Quantifying energy changes in chemical reactions

The coefficients in a balanced equation represents the molar ratio of the amounts of reactants and products involved in the reaction. In thermochemical equations, the coefficients also indicate the number of moles of substances required to result in the enthalpy change (ΔH). If ΔH<0, it means the reaction releases energy (exothermic). If ΔH>0, the reaction absorbs energy (endothermic).

For example, the following balanced thermochemical equation tells us that 1 mole of CH4​ reacts with 2 moles of O2 to produce 1 mole of CO2​ and 2 moles of H2O, releasing 890 kJ of energy:

CH4​(g) + 2O2​(g) → CO2​(g) + 2H2​O(l)          ΔH = −890 kJ

Use this page to revise the following concepts within quantifying energy changes in chemical reactions:

Calculating energy released from fuel

Using the relationship shown in the balanced thermochemical equation, we can calculate the energy released when a given amount of fuel is consumed.

Worked Example

Calculate the heat energy released, in megajoules (MJ), when 12.0 kg of octane (C8H18) undergoes complete combustion. The balanced thermochemical equation for the reaction is:

2C8H18 (l) + 25O2 (g) →16CO2 (g) + 18H2O (l)          ΔH = −10 940 kJ

Step 1: Find moles of fuel.

\[n(C_8H_18)=\frac{m(C_8H_{18})}{M(C_8H_{18}})=\frac{12000}{114.0}=105mol\]

Step 2: Use the ratio between moles of fuel and energy, as given in the balanced thermochemical equation, to calculate the energy released.

According to the given balanced thermochemical equation, the complete combustion of 2mol C8H18 releases 10940 kJ energy

Let \(105\, mol\) of C8H18 release \(x\) kJ of energy

\[\frac{2}{105}=\frac{10940}{x}\]

\[x=5.75\times10^5kJ=575MJ\]

Therefore, 575 MJ of energy would be released.


Calculating fuel for desired amount of energy

Using the relationship shown in the balanced thermochemical equation, we can also calculate the amount of fuel required to release a certain amount of energy.

Worked Example

What volume of hydrogen gas (H2), measured at SLC, is required to produce 1.50 MJ of energy during complete combustion? The balanced thermochemical equation for the combustion of hydrogen gas is:

2H2(g) + O2(g) → 2H2O(l)          ΔH = −572 kJ

Solution

According to the balanced thermochemical equation:

When completely combusted, 2 mol of H2 produces 572 kJ of energy.

Assume it requires \(x\) mol of H2 to produce 1.50 MJ energy, which is 1.5 x 103 kJ

\[\frac{2}{x}=\frac{572}{1500}\]

\[x=5.24mol\]

The volume of the desired hydrogen gas can be calculated as \(V=n\times{V_m}=5.24\times24.8=130L\)