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(C) $1$. The standard enthalpy of formation for alkali metal bromides becomes more negative on descending the group due to the increase in the size of

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Among the alkali metal halides,$LiF$ is least soluble in water.

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(C) $1$. The standard enthalpy of formation for alkali metal bromides becomes more negative on descending the group due to the increase in the size of the metal cation and the decrease in lattice energy being offset by the ionization energy trends.$2$. The low solubility of $CsI$ is not due to high lattice enthalpy,but rather due to the low hydration enthalpy of the large $Cs^+$ and $I^-$ ions.$3$. Among the alkali metal halides,$LiF$ is the least soluble in water because of its very high lattice enthalpy compared to its hydration enthalpy.$4$. The standard enthalpy of formation for $LiF$ is the most negative among alkali metal fluorides because of the very high lattice energy associated with the small $Li^+$ and $F^-$ ions.

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