{CO}$ is a neutral oxide. It does not form an acid with water under ordinary conditions and does not behave like the acidic $\ce{CO2}$. Second, $\ce{CO}$ is a powerful reducing agent: it reduces almost all metal oxides other than those of the alkali and alkaline earth metals, aluminium, and a few transition metals. This reducing power is exploited in metallurgy to extract metals from their oxide ores, where $\ce{CO}$ is oxidised to $\ce{CO2}$:

$\ce{Fe2O3(s) + 3CO(g) ->[\Delta] 2Fe(s) + 3CO2(g)}$

$\ce{ZnO(s) + CO(g) ->[\Delta] Zn(s) + CO2(g)}$

Why CO is poisonous

The highly poisonous nature of carbon monoxide is a direct consequence of the lone pair on carbon. $\ce{CO}$ uses that lone pair to form a complex with haemoglobin, the oxygen-carrying pigment of red blood corpuscles. This carboxyhaemoglobin complex is about 300 times more stable than the oxygen–haemoglobin complex.

Because the $\ce{CO}$–haemoglobin bond is so much stronger, $\ce{CO}$ preferentially occupies the binding sites and prevents haemoglobin from carrying oxygen around the body. The tissues are starved of oxygen, and the result can ultimately be death. This is exactly why incomplete combustion in poorly ventilated spaces is so dangerous — a colourless, odourless gas gives no sensory warning.

Carbon dioxide — preparation

Carbon dioxide is formed by the complete combustion of carbon and carbon-containing fuels in excess of air:

$\ce{C(s) + O2(g) ->[\Delta] CO2(g)}$

$\ce{CH4(g) + 2O2(g) ->[\Delta] CO2(g) + 2H2O(g)}$

In the laboratory it is conveniently prepared by the action of dilute hydrochloric acid on calcium carbonate — the standard bench reaction in a Kipp's apparatus:

$\ce{CaCO3(s) + 2HCl(aq) -> CaCl2(aq) + CO2(g) + H2O(l)}$

On a commercial scale, $\ce{CO2}$ is obtained by heating limestone (thermal decomposition of calcium carbonate):

$\ce{CaCO3(s) ->[\Delta] CaO(s) + CO2(g)}$

Structure of CO₂

In the $\ce{CO2}$ molecule the carbon atom undergoes sp hybridisation. The two sp hybrid orbitals of carbon overlap with two p orbitals of the oxygen atoms to form two sigma bonds, while the remaining two electrons of carbon are involved in pπ–pπ bonding with the oxygen atoms. The molecule is therefore linear, $\ce{O=C=O}$, with both C–O bonds of equal length (115 pm) and a bond angle of $180^\circ$.

Because the two C–O bond dipoles are equal and point in exactly opposite directions, they cancel, so $\ce{CO2}$ has no net dipole moment. The bonding is best represented by a set of resonance structures, the central canonical form being the symmetric $\ce{O=C=O}$.

Properties of CO₂ — the acidic oxide

Carbon dioxide is a colourless and odourless gas. Its low solubility in water gives it immense biochemical and geochemical importance. The defining chemical property is that $\ce{CO2}$ is an acidic oxide. With water it forms carbonic acid, $\ce{H2CO3}$, a weak dibasic acid that dissociates in two steps:

$\ce{CO2 + H2O <=> H2CO3}$

$\ce{H2CO3(aq) + H2O(l) <=> HCO3^-(aq) + H3O^+(aq)}$

$\ce{HCO3^-(aq) + H2O(l) <=> CO3^{2-}(aq) + H3O^+(aq)}$

This $\ce{H2CO3}/\ce{HCO3^-}$ buffer system helps maintain the pH of blood between 7.26 and 7.42. Being acidic, $\ce{CO2}$ also combines with alkalis to form metal carbonates; with limited alkali the carbonate forms, and with excess $\ce{CO2}$ the soluble bicarbonate is produced:

$\ce{CO2 + 2NaOH -> Na2CO3 + H2O}$

$\ce{CO2 + NaOH -> NaHCO3}$   (excess CO₂)

Several characteristic uses of $\ce{CO2}$ also flow from its physical behaviour. Solid $\ce{CO2}$, obtained by allowing liquefied $\ce{CO2}$ to expand rapidly, is dry ice, used as a refrigerant for ice-cream and frozen food. Gaseous $\ce{CO2}$ carbonates soft drinks; being heavy and a non-supporter of combustion it is used in fire extinguishers; and a substantial amount is consumed in the manufacture of urea.

The environmental angle

Carbon dioxide is normally present to the extent of about 0.03% by volume in the atmosphere, and it is removed from the air by photosynthesis, the process by which green plants convert atmospheric $\ce{CO2}$ into carbohydrates such as glucose:

$\ce{6CO2 + 12H2O ->[\text{chlorophyll}][h\nu] C6H12O6 + 6O2 + 6H2O}$

Unlike $\ce{CO}$, carbon dioxide is not poisonous. However, the increased combustion of fossil fuels and the decomposition of limestone for cement manufacture in recent years appear to be raising the atmospheric $\ce{CO2}$ content. This may increase the greenhouse effect and thus raise the temperature of the atmosphere, with potentially serious consequences. This is the chemical basis of the global-warming narrative that NEET tests through statement-matching questions.

CO versus CO₂ at a glance

The entire subtopic can be compressed into a single comparison. Most NEET questions on the oxides of carbon are answered by reading off one row of this table.