Relation between ΔH and ΔE for the reaction in gaseous phase at constant temperature and pressure Which of the following statements is correct for the reaction at constant temperature and pressure:
$$\ce{CO(g) + \frac{1}{2}O_2(g) -&gt; CO2(g)}$$

$\Delta H = \Delta E$
$\Delta H &gt; \Delta E$
$\Delta H &lt; \Delta E$
None of these.


Since this occurs at constant temperature, $\Delta E = 0$ and $W &lt; 0$. So, $\Delta H$ turns out to be less than $\Delta E$, but this is not the answer. Why?
By definition, $$\Delta H=\Delta E+\Delta(PV)$$For a reacting ideal gas mixture at constant temperature, $$\Delta(PV)=(\Delta n)RT$$where $\Delta n$ is the change in the number of moles between reactants and products.  Therefore, $$\Delta H=\Delta E+(\Delta n)RT$$For the reaction under consideration, $$\Delta n=-\frac{1}{2}$$