Organic Fuels and Biofuels: HSC Chemistry Comparison Guide

Two fuels can produce almost the same amount of carbon dioxide when they burn, yet one may add much less fossil carbon to the atmosphere over its full life cycle. At the same time, that “greener” fuel might make a car travel a shorter distance on the same volume of fuel.

So which fuel is better?

Before reading on, make a prediction. Suppose one car burns petrol and another burns ethanol made from recently grown sugarcane. If both cars release carbon dioxide from the exhaust, why might the ethanol still have an environmental advantage?

The key is that choosing a fuel is not just a combustion question. You need to track three things at once: the energy the fuel provides, where its carbon came from, and what had to happen before the fuel reached the engine.

01Fossil fuel and biofuel are not opposites in chemistry

There is an easy terminology trap here.

In chemistry, an organic substance is mainly a carbon-based substance. Petrol contains organic compounds. Ethanol is also an organic compound. So a biofuel such as ethanol is itself an organic fuel.

The more useful comparison is between:

Both ultimately have biological origins. The important difference is the timescale.

Burning a fossil fuel moves carbon from a long-term geological store into the active carbon cycle. Burning a biofuel also releases carbon dioxide, but the plants used to make the fuel may have recently removed carbon dioxide from the atmosphere while growing.

That is why “it produces \(\ce{CO2}\)” is not enough information to compare the environmental impact.

02What a fuel is actually doing

A fuel stores chemical energy. During combustion, it reacts with oxygen and forms more stable products, usually carbon dioxide and water.

For ethanol, complete combustion is:

\[ \ce{C2H5OH(l) + 3O2(g) -> 2CO2(g) + 3H2O(l)} \]

One mole of ethanol therefore produces two moles of carbon dioxide during complete combustion.

A simplified model for petrol often uses octane, \(\ce{C8H18}\). Real petrol is a mixture, so octane is a useful model rather than an exact description.

Complete combustion of octane can be written as:

\[ \ce{2C8H18(l) + 25O2(g) -> 16CO2(g) + 18H2O(l)} \]

At this stage you might predict that octane must be environmentally much worse because one octane molecule contains eight carbon atoms while ethanol contains only two.

That comparison is incomplete. An octane molecule is also much larger and releases much more energy per mole. To compare fuels fairly, we often need to ask how much carbon dioxide is produced for the same amount of energy delivered.

03Comparing the energy available from fuels

The enthalpy of combustion, \(\Delta H_c\), is the enthalpy change when one mole of a substance undergoes complete combustion under specified conditions.

Combustion is exothermic, so \(\Delta H_c\) is negative. A more negative value means more energy is released per mole.

For example, suppose we use:

• ethanol: \(\Delta H_c = -1367\ \text{kJ mol}^{-1}\)
• octane: \(\Delta H_c = -5470\ \text{kJ mol}^{-1}\)

At first glance, octane seems to provide about four times as much energy. But that is per mole. One mole of octane is much heavier than one mole of ethanol.

For practical fuel comparisons, energy per gram or energy per litre is often more useful.

04Why ethanol contains less energy per gram

Ethanol is:

\[ \ce{C2H5OH} \]

A hydrocarbon such as octane contains only carbon and hydrogen:

\[ \ce{C8H18} \]

Ethanol already contains oxygen. In chemical terms, its carbon atoms are partly oxidised compared with the carbon in a hydrocarbon.

Picture the combustion process as an energy downhill run towards \(\ce{CO2}\) and \(\ce{H2O}\). A hydrocarbon begins higher up the hill. Ethanol has already travelled partway down because oxygen is already incorporated into the molecule.

The analogy is useful for comparing chemical energy, but it is not a literal picture of individual bonds “falling downhill”. Enthalpy depends on the total energy difference between reactants and products.

The practical result is the important part: ethanol has a lower energy content per unit mass than common hydrocarbon fuels such as petrol.

05Tailpipe carbon dioxide gives only part of the answer

Now return to the original puzzle.

Could ethanol and a hydrocarbon fuel produce similar masses of carbon dioxide when they supply the same amount of energy?

Surprisingly, yes.

06Fossil carbon and recently captured carbon

Imagine the atmosphere has a carbon bank account.

Growing biomass withdraws carbon dioxide from the atmosphere through photosynthesis. Burning fuel made from that biomass returns some of that carbon.

