Which of the following statements or relationships is not correct in thermodynamic changes?

Identify the correct statements from the following.
$I$. At $0 \ K$,the entropy of pure crystalline materials approaches zero.
$II$. Entropy for the process,$H_2O_{(l)} \longrightarrow H_2O_{(g)}$ decreases.
$III$. Gibbs' energy is a state function.

The enthalpy of formation of $H_2O$ is $-68 \ k.cal/mol$. Calculate the enthalpy of formation of $OH^-$. Given that the enthalpy of neutralization of $H^+$ and $OH^-$ is $-13.7 \ k.cal/mol$.

$2 \ mol$ of $Hg_{(g)}$ is combusted in a fixed volume bomb calorimeter with excess of $O_2$ at $298 \ K$ and $1 \ atm$ into $HgO_{(s)}$. During the reaction,temperature increases from $298.0 \ K$ to $312.8 \ K$. If heat capacity of the bomb calorimeter and enthalpy of formation of $Hg_{(g)}$ are $20.00 \ kJ \ K^{-1}$ and $61.32 \ kJ \ mol^{-1}$ at $298 \ K$,respectively,the calculated standard molar enthalpy of formation of $HgO_{(s)}$ at $298 \ K$ is $X \ kJ \ mol^{-1}$. The value of $|X|$ is. . . . . [Given : Gas constant $R = 8.3 \ J \ K^{-1} \ mol^{-1}$]

An organic compound $C_xH_yO_z$ on complete combustion provides equivolume of products $CO_{2(g)}$ and $H_2O_{(g)}$,which is individually double the volume of the organic compound taken. In the process,the volume of oxygen consumed is the same as the volume of $CO_2$ produced,and the liberated heat during combustion is $500 \ kcal/mol$ at constant pressure and $500 \ K$. The formula of the organic compound and $\Delta U$ for the reaction will be:

The $\Delta E^o$ of combustion of $2-$methylpropene is $-X \ kJ \ mol^{-1}$. The value of $\Delta H^o$ is