A series of stacked-bar charts gives the average bond-energy contributions (in kJ mol⁻¹) for the molecules involved in the chlorination of methane:
CH₄(g): one C–H bond shown at 413 (kJ mol⁻¹) repeated four times in the bar
Cl₂(g): one Cl–Cl bond at 242
CH₃Cl(g): three C–H bonds at 413 each and one C–Cl bond at 328
HCl(g): one H–Cl bond at 431
Consider the reaction:
CH₄(g) + Cl₂(g) → CH₃Cl(g) + HCl(g)
(a) Using the graphical information, determine the total energy absorbed to break all bonds in the reactants and the total energy released when all bonds form in the products.
(b) Calculate the enthalpy change, ΔH, for the reaction. State whether the process is endothermic or exothermic.
(c) When 1.00 g of CH₄ reacts completely under the conditions stated, estimate the quantity of heat absorbed or released (in kJ). Assume the value of ΔH you obtained in part (b) applies per mole of reaction as written.