Carbon and its Compounds: Oxides of Carbon, Carbon(IV) oxide

Oxides of Carbon

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Carbon forms two important gaseous oxides. These are: carbon monoxide or carbon(II) oxide (CO) and carbon dioxide or carbon(IV) oxide (CO₂).

This article focuses on carbon(IV) oxide (CO₂), its structure, physical properties, chemical properties, preparation, test, uses, processes that introduces it into (and removes it from) the atmosphere and its toxicity, alongside the comparison of carbon(II) oxide and carbon(IV) oxide.

 

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Carbon(IV) oxide, (CO₂)

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Structure

Carbon(IV) oxide (CO₂) is a linear molecule, O=C=O, (with a 180° bond angle), with each C=O bond being a double bond. Simply put, one central carbon atom double-bonded (covalently by sharing two pairs of electrons each (through each double bonds) with the two terminal oxygen atoms. It is a typical acidic oxide.

The structure of CO₂ is shown below:

 

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Please read on carbon, its structure and crystalline allotropes here.

 

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Physical properties of Carbon(IV) oxide (CO₂)

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• Carbon(IV) oxide is a colorless, odorless gas with a slightly sharp/sour taste.

• It is denser than air (about 1.5 times as dense). In this regard, it can be collected by downward delivery, and used to "pour" over flames to extinguish them.

• Solid CO₂ (called dry ice) transitions directly from a solid to a gas at -78.5°C (-109.3°F} at standard pressure without passing through the liquid phase. This process is termed sublimation.

• It is fairly soluble in water, forming a weakly acidic solution:

CO₂_(g) + H₂O_(l) ⇌ H₂CO₃_(aq) (carbonic acid)

Solubility increases under higher pressure.

• CO₂ does not burn and does not support combustion. For this reason, it is used in fire extinguishers.

• The molecular structure of CO₂ is linear and has symmetrical geometry with a zero dipole moment.

Please read on bases and alkalis, their classifications, physical and chemical properties here.

 

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Chemical properties of Carbon(IV) oxide (CO₂)

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• It reacts with water by dissolving slightly to form a weakly acidic solution (carbonic acid). This is shown via the equation below:

CO₂_(g) + H₂O_(l) ⇌ H₂CO₃_(aq) (carbonic acid)

Solubility increases under higher pressure.

• As an acidic oxide, it reacts with alkalis to form carbonates and hydrogencarbonates.

CO₂_(g) + 2NaOH_(aq) → Na₂CO₃_(aq) + H₂O_(l)

CO₂_(g) + NaOH_(aq) → NaHCO₃_(aq)

 CO₂_(g) + Ca(OH)₂_(aq) → CaCO₃_(s) + H₂O_(l)

• Passing CO₂ through a clear calcium hydroxide solution or lime water - Ca(OH)₂, turns it milky due to the formation of insoluble calcium carbonate (CaCO₃).

Ca(OH)_2(aq) + CO_2(g) → CaCO_3(s)↓ + H₂O_(l)

Note: Passing excess CO₂ clears the milkiness because it forms soluble calcium bicarbonate - Ca(HCO₃)₂.

CaCO₃_(s) + H₂O_(l) + CO₂_(g) → Ca(HCO₃)₂_(aq)

• CO₂ reacts with active metals, e.g, burning reactive magnesium (Mg), causing the metal to reduce the carbon(IV) oxide to carbon.

2Mg_(s) + CO_2(g) → 2MgO + C_(s)

• In the leaf of green plants, CO₂ reacts with water in the presence of sunlight to synthesize sugar and release oxygen. This reaction is termed photosynthesis.

6CO_2(g) + 6H₂O_(l) - Sunlight→ C₆H₁₂O_6(s) + 6H₂O_(l)

Please read on the physical and chemical properties of acids, preparation, test and uses of acids here.

