Thermodynamics Questions

$\text{C(s)}+2{\text{H}}_2\text{(g)}\to {\text{CH}}_4\text{(g)};\Delta H=-74.8{\text{kJ mol}}^{-1}$. Which of the following diagrams gives an accurate representation of the above reaction? [$R\to$ reactants; $P\to$ products]

Which of the following is correct option for free expansion of an ideal gas under adiabatic condition?

If the enthalpy change for the transition of liquid water to steam is $30{\text{kJ mol}}^{-1}$ at $27^{\circ }\text{C}$, the entropy change for the process would be:

If the $E_{cell}^{\circ }$ for a given reaction has a negative value, then which of the following gives the correct relationships for the values of $\Delta G^{\circ }$ and $K_{eq}$?

Enthalpy change for the reaction $2{\text{H}}_{(\text{g})}\to {\text{H}}_{2(\text{g})}$ is $-869.6\text{kJ}$. The dissociation energy of $\text{H}-\text{H}$ bond is:

Standard enthalpy of vapourisation ${\Delta }_{\text{vap}}{H}^{\circ }$ for water at $100^{\circ }\text{C}$ is $40.66\text{kJ/mol}$. The internal energy of vapourisation of water at $100^{\circ }\text{C}$ (in $\text{kJ/mol}$) is:

Equal volumes of two monoatomic gases, $A$ and $B$, at same temperature and pressure are mixed. The ratio of specific heats ($C_p/C_v$) of the mixture will be:

The enthalpy of fusion of water is $1.435\text{kcal/mol}$. The molar entropy change for the melting of ice at $0^{\circ }\text{C}$ is:

In which of the following reactions, standard reaction entropy change ($\Delta S^{\circ }$) is positive and standard Gibb’s energy change ($\Delta G^{\circ }$) decreases sharply with increasing temperature?

For an endothermic reaction, energy of activation is $E_a$ and enthalpy of reaction is $\Delta H$ (both of these in $\text{kJ/mol}$). The minimum value of $E_a$ will be:

A reaction having equal energies of activation for forward and reverse reactions has:

Using the Gibbs energy change, $\Delta G^{\circ }=+63.3\text{kJ}$, for the reaction:
$${\text{Ag}}_2{\text{CO}}_3(s)\rightleftharpoons 2{\text{Ag}}^{+}(aq)+{\text{CO}}_3^{2-}(aq)$$
The $K_{sp}$ of ${\text{Ag}}_2{\text{CO}}_3(s)$ in water at $25^{\circ }\text{C}$ is: ($R=8.314{\text{J K}}^{-1}{\text{mol}}^{-1}$)

For the reaction ${\text{X}}_2{\text{O}}_4(l)\to 2{\text{XO}}_2(g)$, $\Delta U=2.1\text{k cal}$ and $\Delta S=20{\text{cal K}}^{-1}$ at $300\text{K}$. Hence, $\Delta G$ is:

The formation of the oxide ion, ${\text{O}}^{2-}(g)$, from oxygen atom requires first an exothermic and then an endothermic step as shown below:$\text{O}(g)+e^{-}⟶{\text{O}}^{-}(g);{\Delta }_f{H}^{\ominus }=-141{\text{kJ mol}}^{-1}$${\text{O}}^{-}(g)+e^{-}⟶{\text{O}}^{2-}(g);{\Delta }_f{H}^{\ominus }=+780{\text{kJ mol}}^{-1}$ Thus the process of formation of ${\text{O}}^{2-}$ in gas phase is unfavourable though ${\text{O}}^{2-}$ is isoelectronic with neon. It is due to the fact that:

The heat of combustion of carbon to ${\text{CO}}_2$ is $-393.5\text{kJ/mol}$. The heat released upon formation of $35.2\text{g}$ of ${\text{CO}}_2$ from carbon and oxygen gas is:

The correct thermodynamic conditions for the spontaneous reaction at all temperatures is:

Consider the following liquid – vapour equilibrium: $\text{Liquid}\rightleftharpoons \text{Vapour}$. Which of the following relations is correct?

