Mock Test 1 — Chemical Thermodynamics

The first law of thermodynamics is statement of:

Heat absorbed at constant pressure equals:

For an ideal gas, $\Delta U$ depends on:

Bond dissociation enthalpy of H-Cl is $431$ kJ/mol. This is energy:

For a reversible isothermal process: $W$ for ideal gas:

The criterion for spontaneity at const $T, P$:

Standard enthalpy of formation of O₂(g):

The entropy of universe in spontaneous process:

$\Delta H$ for reaction is $-10$ kcal. Energy released as heat at const $P$:

The standard free energy of formation $\Delta_f G^\ominus$ of elements is:

Number of moles in a system if $\Delta U = nC_V\Delta T = 600$ J, $C_V = 12.47$ J/(mol·K), $\Delta T = 24.06$ K (rounded integer):

For reaction at $T = 500$ K, $K = 10$. $\Delta G^\ominus$ (in kJ/mol, magnitude, integer using $R = 8.314$, $\ln 10 = 2.303$):

$\Delta H$ of reaction = $-200$ kJ/mol. $\Delta S = -100$ J/(mol·K). Find temperature above which reaction becomes non-spontaneous (in K):

Assertion (A): A spontaneous reaction always has negative $\Delta H$. Reason (R): Negative $\Delta H$ favors spontaneity at all temperatures.

Assertion (A): $\Delta H = \Delta U$ for reactions where $\Delta n_g = 0$. Reason (R): $\Delta H = \Delta U + \Delta n_g RT$.