Comparing Fuels

Introduction

Barbecues burn charcoal, cars burn petrol, rockets can burn hydrogen and camping stoves burn gas. Why so many different fuels? No single fuel is best for every job. A good fuel for a spacecraft must pack enormous energy into a small mass; a good fuel for a home must be safe, cheap and clean. To choose sensibly, chemists compare fuels on several measurable properties, starting with how much energy each gram releases.

Core explanation

Energy per gram. The most direct comparison is the energy released when 1 g of fuel burns completely. Typical values:

Fuel Formula or type Energy released (kJ/g) --- --- --- Hydrogen H₂ about 142 Methane (natural gas) CH₄ about 55 Propane (bottled gas) C₃H₈ about 50 Petrol (mainly octane) about C₈H₁₈ about 47 Ethanol C₂H₅OH about 30 Coal mostly carbon about 25–30 Dry wood cellulose and lignin about 15–18

Hydrogen stands out because hydrogen atoms are extremely light. Among hydrocarbons, fuels with a higher proportion of hydrogen tend to release more energy per gram. Ethanol releases less than the alkanes because it already contains an oxygen atom — it is partly oxidised before it burns.

Energy per unit volume matters too. Hydrogen gas has a very low density, so a litre of it holds little energy unless it is compressed or liquefied. Liquid petrol, by contrast, stores a lot of energy in a small tank. This is one reason liquid fuels dominate road transport.

Flame colour and cleanliness. Gaseous fuels such as methane mix easily with air and usually burn with a clean blue flame. Heavier liquid and solid fuels are harder to mix with oxygen and more often give yellow, smoky flames, producing soot and carbon monoxide. Coal and heavy oils may also contain sulfur and release sulfur dioxide. Hydrogen produces only water when it burns.

Measuring energy in the laboratory. A simple method, called calorimetry, uses a fuel to heat a known mass of water. The energy transferred to the water is:

Q = m × c × ΔT

where m is the mass of water in g, c is 4.18 J/g °C and ΔT is the temperature rise. Dividing Q by the mass of fuel burnt gives energy per gram. Simple apparatus loses a lot of heat to the air, so measured values are usually well below textbook values.

Other factors. A practical choice also considers cost, availability, ease of storage and transport, how easily the fuel ignites, the risk of leaks or explosions, and how much carbon dioxide it releases per unit of energy.

Formulae

Q = m × c × ΔT (Q in J, m in g, c = 4.18 J/g °C for water, ΔT in °C). Energy per gram of fuel = Q ÷ mass of fuel burnt.

Step-by-step reasoning

To compare two fuels fairly by experiment:

1. Use the same mass of water and the same container each time. 2. Keep the same distance between flame and container. 3. Measure the temperature rise and the mass of fuel burnt. 4. Calculate Q = mcΔT for each fuel. 5. Divide by the mass of fuel to get energy per gram, then compare.

Visual explanation

Picture a bar chart of energy per gram. The hydrogen bar towers above all others at about 142 kJ/g. Methane, propane and petrol form a cluster around 50, ethanol sits near 30, and coal and wood are shortest. A second chart of energy per litre would look very different, with liquid fuels at the top.

Real-world analogy

Choosing a fuel is like choosing a snack for a hike. A chocolate bar packs lots of energy into little weight, but melts; fruit is refreshing but bulky for its energy. The best choice depends on the trip, just as the best fuel depends on the task.

Real-world example

Space rockets such as those used for upper stages often burn liquid hydrogen with liquid oxygen, because every kilogram lifted is costly and hydrogen gives the most energy per gram. Home heating in many countries uses natural gas instead, because it is piped directly to buildings, burns cleanly and is relatively cheap.

Why?

Why does ethanol release less energy per gram than octane? Ethanol already has an oxygen atom bonded to carbon, so part of its oxidation has already happened. Octane contains only carbon and hydrogen, so all of its atoms still have to form strong bonds with oxygen, releasing more energy.

Common misconception

"The fuel that gives the biggest flame releases the most energy." Flame size depends on burning rate and wick or nozzle size. Energy per gram is found only by measuring energy released against mass of fuel used.

Worked example

Question: Burning 0.50 g of a liquid fuel raises the temperature of 100 g of water by 25 °C. Calculate the energy transferred per gram.

Reasoning: Q = 100 × 4.18 × 25 = 10 450 J = 10.45 kJ. Energy per gram = 10.45 ÷ 0.50 = 20.9 kJ/g.

Answer: About 21 kJ/g. This is lower than the true value because heat escapes to the surroundings.

Quick check

1. Which common fuel releases the most energy per gram? Answer: Hydrogen, at about 142 kJ/g.

Exam focus

Practise Q = mcΔT calculations and remember to convert J to kJ. Know how to control variables when comparing fuels, and give heat loss to the surroundings as the main reason experimental values are too low. Evaluate fuels using several factors, not energy alone.

Advanced insight

Values of energy per gram can be quoted as "higher" or "lower" heating values depending on whether the water produced is counted as liquid or vapour. Condensing boilers recover extra energy by cooling exhaust gases until the water vapour condenses, making them noticeably more efficient.

Summary

Fuels are compared by energy released per gram, energy per unit volume, cleanliness of burning, cost, storage and safety. Hydrogen gives the most energy per gram; alkanes give around 50 kJ/g; oxygen-containing fuels such as ethanol give less. Q = mcΔT lets us measure energy transferred to water, though simple methods lose heat to the surroundings.

Practice questions

1. Calculate the energy needed to raise 250 g of water by 10 °C. Answer: Q = 250 × 4.18 × 10 = 10 450 J, about 10.5 kJ. 2. Give two reasons, other than energy per gram, for choosing a fuel. Answer: For example, it burns cleanly without soot, and it is cheap and easy to store and transport. 3. Why do experimental values of energy per gram tend to be lower than data-book values? Answer: Much of the heat is lost to the air and apparatus instead of the water, and combustion may be incomplete. 4. Why is hydrogen difficult to use as a car fuel despite its high energy per gram? Answer: As a gas it has a very low density, so it must be stored under high pressure or as a very cold liquid to fit enough in a tank. 5. Explain why a gaseous fuel usually burns more cleanly than a solid one. Answer: A gas mixes thoroughly with air, so each molecule meets enough oxygen for complete combustion, giving less soot and carbon monoxide.