For a sample of perfect gas when its pressure is changed isothermally from $p_i$ to $p_f$, the entropy change is given by:

Consider the following reactions:(i) ${\text{H}}^{+}(aq)+{\text{OH}}^{-}(aq)={\text{H}}_2\text{O}(l)$, $\Delta H=-x_1{\text{ kJmol}}^{-1}$ (ii) ${\text{H}}_2(g)+\frac{1}{2}{\text{O}}_2(g)={\text{H}}_2\text{O}(l)$, $\Delta H=-x_2{\text{ kJmol}}^{-1}$ (iii) ${\text{CO}}_2(g)+{\text{H}}_2(g)=\text{CO}(g)+{\text{H}}_2\text{O}(l)$, $\Delta H=-x_3{\text{ kJmol}}^{-1}$ (iv) ${\text{C}}_2{\text{H}}_2(g)+\frac{5}{2}{\text{O}}_2(g)=2{\text{CO}}_2+{\text{H}}_2\text{O}(l)$, $\Delta H=-x_4{\text{ kJmol}}^{-1}$ The enthalpy of formation of ${\text{H}}_2\text{O}(l)$ is:

$\Delta H_f^{\circ }(298\text{K})$ of methanol is given by the chemical equation:

For the reaction of one mole of zinc dust with one mole of sulphuric acid in a bomb calorimeter, $\Delta U$ and $w$ corresponds to:

For the chemical equilibrium, ${\text{CaCO}}_{3(s)}\rightleftharpoons {\text{CaO}}_{(s)}+{\text{CO}}_{2(g)}$, $\Delta H_r^{\circ }$ can be determined from which one of the following plots?

Identify the correct statement for change of Gibbs energy for a system ($\Delta G_{\text{system}}$) at constant temperature and pressure:

Assume each reaction is carried out in an open container. For which reaction will $\Delta H=\Delta E$ ?

The enthalpy and entropy change for the reaction:${\text{Br}}_2(\text{l})+{\text{Cl}}_2(\text{g})\to 2\text{BrCl}(\text{g})$ Are $30{\text{kJ mol}}^{-1}$ and $105{\text{JK}}^{-1}{\text{mol}}^{-1}$ respectively. The temperature at which the reaction will be in equilibrium is:

The enthalpy of combustion of ${\text{H}}_2$, cyclohexene (${\text{C}}_6{\text{H}}_{10}$) and cyclohexane (${\text{C}}_6{\text{H}}_{12}$) are -241, -3800 and -3920 kJ per mol respectively. Heat of hydrogenation of cyclohexene is :

In which of the following the hydration energy is higher than the lattice energy?

Given:$(1)$ $A\to 2B,\Delta H=+150$$(2)$ $3B\to 2C+D,\Delta H=-125$$(3)$ $E+A\to 2D,\Delta H=+350$ For $B+D\to E+2C$, $\Delta H$ will be:

Given that bond energy of $\text{H—H}$ and $\text{Cl-Cl}$ is $430{\text{ kJ mol}}^{-1}$ and $240{\text{ kJ mol}}^{-1}$ respectively and $\Delta H_f$ for $\text{HCl}$ is $-90{\text{ kJ mol}}^{-1}$. The bond enthalpy of $\text{HCl}$ is:

Which of the following are not state functions?(I) $q+w$ (II) $q$ (III) $w$ (IV) $H-TS$

Bond dissociation enthalpy of ${\text{H}}_2$, ${\text{Cl}}_2$ and $\text{HCl}$ are $434$, $242$ and $431\text{kJ/mol}$ respectively. Enthalpy of formation of $\text{HCl}$ is:

For the gas phase reaction ${\text{PCl}}_5(g)\rightleftharpoons {\text{PCl}}_3(g)+{\text{Cl}}_2(g)$, which of the following conditions is correct?

The values of $\Delta H$ and $\Delta S$ for the reaction $\text{C}(\text{graphite})+{\text{CO}}_2(g)\to 2\text{CO}(g)$ are $170\text{kJ}$ and $170{\text{JK}}^{-1}$ respectively. This reaction will be spontaneous at:

From the following bond energies: $\text{H-H}:431.37\text{kJ/mol}$ $\text{C=C}:606.10\text{kJ/mol}$ $\text{C-C}:336.49\text{kJ/mol}$ $\text{C-H}:410.50\text{kJ/mol}$ Enthalpy for the reaction will be:

Standard entropies of $X_2,Y_2$ and $X{Y}_3$ are $60,40$ and $50{\text{JK}}^{-1}{\text{mol}}^{-1}$ respectively. For the reaction $\frac{1}{2}{X}_2+\frac{3}{2}{Y}_2\rightleftharpoons X{Y}_3,\Delta H=-30\text{kJ}$ to be at equilibrium, the temperature should be: