Thermodynamics practice · Medium

Entropy change of vaporisation at constant pressure is given by:

The enthalpy of neutralization of a strong monobasic acid by a strong monoacidic base is $-13,700 \mathrm{~cal}.$ A certain monobasic weak acid is $10 \%$ ionized in a molar solution. If the enthalpy of ionization of the weak acid is $+400 \mathrm{~cal} / \mathrm{mole},$ then what is the enthalpy of neutralization of one molar solution of the weak acid?

For a system in equilibrium

When $50 \mathrm{~mL}$ of $1 \mathrm{~M} \mathrm{~HCl}$ (aq) with $50 \mathrm{~mL}$ of $1 \mathrm{~M} \mathrm{~NaOH}(\mathrm{aq})$, the enthalpy of neutralisation will be

The $\Delta_{\mathrm{f}} \mathrm{H}^{\circ}$ for $\mathrm{N}_2 \mathrm{O}_3(\mathrm{~g})$ in $\mathrm{kJ} / \mathrm{mol}$ on the basis of the following data is :

$\begin{array}{ll} 2 \mathrm{NO}(\mathrm{g})+\mathrm{O}_2(\mathrm{~g}) \longrightarrow 2 \mathrm{NO}_2(\mathrm{~g}) & \Delta_{\mathrm{r}} \mathrm{H}^{\circ}=-114 \mathrm{~kJ} / \mathrm{mol} \\ 4 \mathrm{NO}_2(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \longrightarrow 2 \mathrm{~N}_2 \mathrm{O}_5(\mathrm{~g}) & \Delta_{\mathrm{r}} \mathrm{H}^{\circ}=-102.6 \mathrm{~kJ} / \mathrm{mol} \\ \Delta_{\mathrm{f}} \mathrm{H}^{\circ} \mathrm{NO}(\mathrm{g})=90.2 \mathrm{~kJ} / \mathrm{mol} & \end{array}$

Tea placed in thermos flask is an example of

3.0 moles of an ideal gas is heated at constant pressure from $27^{\circ} \mathrm{C}$ to $127^{\circ} \mathrm{C}$. The work done by the gas in this expansion is:

Under identical conditions, how many millilitres of $1 \mathrm{~M}-\mathrm{KOH}$ and $2 \mathrm{~M}-\mathrm{H}_2 \mathrm{SO}_4$ solutions are required to produce a resulting volume of $100 \mathrm{~ml}$ with the highest rise in temperature?

One gram sample of ${NH}_4 {NO}_3$ is decomposed in a bomb calorimeter, the temperature of the calorimeter increases by $6.12 {~K}$. The heat capacity of the system is $1.23 {~kJ} / {g} {deg}$. What is the molar heat of decomposition for ${NH}_4 {NO}_3$ ?

If 900 J/g of heat is exchanged at boiling point of water then increase in entropy

The equation that represents general Van't
Hoff equation is

A diatomic ideal gas initially at $273 \mathrm{~K}$ is given $100 \text{ cal}$ heat due to which system did $210 \mathrm{~J}$ work. The molar heat capacity of the gas for the process is $(1 \mathrm{~cal}=4.2 \mathrm{~J})$

Heat of combustion of methane is $-800 \mathrm{~kJ}$. What is the heat of combustion for $4 \times 10^{-4} \mathrm{~kg}$ of methane?

Heat of combustion ∆H for C(s), ${\mathrm{H}_2}$(g) and $\mathrm{C}{\mathrm{H}_4}$ (g) are -94,-68 and -213 kcal/mole then ∆H for $\mathrm{C}\left( s \right) + 2{\mathrm{H}_2}\left( \mathrm{g} \right) \to \mathrm{C}{\mathrm{H}_4}\left( \mathrm{g} \right)\mathrm{i}s:$

The molar heat capacity of water in equilibrium with ice at constant pressure is

The sole criterion for the spontaneity of a process is

Statement I : Heat of neutralisation of $HF(aq)$, a weak acid, with $NaOH(aq)$ is less than 13.7 kcal, in an exothermic reaction
Statement II : Some heat is lost in the ionisation of a weak acid

