Final answer:
The enthalpy change for the reaction H₂(g) + Cl₂(g) → 2 HCl(g) can be estimated using bond energies. The total energy required to break H-H and Cl-Cl bonds is 679 kJ/mol, and the energy released from forming two moles of H-Cl bonds is 864 kJ/mol, resulting in an enthalpy change of -185 kJ/mol, indicating an exothermic reaction.
Step-by-step explanation:
Estimating the Enthalpy Change for H₂ + Cl₂ → 2 HCl Reaction
To estimate the enthalpy change (ΔH) for the reaction H₂(g) + Cl₂(g) → 2 HCl(g), we can use the bond energies for the bonds being broken and formed during the reaction. The energy needed to break one mole of H-H bonds is 436 kJ/mol, and for the Cl-Cl bonds, it is 243 kJ/mol. When forming the products, two moles of H-Cl bonds are created, with each releasing 432 kJ/mol for a total of 864 kJ released. Thus, the total energy required to break the reactant bonds is the sum of 436 kJ/mol and 243 kJ/mol. Subtracting the energy released during bond formation in the products, we calculate the enthalpy change for the reaction as follows:
(436 kJ + 243 kJ) - (2 × 432 kJ) = ΔH
The enthalpy change (ΔH) for this reaction using bond energies would be:
ΔH = (436+243) - (864) = -185 kJ/mol
This indicates that the reaction is exothermic, releasing 185 kJ/mol of energy.