Answer
The change in energy (ΔE) for the given reaction is -1189 kJ/mol, which indicates that the reaction is exothermic (releases energy).
Work Step by Step
To predict the change in energy (ΔE) for the given reaction using bond energies, we need to calculate the energy released or absorbed during the breaking and formation of bonds.
The reaction is:
$H_{2}S(g) + 3F_{2}(g) \rightarrow SF_{4}(g) + 2HF(g)$
The bond energies involved are:
- H-H bond energy: 436 kJ/mol
- H-S bond energy: 339 kJ/mol
- F-F bond energy: 158 kJ/mol
- S-F bond energy: 327 kJ/mol
- H-F bond energy: 565 kJ/mol
To calculate the change in energy (ΔE), we need to consider the energy required to break the reactant bonds and the energy released by the formation of the product bonds.
Energy required to break the reactant bonds:
- 1 H-H bond (436 kJ/mol)
- 1 H-S bond (339 kJ/mol)
- 3 F-F bonds (3 × 158 kJ/mol = 474 kJ/mol)
Total energy required to break the reactant bonds = 436 + 339 + 474 = 1249 kJ/mol
Energy released by the formation of product bonds:
- 4 S-F bonds (4 × 327 kJ/mol = 1308 kJ/mol)
- 2 H-F bonds (2 × 565 kJ/mol = 1130 kJ/mol)
Total energy released by the formation of product bonds = 1308 + 1130 = 2438 kJ/mol
The change in energy (ΔE) is the difference between the energy released and the energy required:
ΔE = Energy released - Energy required
ΔE = 2438 kJ/mol - 1249 kJ/mol = 1189 kJ/mol
Therefore, the change in energy (ΔE) for the given reaction is -1189 kJ/mol, which indicates that the reaction is exothermic (releases energy).
