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# Combustion

Alkanes are great fuels. One litre of petrol contains roughly 33.6 megajoules of energy and provides over 8000 calories1. If you could eat petrol, this would be enough to feed an average man for over three days! When burnt, all the stored energy is released into the environment as heat. This is why we combust alkanes as fuels for our cars, planes, and other engines; to heat our homes, and to power our electronics.

• In particular, we'll focus on the combustion of hydrocarbons like alkanes.
• We'll start by defining combustion before looking at both complete and incomplete combustion.
• After that, we'll practice writing combustion equations.
• To finish, we'll consider the environmental impacts of combustion. We'll also learn how we can limit these impacts through measures such as flue gas sulphurisation and catalytic converters.

## Combustion definition

Many of us have appreciated the joy of standing around a bonfire on a chilly winter's evening, perhaps toasting marshmallows or sipping hot chocolate. In fact, fire is no secret - it has been used by our ancestors for over 1 million years2! The flames you see are the visible result of a combustion reaction.

Combustion is a reaction that involves burning a fuel such as coal, gas, or petrol, usually in oxygen. It is exothermic, meaning that it releases lots of energy into the environment in the form of heat.

Combustion reactions have multiple different uses in everyday life, thanks to the heat energy they release. For example, we use combustions reactions to:

• Warm our homes.
• Cook our food.
• Run our factories and power stations.
• Produce electricity.

In fact, it is quite hard to think of a part of our modern existence that doesn't rely on combustion in one way or another!

## Types of combustion reaction

There are two different types of combustion reactions:

• Complete combustion.
• Incomplete combustion.

They vary in their conditions, products, and the relative amounts of energy that they release. Let’s take a brief look.

### Complete combustion

In complete combustion, a fuel is burnt in excess oxygen. We commonly use hydrocarbons as fuel. Burning any hydrocarbon in this way oxidises its carbon and hydrogen atoms, producing both carbon dioxide and water. Perhaps most importantly, complete combustion also releases lots of heat energy into the environment.

The equation for the complete combustion of methane is given below:

$$CH_4(g)+2O_2(g)\rightarrow CO_2(g)+2H_2O(g)\qquad \Delta H^\circ = -802.3\space kJ\space mol^{-1}$$

Note the negative enthalpy change. This shows that the reaction is exothermic.

Don't worry - we'll go through how you write combustion equations in just a second. Before we do, make sure you check out Enthalpy Changes if you aren't too sure about what we mean by exothermic reactions.

### Incomplete combustion

In incomplete combustion, the fuel is burnt in limited oxygen. This means that there isn’t enough oxygen to fully oxidise all the fuel's carbon atoms into carbon dioxide. Instead, they are partially oxidised into carbon monoxide. If oxygen is really limited, the carbon atoms aren't oxidised at all - instead, they are released as pure carbon, in the form of soot. Although still exothermic, incomplete combustion is much less efficient than complete combustion and so releases less energy.

Look at the following two equations for the incomplete combustion of methane:

$$CH_4(g)+\frac{3}{2}O_2(g)\rightarrow CO(g)+2H_2O(g)\qquad \Delta H^\circ = -519.3\space kJ\space mol^{-1}$$ $$CH_4(g)+O_2(g)\rightarrow C(s)+2H_2O(g)\qquad \Delta H^\circ = -408.8\space kJ\space mol^{-1}$$

Compare them to the equation we gave earlier for the complete combustion of methane.

• Complete combustion requires more oxygen than incomplete combustion. Here, complete combustion needs a minimum of two moles of oxygen for each mole of methane, whereas incomplete combustion only requires just one.
• Complete combustion releases more energy than incomplete combustion.

Be aware that combustion typically releases a mixture of all three carbon products: CO2, CO, and C. However, we can favour one product or another by controlling the amount of oxygen present when we burn the fuel.

Want to see complete and incomplete combustion in action? Just look towards a common piece of chemical apparatus: the Bunsen burner.

When working with a Bunsen burner in the lab, you’ll know that closing the air hole results in the safety flame. This flame is orange-yellow. On the other hand, opening the air hole produces a blue flame. The safety flame is so named because it is a lot easier to see than the blue flame, and is also a lot less hot. This is because oxygen is limited and so the fuel burns in an incomplete combustion reaction. Using the safety flame causes the bottom of any beakers held above the Bunsen burner to go black and sooty due to the carbon particles formed, which is why the safety flame is sometimes nicknamed the 'dirty' flame. In contrast, the blue flame uses complete combustion. It releases a lot more energy when burning, resulting in a cleaner, hotter, more dangerous flame.