A fossil fuel does something different. Its carbon has been locked away underground on geological timescales. Burning it transfers that previously stored carbon into the atmosphere.

The bank-account analogy breaks if we pretend every deposit and withdrawal is automatic or equal. Growing, harvesting, transporting, processing, and refining biomass all require resources and may produce additional greenhouse gas emissions.

That is why biofuel should not automatically be described as “carbon neutral”.

07How ethanol can be produced as a biofuel

One common route begins with carbohydrates obtained from biomass.

Sugars can undergo fermentation:

\[ \ce{C6H12O6(aq) -> 2C2H5OH(aq) + 2CO2(g)} \]

Glucose produces ethanol and carbon dioxide.

The ethanol then needs to be separated from the fermentation mixture. Distillation requires energy because the mixture must be heated and ethanol must be separated from a large amount of water.

For fuel applications, further water removal may also be required.

This creates an important environmental question: where does the energy for processing come from?

If a bioethanol plant burns a large amount of fossil fuel to grow crops, manufacture fertiliser, transport biomass, and distil ethanol, its overall greenhouse impact can be much less favourable than the simple plant-growth diagram suggests.

08A useful way to evaluate any fuel

An HSC response comparing fuels should avoid one-word judgements such as “renewable” or “clean”.

Instead, separate the comparison into energy, environmental, and practical criteria.

The phrase “potentially renewable” matters. If forests are cleared faster than biomass regrows, calling the resulting fuel renewable does not make the environmental damage disappear.

09The most tempting misconception: biofuel means carbon neutral

A student might reason like this:

Plants absorb \(\ce{CO2}\). We turn the plants into ethanol. Burning the ethanol releases that same \(\ce{CO2}\). Therefore the net emissions are zero.

That is a useful first model, but it leaves out several parts of the system.

A fuller life-cycle assessment may include:

• energy used to plant, harvest, and process the feedstock
• fuel used to transport biomass and ethanol
• energy required for fermentation products to be separated and purified
• fertiliser production and agricultural emissions
• changes in soil carbon
• clearing of existing vegetation
• water use
• whether the crop replaces food production or another land use
• what happens to by-products and waste.

A biofuel can therefore have lower life-cycle greenhouse emissions than a fossil fuel without having zero emissions.

It can also perform poorly if its production pathway is inefficient or causes major land-use change.

The word “bio” tells you the source. It does not give you the final environmental verdict.

10Complete combustion is not guaranteed either

The combustion equations above assume complete combustion.

That requires enough oxygen for carbon to form \(\ce{CO2}\) and hydrogen to form \(\ce{H2O}\).

With insufficient oxygen, incomplete combustion can produce carbon monoxide and carbon-containing particulates.

For example, ethanol may form carbon monoxide:

\[ \ce{C2H5OH + 2O2 -> 2CO + 3H2O} \]

Real engines are more complicated than a single equation. Temperature, oxygen supply, engine design, and combustion conditions all affect the products.

So another common shortcut fails: simply knowing that a fuel contains oxygen does not guarantee perfectly complete or pollution-free combustion.

11Why ethanol’s physical properties matter

A fuel has to do more than release energy. It has to survive storage, transport, and use.

Ethanol contains a polar \(\ce{-OH}\) group and can form hydrogen bonds. That makes its physical behaviour quite different from non-polar hydrocarbons in petrol.

In particular, ethanol is highly miscible with water. Water contamination can therefore matter during storage and fuel handling.

Its intermolecular forces also affect properties such as boiling behaviour, which becomes important when ethanol is separated from fermentation mixtures. If that connection feels unfamiliar, the intermolecular bonding in alcohols guide develops it from the molecular level.

There is no point finding a chemically impressive fuel if storing, transporting, separating, or using it requires unreasonable amounts of energy or unsuitable infrastructure.

12A better decision rule than “renewable versus non-renewable”

When comparing fuels, work through this sequence.

1. How much useful energy can the fuel provide per mass or volume?
2. What products form when it burns?
3. Where did its carbon come from?
4. What energy and resources were required to produce the fuel?
5. Can its feedstock be replaced sustainably?
6. What practical problems arise in storage, transport, and use?

Notice that no single criterion decides the answer.

Petrol has an energy-density advantage. Bioethanol may have a carbon-cycle advantage. A particular ethanol production pathway may then gain or lose environmental advantages depending on farming, processing, and land use.