 

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Preparation of Carbon(IV) oxide (CO₂)

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Laboratory preparation

Carbon(IV) oxide can be produced in the laboratory by the action of dilute acid on a carbonate. CO₂ gas is liberated during this reaction (as salt and water are also produced). This preparation is carried out in a Kipp's apparatus (or flask + thistle funnel).

CaCO₃_(s) + 2HCl_(aq) → CaCl₂_(aq) + H₂O_(l) + CO₂_(g)

The CO₂ gas liberated in the above reaction is collected by downward delivery (or upward displacement of air), since CO₂ is denser than air. Pure samples of CO₂ is not usually collected over water since it is only fairly soluble in water.

 

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Industrial/large-scale preparation of CO2

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CO₂ is produced industrially through the following methods:

1. By the thermal decomposition (or calcination) of limestone. See the equation below:

•  CaCO₃_(s) --(heat) --> CaO(l_(s)+ CO₂_(g)

2. As a by-product of fermentation (respiration of yeast in brewing/baking). See the equation below:

•  C₆H₁₂O₆_(s) → 2C₂H₅OH_(aq) + 2CO₂_(g)

3. As a by-product of the combustion of fossil fuels.

• Fuel + Oxygen → CO₂ + 2H₂O + Energy

• CH_4(g) + 2O_2(g) → CO_2(g) + H₂O_(l) + Energy

Please read more on the amorphous allotropes of carbon here.

 

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Test for Carbon(IV) oxide (CO₂)

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Pass the unknown gas into freshly prepared limewater [Ca(OH)₂ solution]. It turns the limewater milky (or cloudy) due to formation of insoluble calcium carbonate:

•  CO₂_(g) + Ca(OH)₂_(aq) → CaCO₃_(s)↓ + H₂O_(l)

If excess CO₂ is passed in for a long time, the milkiness disappears again as the insoluble CaCO₃ is converted to soluble calcium hydrogencarbonate:

•  CaCO₃_(s) + H₂O_(l) + CO₂_(g) → Ca(HCO₃)₂_(aq)

This "milky appearance that then clears" behavior can help us confirm CO₂ specifically (and explains related natural phenomena like the formation and dissolving of stalactites/stalagmites and hard water).

You can read on relative acidity and alkalinity (pH scale) and its calculations here.

 

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Processes that introduce Carbon(IV) oxide (CO₂) into the Atmosphere

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Carbon(IV) oxide (CO₂) is introduced into the atmosphere through the following processes:

• Respiration: During respiration, living organisms (plants, animals and microorganisms) releases CO₂ as a waste product of cellular respiration into the environment.

• Combustion: The oxidation of fuel is called combustion when it happens rapidly, releases a large amount of heat and light, and usually creates a visible flame. The combustion of fuels, both fossil fuels (coal, petroleum, natural gas) and biomass (wood) releases CO₂ into the environment.

• Decay/Decomposition/Putrefaction: Putrefaction is the process whereby bacteria and fungi break down dead animal or plant matter, especially proteins, causing a very bad smell. The decomposition of dead plant and animal matter by microorganisms releases CO₂ into the environment.

• Thermal Decomposition: Thermal decomposition is a chemical reaction whereby a single compound breaks down into two or more simpler substances when heated. Rock that contain carbon (carbonate rocks) e.g. limestone, can undergo thermal decomposition through chemical weathering to release CO₂ into the environment. Also, industrial calcination also releases CO₂ into the environment.

• Fermentation: This refers to a natural metabolic process where microorganisms like yeast and bacteria convert carbohydrates into acids, gases, or alcohol in the absence of oxygen. CO₂ is released during alcoholic fermentation.

• Volcanic activities introduce carbon(IV) oxide (CO₂) into the atmosphere through two primary mechanisms. These are explosive eruptions and passive degassing. CO₂ is typically the second most abundant gas evolved during this process, making up 10% to 40% of a volcano's gas emissions. Also, natural outgassing introduces CO₂ into the atmosphere.

You can read on salts, types, properties and uses of salts in chemistry here.

