Calculate the enthalpy change for the reaction C2H4 + H2 -> C2H6 using the provided bond energies.

Bonds broken:

• $1 \times \text{C}=\text{C}$ bond in $\text{C}_2\text{H}_4 = 1 \times 610 \text{ kJ/mol} = 610 \text{ kJ/mol}$
• $4 \times \text{C}-\text{H}$ bonds in $\text{C}_2\text{H}_4 = 4 \times 410 \text{ kJ/mol} = 1640 \text{ kJ/mol}$
• $1 \times \text{H}-\text{H}$ bond in $\text{H}_2 = 1 \times 436 \text{ kJ/mol} = 436 \text{ kJ/mol}$ Total energy to break bonds = $610 + 1640 + 436 = 2686 \text{ kJ/mol}$

Bonds formed:

• $1 \times \text{C}-\text{C}$ bond in $\text{C}_2\text{H}_6 = 1 \times 150 \text{ kJ/mol} = 150 \text{ kJ/mol}$
• $6 \times \text{C}-\text{H}$ bonds in $\text{C}_2\text{H}_6 = 6 \times 410 \text{ kJ/mol} = 2460 \text{ kJ/mol}$ Total energy released when bonds form = $150 + 2460 = 2610 \text{ kJ/mol}$

B: The enthalpy change of this reaction was found by experiment to be $-136 \text{ kJ/mol}$. Explain why this value is different from that calculated above.
The calculated enthalpy change uses average bond energies, which are values averaged over many different compounds. The actual bond energies in specific molecules like ethene and ethane can differ from these average values. Additionally, bond energies are typically defined for substances in the gaseous state, and the experimental conditions might not perfectly match this.