That is the level of comparison an HSC question usually needs.

13Questions and solutions

Question 1

A sample containing \(2.00\ \text{mol}\) of ethanol undergoes complete combustion.

Use \(\Delta H_c(\text{ethanol})=-1367\ \text{kJ mol}^{-1}\).

Calculate:

a. the energy released
b. the amount of carbon dioxide produced
c. the mass of carbon dioxide produced.

Use \(M(\ce{CO2})=44.01\ \text{g mol}^{-1}\).

Solution 1

Complete combustion of \(2.00\ \text{mol}\) of ethanol releases \(2.73\times10^3\ \text{kJ}\) and produces \(176\ \text{g}\) of carbon dioxide.

The combustion equation is:

\[ \ce{C2H5OH + 3O2 -> 2CO2 + 3H2O} \]

For part a, the magnitude of the energy released is:

\[ E = n|\Delta H_c| = 2.00\ \text{mol}\times1367\ \text{kJ mol}^{-1} = 2734\ \text{kJ} \]

Therefore:

\[ E=2.73\times10^3\ \text{kJ} \]

For part b, the mole ratio is:

\[ 1\ \text{mol ethanol}:2\ \text{mol }\ce{CO2} \]

Therefore:

\[ n(\ce{CO2})=2(2.00)=4.00\ \text{mol} \]

For part c:

\[ m=nM = 4.00\ \text{mol}\times44.01\ \text{g mol}^{-1} = 176.04\ \text{g} \]

So the mass is approximately:

\[ 176\ \text{g} \]

A tempting mistake is to say bioethanol produces no carbon dioxide because plants absorbed carbon dioxide while growing. That is a life-cycle argument, not a combustion result. The chemical equation shows that ethanol itself definitely produces \(\ce{CO2}\) when completely burned.

Question 2

A simplified petrol model uses octane with \(\Delta H_c=-5470\ \text{kJ mol}^{-1}\) and molar mass \(114.2\ \text{g mol}^{-1}\).

Ethanol has \(\Delta H_c=-1367\ \text{kJ mol}^{-1}\) and molar mass \(46.1\ \text{g mol}^{-1}\).

A vehicle needs \(30.0\ \text{MJ}\) of chemical energy.

Calculate the mass of each fuel required. Hence determine how many times greater the required mass of ethanol is.

Assume equal engine efficiency.

Solution 2

The vehicle requires about \(1.01\ \text{kg}\) of ethanol or \(0.626\ \text{kg}\) of octane, so the required ethanol mass is about \(1.62\) times greater.

For ethanol:

\[ n = \frac{30\,000\ \text{kJ}}{1367\ \text{kJ mol}^{-1}} = 21.9\ \text{mol} \]

Then:

\[ m = nM = 21.9\ \text{mol}\times46.1\ \text{g mol}^{-1} \approx1010\ \text{g} \]

So:

\[ m(\text{ethanol})\approx1.01\ \text{kg} \]

For octane:

\[ n = \frac{30\,000\ \text{kJ}}{5470\ \text{kJ mol}^{-1}} = 5.48\ \text{mol} \]

Then:

\[ m = 5.48\ \text{mol}\times114.2\ \text{g mol}^{-1} \approx626\ \text{g} \]

So:

\[ m(\text{octane})\approx0.626\ \text{kg} \]

The ratio is:

\[ \frac{1.01}{0.626}\approx1.62 \]

The tempting route is to compare only the molar enthalpies and claim octane provides roughly four times the energy. That ignores the much greater mass of one mole of octane. For a practical mass comparison, energy per gram is the relevant quantity.

Question 3

A student makes the following argument:

Bioethanol is always environmentally better than petrol because ethanol is renewable and petrol is non-renewable.

Evaluate this statement. Your response must consider combustion, the carbon cycle, production, and one practical property of ethanol.

Solution 3

The statement is too strong: bioethanol can offer environmental advantages over petrol, but being renewable does not make it automatically better in every situation.

Both fuels produce carbon dioxide during complete combustion. For ethanol:

\[ \ce{C2H5OH + 3O2 -> 2CO2 + 3H2O} \]

The important difference is the carbon source. Petrol transfers carbon from long-term geological stores into the atmosphere. Ethanol made from recently grown biomass may recycle carbon that plants recently removed from atmospheric \(\ce{CO2}\).

