For the gas phase reaction A_{2(g)} + B_{2(g) | vedclass

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Question

(B) For an ideal gas reaction,the relationship between enthalpy change $(\Delta H)$ and internal energy change $(\Delta U)$ is given by $\Delta H = \D

Vedclass · Accepted Answer

$\Delta H 
eq \Delta U$

Vedclass · Answer

(B) For an ideal gas reaction,the relationship between enthalpy change $(\Delta H)$ and internal energy change $(\Delta U)$ is given by $\Delta H = \Delta U + \Delta n_g RT$. In the reaction $A_{2(g)} + B_{2(g)} ightarrow 2AB_{(g)}$,the change in the number of moles of gaseous species is $\Delta n_g = n_{products} - n_{reactants} = 2 - (1 + 1) = 0$. Since $\Delta n_g = 0$,the equation becomes $\Delta H = \Delta U + (0)RT$,which implies $\Delta H = \Delta U$. Therefore,option $B$ is incorrect as it states $\Delta H 
eq \Delta U$. Option $A$ and $C$ are not necessarily true for all such reactions. According to Kirchhoff's law,$\left[ \frac{\partial (\Delta_r H)}{\partial T} ight]_P = \Delta C_p$. For this specific reaction,$\Delta C_p = C_p(2AB) - [C_p(A_2) + C_p(B_2)]$. While $\Delta C_p$ is generally not zero,in the context of standard textbook problems regarding ideal gas reactions where $\Delta n_g = 0$,the relationship $\Delta H = \Delta U$ is the fundamental identity. However,looking at the options provided,if $\Delta n_g = 0$,then $\Delta H = \Delta U$ is the correct identity. Since the question asks what 'must' be correct and option $B$ says they are not equal,$B$ is false. Given the constraints,the question likely implies identifying the relationship between $\Delta H$ and $\Delta U$.

For the gas phase reaction $A_{2(g)} + B_{2(g)} \rightarrow 2AB_{(g)}$ taking place at constant pressure and temperature where all gases are behaving ideally,which of the following must be correct?

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