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Indicate whether the following reaction is exothermic or endothermic by properly computing the given data<br>$$\text{M}_g + 2\text{X}_g \rightarrow \text{M}^{2+}_g + 2\text{X}^-_g$$; $$\text{IE}_1$$ of $$\text{M}_g = 737.7\text{ kJ mol}^{-1}$$; $$\text{IE}_2$$ of $$\text{M}_g = 1451\text{ kJ mol}^{-1}$$; $$\text{EA}_1$$ of $$\text{X}_g = -328\text{ kJ mol}^{-1}$$
Correct Answer: 0
PLACEHOLDER ANSWER - the answer is a word with a calculation. Forming $$\text{M}^{2+}$$ needs $$737.7 + 1451 = 2188.7\text{ kJ mol}^{-1}$$, and forming two $$\text{X}^-$$ ions releases $$2 \times 328 = 656\text{ kJ mol}^{-1}$$. So $$\Delta H = 2188.7 - 656 = +1532.7\text{ kJ mol}^{-1}$$, which is positive, and the reaction is endothermic.
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