However, the life-cycle impact of ethanol also depends on cultivation, fertiliser use, harvesting, transport, fermentation, and energy-intensive separation processes such as distillation. Clearing high-carbon vegetation to grow a fuel crop can also create emissions that reduce or reverse the expected benefit.

Ethanol additionally has lower energy content per unit mass than petrol-like hydrocarbons, so more fuel may be required to provide the same chemical energy.

A practical difference is ethanol’s interaction with water. Because ethanol contains a polar \(\ce{-OH}\) group and forms hydrogen bonds, it mixes readily with water. This can create storage and fuel-handling issues that differ from those of non-polar hydrocarbon fuels.

The tempting mistake is to treat “renewable” as a complete environmental assessment. Renewability describes whether the resource can be replenished on a useful timescale. It does not measure every emission, resource input, or practical consequence.

Question 4

One tonne of dry waste biomass contains \(18.0\ \text{GJ}\) of chemical energy. A council is considering two uses for it.

Option A: Burn the biomass in a heating system that converts \(85\%\) of its chemical energy into useful heat. This useful heat replaces heat from a natural-gas system that is \(90\%\) efficient. Burning natural gas releases \(56.0\ \text{kg}\) of fossil \(\ce{CO2}\) per gigajoule of natural gas burned.

Option B: Convert the biomass into ethanol. The ethanol contains \(55\%\) of the original biomass energy. It replaces an equal amount of petrol chemical energy in vehicles. Petrol releases \(73.0\ \text{kg}\) of fossil \(\ce{CO2}\) per gigajoule burned. Fossil energy used during ethanol manufacture causes an additional \(150\ \text{kg}\) of \(\ce{CO2}\) emissions per tonne of biomass processed.

Assume the biomass itself is sustainably regenerated and ignore all other emissions.

Which option avoids more fossil carbon dioxide emissions? Calculate the difference and explain why simply choosing the transport biofuel is not automatically the better use of the biomass.

Solution 4

Option A avoids more fossil carbon dioxide: about \(952\ \text{kg}\), compared with about \(573\ \text{kg}\) for Option B, a difference of approximately \(379\ \text{kg}\) of \(\ce{CO2}\) per tonne of biomass.

For Option A, useful heat obtained from the biomass is:

\[ E_{\text{useful}} = 18.0\ \text{GJ}\times0.85 = 15.3\ \text{GJ} \]

A natural-gas system with \(90\%\) efficiency would need:

\[ E_{\text{gas}} = \frac{15.3\ \text{GJ}}{0.90} = 17.0\ \text{GJ} \]

The avoided fossil carbon dioxide is therefore:

\[ m(\ce{CO2}) = 17.0\ \text{GJ}\times56.0\ \text{kg GJ}^{-1} = 952\ \text{kg} \]

For Option B, the ethanol contains:

\[ E_{\text{ethanol}} = 18.0\ \text{GJ}\times0.55 = 9.90\ \text{GJ} \]

It therefore replaces \(9.90\ \text{GJ}\) of petrol.

The petrol emissions avoided are:

\[ m(\ce{CO2}) = 9.90\ \text{GJ}\times73.0\ \text{kg GJ}^{-1} = 723\ \text{kg} \]

But ethanol manufacture causes an additional \(150\ \text{kg}\) of fossil carbon dioxide emissions.

The net avoided emissions are:

\[ 723-150=573\ \text{kg} \]

Therefore:

\[ 952-573=379\ \text{kg} \]

Option A avoids about \(379\ \text{kg}\) more fossil \(\ce{CO2}\).

The tempting answer is Option B because “biofuel replacing petrol” sounds like the more environmentally advanced use. The missing constraint is that biomass is limited. The correct comparison is therefore not just “biofuel versus fossil fuel”, but “what alternative use of this biomass displaces the greatest amount of fossil fuel?”

Conversion losses matter. Option B retains only \(55\%\) of the biomass energy in the ethanol, and its processing also creates fossil emissions. Option A uses more of the original energy to displace another fossil fuel.

This does not prove that biomass should always be burned directly. Change the efficiencies, displaced fuels, processing emissions, or required form of energy, and the result can change.

Question 5

Consider an idealised bioethanol cycle beginning with glucose produced by photosynthesis.