Fig. 1 - A diagram showing complete and incomplete combustion when using Bunsen burners. Closing the air hole results in incomplete combustion (left) and produces a yellow safety flame, whilst opening the air hole results in complete combustion and produces a hotter blue flame

## Combustion equations

Ready to learn how to write combustion equations? Let's give it a go. But first, one quick note: Throughout your time as a chemist, you’ll probably have been told to always write whole number equations. Combustion equations are an exception to this rule. Here, reacting half a mole of oxygen molecules is perfectly acceptable.

How come we can get away with half a mole of oxygen? Well, it is because it is considered standard to write combustion reactions using one mole of the fuel. This is due to the definition of standard enthalpy change of combustion, which looks at the enthalpy change when one mole of a substance is burnt in excess oxygen under standard conditions. Don't worry if you haven't heard this term before - we have a whole article dedicated to standard enthalpies over at Enthalpy Changes. Click the link to find out more.

Here's how you go about writing and balancing equations for the (complete) combustion of hydrocarbons:

1. Write out an unbalanced equation, remembering that the products will be carbon dioxide and water. You should also remember that combustion equations use just one mole of fuel, but the number of moles of carbon dioxide, water, and oxygen can vary.
2. To balance your equation, first focus on carbon. Count up the number of carbon atoms in your fuel and balance the equation by adding extra moles of carbon dioxide to the right-hand side.
3. Next, look at hydrogen. Count up the number of hydrogen atoms in your fuel and balance the equation by adding extra moles of water to the right-hand side.
4. Finally, check on oxygen. Count up the number of oxygen atoms in both the carbon dioxide and water, and balance the equation by adding extra moles of oxygen to the left-hand side.

Writing equations for incomplete combustion follows a similar process. However, you should remember that you produce carbon monoxide (or pure carbon) instead of carbon dioxide. As a result, you'll need fewer moles of oxygen. However, you will still need the same number of moles of the carbon-based product and the same number of moles of water.

Write an equation for:

1. The complete combustion of propane.
2. The incomplete combustion of propane, producing just carbon monoxide.
3. The incomplete combustion of propane, producing one mole of carbon monoxide and two moles of carbon.

Let's start with part a. In the unbalanced equation, we burn one mole of propane (C3H8) in an unknown quantity of oxygen (O2) to produce unknown quantities of carbon dioxide (CO2) and water (H2O):

$$C_3H_8+\_ O_2\rightarrow \_ CO_2+\_ H_2O$$

C3H8 contains three carbon atoms. Therefore, we produce three moles of CO2:

$$C_3H_8+\_ O_2\rightarrow 3CO_2+\_ H_2O$$

C3H8 also contains eight hydrogen atoms. Note that each H2O molecule contains two hydrogen atoms. Therefore, we produce just four moles of H2O:

$$C_3H_8+\_ O_2\rightarrow 3CO_2+4H_2O$$

Now we just need to balance the oxygen atoms. We have $$3(2)=6$$ oxygen atoms from our three moles of CO2, and $$4(1)=4$$ oxygen atoms from our four moles of H2O. This gives ten oxygen atoms in total. Once again, note that each O2 molecule contains two oxygen atoms. Therefore, we need 5 moles of O2. Here's our final answer:

$$C_3H_8+ 5O_2\rightarrow 3CO_2+4H_2O$$

In part b, we produce carbon monoxide (CO) instead of carbon dioxide. Again, we produce three moles of the carbon-based product and four moles of H2O. But this time, we only produce $$3(1)=3$$ plus $$4(1)=4$$ oxygen atoms, giving us a total of seven. Therefore, we only need $$\frac{7}{2}$$ moles of O2:

$$C_3H_8+\frac{7}{2}O_2\rightarrow 3CO+4H_2O$$

In part c, we produce just one mole of carbon monoxide (CO) and two moles of pure carbon (C). This gives us our three moles of carbon-based products. Again, we also produce four moles of H2O. In total, we have $$1(1)=1$$ plus $$2(0)=0$$ plus $$4(1)=4$$ oxygen atoms, which equals five. Therefore, we need just $$\frac{5}{2}$$ moles of O2:

$$C_3H_8+\frac{5}{2}O_2\rightarrow CO+2C+4H_2O$$

## Environmental impacts of combustion

As mentioned at the start of this article, combustion reactions play a big role in our lives. In particular, we rely on fuels made from hydrocarbons, such as coal, gas, and crude oil derivatives. For example, your car probably runs on petrol or diesel, both derived from crude oil, and your house might be warmed by a gas boiler. However, the pressure on governments, businesses, and consumers to move away from hydrocarbons and look toward renewable energy sources is steadily increasing. This is partly due to the products of hydrocarbon combustion reactions and their negative environmental impacts. We’ll explore them now.

### Carbon dioxide

As we learned earlier, the complete combustion of hydrocarbons releases carbon dioxide. Carbon dioxide can be a problem because of its status as a powerful greenhouse gas.

A greenhouse gas is a gas that traps solar radiation reflected from the Earth, instead of letting it pass through the atmosphere and back into outer space.

