Consider the reversible processes for $1.0 \ mol$ of an ideal gas. $w_1, w_2, w_3$ and $w_4$ represent work done (in calories) in the processes $1, 2, 3$ and $4$, respectively; $\Delta U_2$ and $\Delta U_4$ are changes in the internal energy for the processes $2$ and $4$, respectively. (use $R = 2 \ cal \ K^{-1} \ mol^{-1}$). The correct option is:
- A$w_1 + w_3 = -2T_1 \ln \frac{V_2}{V_1} - 2T_2 \ln \frac{V_4}{V_3}$
- B$w_2 + w_4 = \Delta U_2 - \Delta U_4$
- C$w_1 + w_2 = 2T_1 \ln \frac{V_2}{V_1}$
- D$w_1 + w_2 + w_3 + w_4 = 0$
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For the reaction $2X_{(s)} + 2Y_{(s)} \to 2C_{(\ell)} + D_{(g)}$,the value of $q_p$ is $-28 \, \text{kcal mol}^{-1}$ at $2700 \, ^\circ\text{C}$. Calculate the value of $q_v$ in $\text{kcal mol}^{-1}$.
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View SolutionFor the complete combustion of methanol
$CH_3OH_{(l)} + \frac{3}{2} O_{2(g)} \rightarrow CO_{2(g)} + 2H_2O_{(l)}$
the amount of heat produced as measured by a bomb calorimeter is $726 \ kJ \ mol^{-1}$ at $27^{\circ}C$. The enthalpy of combustion for the reaction is $-x \ kJ \ mol^{-1}$,where $x$ is $.....$ (Nearest integer).
(Given: $R = 8.3 \ J \ K^{-1} \ mol^{-1}$)
$CH_3OH_{(l)} + \frac{3}{2} O_{2(g)} \rightarrow CO_{2(g)} + 2H_2O_{(l)}$
the amount of heat produced as measured by a bomb calorimeter is $726 \ kJ \ mol^{-1}$ at $27^{\circ}C$. The enthalpy of combustion for the reaction is $-x \ kJ \ mol^{-1}$,where $x$ is $.....$ (Nearest integer).
(Given: $R = 8.3 \ J \ K^{-1} \ mol^{-1}$)
For the reaction,$2 CO + O_2 \longrightarrow 2 CO_2$; $\Delta H = -560 \ kJ$. Two moles of $CO$ and one mole of $O_2$ are taken in a container of volume $1 \ L$. They completely form two moles of $CO_2$. The gases deviate appreciably from ideal behavior. If the pressure in the vessel changes from $70 \ atm$ to $40 \ atm$,find the magnitude (absolute value) of $\Delta U$ at $500 \ K$. $(1 \ L \ atm = 0.1 \ kJ)$
If the standard enthalpy change $\left(\Delta_{r} H^{\theta}\right)$ for a certain reaction at $298 \ K$ and constant pressure is $-1860 \ kJ \ mol^{-1}$,and the standard entropy change $\left(\Delta_{\text{sys}} S^{\theta}\right)$ of the same reaction is $-550 \ J \ K^{-1} \ mol^{-1}$,which one of the following statements is correct?
The $\Delta H$ and $\Delta S$ for a reaction at $1 \ \text{atm}$ pressure are $+30.558 \ \text{kJ}$ and $0.066 \ \text{kJ K}^{-1}$ respectively. The temperature at which the free energy change will be zero and below this temperature the nature of the reaction would be:
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