Calculate the enthalpy change when $50 \ mL$ of $0.01 \ M$ $Ca(OH)_2$ reacts with $25 \ mL$ of $0.01 \ M$ $HCl$. Given that $\Delta H^o_{\text{neutralisation}}$ of strong acid and strong base is $-57.1 \ kJ \ mol^{-1}$. [Assuming that $Ca(OH)_2$ is a strong base]

One mole of an ideal gas is allowed to expand reversibly and adiabatically from a temperature of $27\,^oC$. If the work done during the process is $3\,kJ$,the final temperature will be equal to $(C_v = 20\,J\,K^{-1} \, mol^{-1})$

For the combustion of one mole of acetic acid,the work done at $298 \ K$ is (in $J$)

One mole of an ideal monoatomic gas is subjected to changes as shown in the graph. The magnitude of the work done (by the system or on the system) is $...........J$ (nearest integer). Given: $\log 2=0.3, \ln 10=2.3$

$1 \ mol$ of $NH_3$ $(\gamma = 4/3)$ gas at $27 \ ^\circ C$ is expanded under reversible adiabatic conditions to make the volume $8$ times. Calculate the work done in $cal$.

One mole of an ideal monoatomic gas undergoes two reversible processes ($A \rightarrow B$ and $B \rightarrow C$) as shown in the given figure:
$A \rightarrow B$ is an adiabatic process. If the total heat absorbed in the entire process ($A \rightarrow B$ and $B \rightarrow C$) is $R T_2 \ln 10$,the value of $2 \log V_3$ is . . . . . [Use,molar heat capacity of the gas at constant pressure,$C_{p, m} = \frac{5}{2} R$ ]