By trapping solar heat, greenhouse gases warm the planet and contribute to global warming. There's a clear trend: ever since the start of the Industrial Revolution, when we started burning hydrocarbons on a mass scale (and so rapidly increased the atmospheric levels of carbon dioxide particles), average global temperatures have rocketed upwards. Currently, our planet is about 1.1 °C warmer than in 1880, when temperature records began3. The rise in temperature has been accompanied by more and more cases of extreme weather, crop failure, and mass extinction. Of course, the correlation between rising CO2 levels and temperature could just be one massive coincidence - but the evidence increasingly says otherwise.

Fig. 2 - A diagram of the greenhouse effect. Greenhouse gas particles trap the sun’s radiation reflected back from the Earth and reemit part of it back at the planet, instead of letting it pass through into outer space

### Carbon monoxide and carbon particles

Unlike complete combustion, incomplete combustion doesn't produce carbon dioxide. However, it releases carbon monoxide and pure carbon instead. These products aren't much better than their greenhouse gas cousin! For example, carbon monoxide is highly toxic to humans and animals. It is also odourless and colourless, making it hard to detect and so even deadlier. Carbon particles from soot, on the other hand, can cause respiratory irritation, certain cancers, and global dimming. You don't want to be breathing any of these substances in!

### Sulphur dioxide

Unfortunately, carbon-based products and water aren't actually the only products of combustion. This is because our fuels are typically impure - they are riddled with contaminants. These impurities are then released back into the atmosphere when the fuels are burned. One such common impurity is sulphur, which burns to produce sulphur dioxide. Sulphur dioxide then reacts with oxygen and water in the air to form acid rain - a corrosive substance that damages buildings, statues, habitats, and plant life.

### Nitrous oxides

A further pollutant produced in combustion reactions is nitrous oxide. It is formed because some fuels, such as those used in combustion engines, require high temperatures to burn effectively. Such high temperatures cause nitrogen and oxygen from the air to react to produce nitrous oxides. Like sulphur dioxide, nitrous oxides can form acid rain. They also cause breathing difficulties and photochemical smog.

## Limiting the impacts of combustion

Global warming, acid rain, smog - after looking at the negative impacts of burning fuels such as hydrocarbons, you might wonder why we don't stop these combustion reactions altogether. Well, it is all a question of weighing up the pros and cons. Hydrocarbons are not only abundant and easy to get hold of, but also release a lot of energy when burned. This is why they make such great fuels, and so moving away from them completely is not an easy task. However, scientists have worked hard to come up with ways of reducing the effects of hydrocarbon combustion on the environment. These include:

• Flue gas desulphurisation.
• Catalytic converters.
• Carbon-neutral fuel alternatives.

### Flue gas desulphurisation

Flue gas is the gas produced when burning coal in power stations. It contains sulphur dioxide because of impurities in the fuel. We can remove the sulphur through flue gas desulphurisation, by reacting flue gas with either calcium oxide and water, or calcium carbonate and oxygen. This forms gypsum, which is a saleable product used to make plasterboard.

For example:

$$CaCO_3(s)+\frac{1}{2} O_2(g)+SO_2(g)\rightarrow CaSO_4(s)+CO_2(g)$$

### Catalytic converters

In 1993, it became law for all new cars in the UK to include catalytic converters fitted to their exhausts. This is because catalytic converters reduce the amounts of harmful carbon monoxide, nitrous oxides, and unburnt hydrocarbons in the exhaust fumes from vehicle internal combustion engines. Here's how they work:

• Catalytic converters have a honeycomb shape to maximise their surface area.
• They are coated in platinum and rhodium metals, which act as catalysts. These metals are spread thinly to minimise the amount needed.
• As the exhaust fume gases leave the engine, they pass over the catalytic converter and react to form less harmful products:
• Carbon monoxide (CO) reacts with nitrous oxides (NOx) to produce nitrogen (N2) and carbon dioxide (CO2).
• Unburnt hydrocarbons (CxHy) react with nitrous oxides to produce nitrogen, carbon dioxide (CO2), and water (H2O).

The exact reactions themselves depend on the types of nitrous oxides and unburnt hydrocarbons present in the flue gas. Here are a few examples of balanced chemical equations:

$$8CO(g)+4NO_2(g)\rightarrow 2N_2(g)+8CO_2(g)$$

$$C_9H_20(g)+28NO(g)\rightarrow 14N_2(g)+9CO_2(g)+10H_2O(g)$$

### Carbon-neutral alternatives

As you now know, burning fossil fuels such as coal and gas causes a net increase in atmospheric carbon dioxide levels. This is strongly linked to global warming. But what if we could find a fuel that didn't produce any overall carbon emissions and so didn’t heat our planet? Well, these fuels actually exist! We say that they are carbon-neutral.