 

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Processes that remove Carbon(IV) oxide (CO₂) from the Atmosphere

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Carbon(IV) oxide (CO₂) is removed from the atmosphere through various processes. They include:

• Photosynthesis by green plants, algae, and some bacteria, which absorb CO₂ and convert it into glucose (simple sugar) and other organic compounds, releasing O₂ as a byproduct.

• Dissolution or (dissolving) of CO₂ in oceans, lakes, and rainwater, forming carbonic acid, bicarbonate and carbonate systems. In this regard, the oceans act as a major carbon "sink".

• CO₂ is used in the formation of carbonate sediments/rocks, (e.g. shells and skeletons of marine organisms made of calcium carbonate, eventually forming limestone).

• Through the uptake and storage of CO₂ in plant biomass and soils.

Note: This continuous exchange (addition and removal) of CO₂ between the atmosphere, living organisms, oceans, and rocks is known as the carbon cycle, and it normally keeps atmospheric CO₂ levels roughly balanced. However, it is noteworthy to state that excess CO₂ from fossil fuel combustion and deforestation is currently increasing atmospheric concentration, contributing to the greenhouse effect and global warming.

The image below shows the carbon cycle.

You can read on types of mixtures in chemistry here.

 

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Toxicity of Carbon(IV) oxide (CO₂)

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CO₂ is not directly poisonous in the same chemical sense as carbon(II) oxide (CO) (since it does not bind to haemoglobin). However, it is still dangerous, mainly because of oxygen displacement and asphyxiation; and its direct effect on breathing regulation.

In high concentrations, CO₂ displaces oxygen in an enclosed space, reducing the available oxygen and causing suffocation (asphyxiation).

Elevated CO₂ levels in the blood directly stimulate faster, labored breathing and, at high enough levels, cause headache, dizziness, confusion, and can lead to unconsciousness or death if oxygen is insufficient.

Since CO₂ is denser than air, it can accumulate in low-lying, poorly ventilated spaces (e.g. wells, mines, silos and fermentation cellars), creating serious suffocation hazards even though the gas itself is not "poisonous" in the way CO is.

Please read more on carbon(II) oxide (CO) here.

 

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Uses of Carbon(IV) oxide (CO₂)

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• It is used in fire extinguishers. This is so because it does not support combustion, and is also denser than air. As a result, it blankets the fire, cutting off oxygen supply.

• It is used to carbonate soft drinks and sparkling water (when dissolved under pressure).

• Solid CO₂ (dry ice) is used for refrigeration, and in creating stage/theatrical fog effects.

• It is crucial in photosynthesis, a process through which green plants manufacture food (sugars) and indirectly, other organic compounds:

6CO₂ + 6H₂O --(light, chlorophyll) → C₆H₁₂O₆ + 6O₂

• It can be used as an inert atmosphere for certain chemical processes, and in some fire-suppression systems.

• It is used industrially in the Solvay process for manufacturing sodium carbonate (Na₂CO₃).

You can attempt questions and answers in chemistry here.

 

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Summary Comparison Table for CO vs CO₂

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Property	CO	CO₂
Bonding	Triple bond (C≡O)	Double bonds (O=C=O), linear
Nature of oxide	Neutral oxide	Acidic oxide
Color/odor	Colorless, odorless	Colorless, odorless
Density vs air	Slightly less dense	About 1.5× denser
Solubility in water	Almost insoluble	Fairly soluble (forms weak acid)
Combustion	Burns with blue flame	Does not burn/support combustion
Test	Burns blue; product turns limewater milky	Turns limewater milky, then clear in excess
Main hazard	Highly toxic. Binds haemoglobin, causes hypoxia / Asphyxiant	Displaces oxygen, not directly poisonous
Key source	Incomplete combustion	Respiration, complete combustion, decay and volcanic activity
Key removal process	Slow atmospheric oxidation to CO₂	Photosynthesis, dissolving in oceans, carbonate rock formation

 

Please read the next topic (Trioxocarbonates) here.