$1 \mathrm{~mole} \mathrm{~H_2} \mathrm{O}(l)$ at $-10^{\circ} \mathrm{C}$ is converted to 1 mole $\mathrm{H}_2 \mathrm{O}(s)$ at $-10^{\circ} \mathrm{C}$. Entropy change is the sum of all steps but not of

Assertion : Many endothermic reactions that are not spontaneous at room temperature become spontaneous at high temperature.
Reason : Entropy of the system increases with increase in temperature.

Find the bond energy of $\mathrm{S}-\mathrm{S}$ bond from the following data.

$\begin{aligned} & \mathrm{C}_2 \mathrm{H}_5-\mathrm{S}-\mathrm{C}_2 \mathrm{H}_5(\mathrm{g}); \Delta H_{\mathrm{f}}^{\mathrm{o}}=-148 \mathrm{~kJ} \\ & \mathrm{C}_2 \mathrm{H}_5-\mathrm{S}-\mathrm{S}-\mathrm{C}_2 \mathrm{H}_5(\mathrm{g}); \Delta H_{\mathrm{f}}^{\circ}=-202 \mathrm{~kJ} \\ & \mathrm{~S}(\mathrm{g});\Delta H_{\mathrm{f}}^{\circ}=222 \mathrm{~kJ} \end{aligned}$

Why is the heat evolved in neutralisation of HF
is highest ?

In the reversible adiabatic expansion of an ideal monoatomic gas, the final volume is 8 times the initial volume. The ratio of final temperature to initial temperature is

When $5 \mathrm{~mol}$ of an ideal gas expands from $10 \mathrm{~L}$ to $20 \mathrm{~L}$, temperature of the gas changes from $300 \mathrm{~K}$ to $400 \mathrm{~K}$. If the heat capacity ${C}_{{p}}$ of the gas is $4.0 \mathrm{~cal} \mathrm{~mol}^{-1} \mathrm{~K}^{-1}$, then entropy change $\left(\text{in cal } \mathrm{K}^{-1}\right)$ of the gas is $(\log 2=0.30$, $\log 3=0.47$ and $R \simeq 2 \mathrm{~cal} \mathrm{~mol}^{-1} \mathrm{~K}^{-1}$)

$\Delta_{\mathrm{f}} H^{\circ}$ for $\mathrm{NF}_3(\mathrm{g})$ is $-113 \mathrm{~kJ} / \mathrm{mol}.$ Bond energy for the $\mathrm{N}-\mathrm{F}$ bond is $273.5 \mathrm{~kJ} / \mathrm{mol}.$ The bond energies of $\mathrm{N}_2$ and $\mathrm{F}_2,$ if their magnitudes are in the ratio $6: 1,$ are respectively.

The term $-T \Delta S_{\text {total}}$ at constant temperature and pressure is known as

Assertion: The enthalpy of atomization increases with an increase in atomic number in a period.

Reason: The increase in atomic number results in an increase in the number of protons and electrons in the atom.

Internal energy and pressure of a gas of unit volume are related as:

Which compound will absorb the maximum amount of heat when dissolved in the same amount of water? (Integral heats of solution at ${25^0}c$ in kcal/mol of each solute are given in brackets):

If the bond energies of $\mathrm{C}=\mathrm{C}$ and $\mathrm{C}-\mathrm{C}$ at $298 \mathrm{~K}$ are $590$ and $331 \mathrm{~kJ} \mathrm{~mol}^{-1}$ respectively then enthalpy of polymerization per mole of ethylene is:

Enthalpies of combustion of $\mathrm{CH}_4, \mathrm{C}_2 \mathrm{H}_6$ and $\mathrm{C}_3 \mathrm{H}_8$ are $-210,-370$ and $-526 \mathrm{~kcal} / \mathrm{mol},$ respectively. Enthalpy of combustion of hexane can be predicted as