Carbon-neutral substances are substances that produce no net overall carbon dioxide emissions in their lifetime. All the carbon dioxide they release is counterbalanced by carbon dioxide taken in at different points in their lives.

Examples of carbon-neutral fuels include biofuels and synthetic fuels:

• Biofuels are made from biomass, such as sugar cane or wood. They are quick to grow, and all the carbon they release upon combustion was taken in during their lifetime. However, the fertilisers, water, and land required to produce biofuels can create other problems like food security issues, monoculture, and deforestation.
• Synthetic fuels are artificial hydrocarbons. They are made from carbon dioxide captured from the air and hydrogen produced in electrolysis.

Carbon capture is still a relatively small industry but is growing rapidly as techniques improve. For example, the Swiss company Climeworks directly captures carbon from the air using special filters and either stores it deep within the ground or recycles it into fuels4.

## Combustion - Key takeaways

• Combustion is an exothermic reaction that involves burning a fuel such as coal, gas, or petrol, usually in oxygen.
• Combustion can be complete or incomplete.
• Complete combustion takes place in excess oxygen and produces carbon dioxide and water. It releases more energy than incomplete combustion.
• Incomplete combustion takes place in limited oxygen and produces carbon monoxide, carbon, and water. It releases less energy than complete combustion.
• The combustion of hydrocarbon fuels, such as coal, gas, and petrol, has several negative environmental effects.
• To reduce the environmental impact of combustion reactions, we can carry out flue gas desulphurisation, install catalytic converters, and switch to carbon-neutral fuels.

## References

1. K Kris Hirst, 'The Discovery of Fire'. ThoughtCo (03/05/2019)
2. 'World of Change: Global Temperature'. NASA Earth Observatory
3. 'World of Change: Global Temperature'. NASA Earth Observatory
4. 'Direct air capture: a technology to remove CO₂'. Climeworks

Combustion is an exothermic reaction that involves burning a fuel, usually in oxygen. It releases lots of energy in the form of heat and light.

Incomplete combustion is combustion that occurs in insufficient oxygen.

The products of complete combustion are carbon dioxide and water.

To stop combustion, you restrict oxygen. This can be done by covering the fire in sand or a heavy fire blanket, or dousing it in water.

Oxidising hydrogen in combustion forms water (H2O).

## Final Combustion Quiz

Question

What is combustion?

Combustion is an exothermic reaction that involves burning a fuel, usually in oxygen. It releases lots of energy in the form of heat and light.

Show question

Question

Compare complete and incomplete combustion.

• Both produce water and carbon dioxide.
• Both are exothermic reactions.
• Complete combustion takes place in excess oxygen.
• Incomplete combustion takes place in limited oxygen.
• Complete combustion releases more energy than incomplete combustion.

Show question

Question

What are the products of complete combustion?

Water

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Question

Give the two carbon-based products that can be produced in incomplete combustion.

Carbon particles

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Question

Why does using the safety flame of a Bunsen burner cause sooty deposits?

Using the safety flame encourages incomplete combustion. One of the products is carbon, which accumulates as soot.

Show question

Question

Write an equation to show the complete combustion of ethane.

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Question

Write an equation to show the incomplete combustion of ethane to produce carbon monoxide and water only.

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Question

Which of the following two alkanes releases more energy when burnt? Their enthalpies of combustion are given:

Butane.

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Question

Name four pollutants produced in combustion.

• Sulfur dioxide.
• Nitrous oxides.
• Carbon.
• Carbon monoxide.
• Carbon dioxide.

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Question

Hydrocarbon fuels can contain sulphur impurities. What are the negative effects caused by these impurities?

Sulphur burns to form sulphur dioxide, which can cause acid rain and breathing difficulties.

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Question

What are catalytic converters made of?

Platinum and rhodium.

Show question

Question

Write an equation to show the reaction of a nitrous oxide, NO, with an unburnt hydrocarbon C6H14, in a catalytic converter to form carbon dioxide, water, and nitrogen.

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Question

Discuss the pros and cons of biofuels.

• Biofuels are cheap, easy to grow and are carbon-neutral.
• However, they are relatively inefficient to burn.
• They take up valuable growing space that could be used for food crops.
• They may also require processing before use, which uses energy, and can be big and bulky to transport.

Show question

Question

How are nitrous oxides formed in combustion reactions?

The temperatures inside internal combustion engines are so high that oxygen and nitrogen from the air react to produce nitrous oxides.

Show question

Question

What are the negative effects of nitrous oxides released in combustion reactions?

• They cause acid rain.
• They contribute to photochemical smog.
• Furthermore, they cause breathing difficulties.

Show question

Question

Give the disadvantages of incomplete combustion.

• Incomplete combustion releases less energy than complete combustion.
• It produces carbon monoxide, which is a toxic gas.
• It also produces carbon particles, which cause global dimming and respiratory irritation.

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