Combustion (HSC SSCE Chemistry): Revision Notes

Combustion

Introduction to alcohol combustion

Alcohols are versatile organic compounds with numerous applications in both everyday life and industry. One increasingly important use is as a fuel additive. When you visit a petrol station, you may notice E10 fuel, which contains $10\%$ ethanol mixed with petrol.

The combustion of alcohols follows similar patterns to hydrocarbon fuels, releasing energy that can be harnessed for practical purposes.

Complete combustion of alcohols

When alcohols burn in sufficient oxygen, they undergo complete combustion. This reaction produces carbon dioxide gas, liquid water, and releases significant amounts of energy.

For ethanol, the complete combustion reaction is:

$\text{C}_2\text{H}_5\text{OH}(\text{l}) + 3\text{O}_2(\text{g}) \rightarrow 2\text{CO}_2(\text{g}) + 3\text{H}_2\text{O}(\text{l})$

This reaction releases approximately 1370 kJ per mole of ethanol burned. All combustion reactions are exothermic, meaning they release heat energy to the surrounding environment.

Incomplete combustion of alcohols

When insufficient oxygen is available, alcohols undergo incomplete combustion. This produces harmful products including carbon monoxide gas and solid carbon (soot), whilst releasing considerably less energy than complete combustion.

For ethanol, incomplete combustion follows this reaction:

$2\text{C}_2\text{H}_5\text{OH}(\text{l}) + 3\text{O}_2(\text{g}) \rightarrow 2\text{CO}(\text{g}) + 2\text{C}(\text{s}) + 6\text{H}_2\text{O}(\text{l})$

Understanding enthalpy of combustion

Enthalpy of combustion ($\Delta H$) measures the heat energy released when exactly one mole of fuel burns completely in excess oxygen at standard atmospheric pressure. This standardised measurement allows us to compare different fuels fairly.

Specific heat capacity and calorimetry

To measure enthalpy of combustion, we use an indirect method based on heating water. The specific heat capacity ($c$) of a substance tells us how much energy is needed to raise the temperature of one gram of that substance by one degree Celsius.

For water, the specific heat capacity is:

$c = 4.18 \text{ J K}^{-1} \text{ g}^{-1} \text{ or } 4180 \text{ J K}^{-1} \text{ kg}^{-1}$

The heat energy absorbed or released ($q$) can be calculated using:

$q = mc\Delta T$

Where:

• $m$ = mass of the substance being heated (in grams)
• $c$ = specific heat capacity of the substance ($\text{J K}^{-1} \text{ g}^{-1}$)
• $\Delta T$ = change in temperature (final temperature - initial temperature)

Calorimetry is the scientific process of measuring heat changes during chemical reactions.

Experimental measurement of enthalpy

Because directly measuring the heat from burning fuel is difficult, scientists use an indirect approach based on the law of conservation of energy. A known mass of water is heated by the burning fuel, and we measure the temperature change in the water.

The experimental setup typically includes:

• A spirit burner containing the alcohol fuel
• A conical flask or beaker holding a measured mass of water
• A thermometer to monitor temperature changes
• Clamps and stands to position the glassware safely above the flame

By calculating the heat absorbed by the water, we can determine the heat released by the fuel, since energy cannot be created or destroyed (conservation of energy principle).

Worked example: calculating enthalpy of combustion

Sources of experimental error

Enthalpy values measured using simple calorimetric experiments are typically very inaccurate, often producing results 80% lower than theoretical values. These are systematic errors that consistently affect results in the same direction.

Comparing energy values for different fuels

Energy per gram of fuel

In everyday situations, we don't purchase fuel by the mole. For liquid fuels like petrol, we buy litres; for solid fuels, we buy kilograms. Converting enthalpy values to these practical units allows meaningful comparisons.

To calculate heat released per gram:

$\text{Heat released per gram} = \frac{\Delta H \text{ (kJ mol}^{-1}\text{)}}{M \text{ (g mol}^{-1}\text{)}}$

Energy per litre of fuel

To convert to heat released per litre, we must also consider the fuel's density:

$\text{Heat released per litre} = \frac{\Delta H \text{ (kJ mol}^{-1}\text{)}}{M \text{ (g mol}^{-1}\text{)}} \times \text{density (g L}^{-1}\text{)}$

Comparison of different alcohols

The following table compares five different alcohol fuels, showing how energy values differ depending on the measurement units used:

Fuel	Enthalpy of combustion ($\text{kJ mol}^{-1}$)	Heat released per gram ($\text{kJ g}^{-1}$)	Heat released per litre ($\text{kJ L}^{-1}$)
Methanol	$-726$	$22.7$	$17\,978$
Ethanol	$-1360$	$29.7$	$23\,433$
1-propanol	$-2021$	$33.6$	$26\,980$
1-butanol	$-2676$	$36.1$	$29\,241$
1-pentanol	$-3329$	$37.7$	$30\,687$

Summary