Consider the reaction \(\Lambda_{2}+\mathrm{B}_{2} \rightarrow 2 \Lambda \mathrm{B}\). Breaking \(1 \mathrm{molc}\) of \(\Lambda-\Lambda\) bonds and \(1 \mathrm{molc}\) of \(\mathrm{B}-\mathrm{B}\) bonds requires \(2200 \mathrm{~kJ}\). Forming 1 mole of \(\mathrm{A}-\mathrm{B}\) bonds releases \(1000 \mathrm{~kJ}\). (a) Is this reaction exothermic or endothermic? Explain. (b) What is the value of \(\Delta E_{\mathrm{rxn}} ?\) (Get the sign right.) How is it to be interpreted?

We are given that breaking 1 mole of \(\Lambda-\Lambda\) bonds and 1 mole of \(\mathrm{B}-\mathrm{B}\) bonds requires 2200 kJ. Forming 1 mole of \(\Lambda\mathrm{B}\) bonds releases 1000 kJ. Since 2 moles of \(\Lambda\mathrm{B}\) are formed in the reaction, the total energy released during bond formation is 2 x 1000 kJ = 2000 kJ. Now, compare the energy required for bond breaking and energy released during bond formation: - If the energy required for bond breaking is greater than the energy released during bond formation, the reaction will be endothermic (absorbing energy). - If the energy required for bond breaking is less than the energy released during bond formation, the reaction will be exothermic (releasing energy). In this case, 2200 kJ (energy required for bond breaking) is greater than 2000 kJ (energy released during bond formation), so the reaction is endothermic.

ΔErxn (change in energy during the reaction) can be calculated using the following formula: ΔErxn = energy released during bond formation - energy required for bond breaking ΔErxn = 2000 kJ - 2200 kJ = -200 kJ Since the value of ΔErxn is negative, it indicates that the reaction is endothermic. This means that the reaction absorbs 200 kJ of energy from its surroundings.