Suppose formation of one mole of glucose from atmospheric carbon dioxide can be represented by:

\[ \ce{6CO2 + 6H2O -> C6H12O6 + 6O2} \]

Fermentation is:

\[ \ce{C6H12O6 -> 2C2H5OH + 2CO2} \]

Combustion is:

\[ \ce{C2H5OH + 3O2 -> 2CO2 + 3H2O} \]

A plant converts one mole of glucose into ethanol and burns both moles of ethanol as fuel.

a. Track the atmospheric carbon dioxide absorbed and released over the complete idealised cycle.

b. A student says, “Each mole of ethanol produces two moles of \(\ce{CO2}\), so the crop only needs to absorb two moles of \(\ce{CO2}\) per mole of ethanol for the fuel to be carbon neutral.” Explain the flaw.

c. Now suppose all fermentation \(\ce{CO2}\) is captured and permanently stored instead of released. Fossil energy used to grow and process enough biomass for two moles of ethanol produces \(0.80\ \text{mol}\) of additional \(\ce{CO2}\). Calculate the net change in atmospheric \(\ce{CO2}\) for the full cycle, assuming the biomass regrows exactly as stated.

Solution 5

The ordinary idealised cycle has zero net change in atmospheric \(\ce{CO2}\); if the fermentation carbon dioxide is permanently stored and the stated processing emissions are included, the cycle removes a net \(1.20\ \text{mol}\) of \(\ce{CO2}\) from the atmosphere.

For part a, photosynthesis absorbs:

\[ 6\ \text{mol }\ce{CO2} \]

Fermentation then releases:

\[ 2\ \text{mol }\ce{CO2} \]

The one mole of glucose produces two moles of ethanol. Each mole of ethanol produces two moles of carbon dioxide during combustion, so combustion releases:

\[ 2\times2=4\ \text{mol }\ce{CO2} \]

Total carbon dioxide released is therefore:

\[ 2+4=6\ \text{mol }\ce{CO2} \]

The idealised net change is:

\[ 6\ \text{mol absorbed}-6\ \text{mol released}=0 \]

For part b, the student’s conclusion accidentally ignores the carbon dioxide released during fermentation.

For every mole of ethanol ultimately produced, the original glucose contained three carbon atoms’ worth of carbon. Two carbon atoms remain in the ethanol and are later released as two molecules of \(\ce{CO2}\) during combustion. The remaining carbon atom is released as \(\ce{CO2}\) during fermentation.

So, averaged per mole of ethanol, photosynthesis originally removed three moles of atmospheric carbon dioxide, not two:

\[ 3\ \text{mol absorbed} = 1\ \text{mol released during fermentation} + 2\ \text{mol released during combustion} \]

This is the hinge in the problem. Looking only at the exhaust loses track of carbon that left the biomass earlier in the production pathway.

For part c, the crop still removes:

\[ 6\ \text{mol }\ce{CO2} \]

The \(2\ \text{mol}\) of fermentation carbon dioxide is now captured and permanently stored, so it does not return to the atmosphere.

Combustion releases:

\[ 4\ \text{mol }\ce{CO2} \]

Processing adds another:

\[ 0.80\ \text{mol }\ce{CO2} \]

Total atmospheric release is therefore:

\[ 4+0.80=4.80\ \text{mol }\ce{CO2} \]

The net atmospheric change is:

\[ 4.80-6.00=-1.20\ \text{mol }\ce{CO2} \]

So the atmosphere loses a net:

\[ 1.20\ \text{mol }\ce{CO2} \]

under the assumptions given.

A strong student may initially expect the best possible result to be “carbon neutral”, because the carbon in a biofuel originally came from the atmosphere. The important interaction is that carbon can leave the cycle at more than one point. If part of the recently captured biomass carbon is permanently stored while the remaining biomass regrows, the idealised system can remove atmospheric carbon overall.

That conclusion depends completely on the assumptions. If the biomass is not regrown, storage is not permanent, land clearing releases additional carbon, or processing emissions are sufficiently large, the result changes.

The broader lesson is that the fuel molecule alone cannot tell you its environmental impact. Combustion chemistry tells you what happens inside the reaction. A life-cycle comparison tracks where the reactants came from and what happened before and after that reaction.

The next useful step is to connect this comparison to measured energy changes. The enthalpy of combustion of alcohols guide shows how \(\Delta H_c\) is determined and why experimental values often differ from accepted values.