Chemical Thermodynamics

The standard state Gibbs free energies of formation of C(graphite) and C(diamond) at T = 298 K are

$$\Delta_f G^o[\mathrm{C(graphite)}] = 0\ \mathrm{kJ\,mol^{-1}}$$
$$\Delta_f G^o[\mathrm{C(diamond)}] = 2.9\ \mathrm{kJ\,mol^{-1}}.$$

The standard state means that the pressure should be 1 bar, and substance should be pure at a given temperature. The conversion of graphite [C(graphite)] to diamond [C(diamond)] reduces its volume by $2 \times 10^{-6}\ \mathrm{m^3\,mol^{-1}}$. If C(graphite) is converted to C(diamond) isothermally at T = 298 K, the pressure at which C(graphite) is in equilibrium with C(diamond), is

[Useful information: 1 J = 1 kg m$^2$ s$^{-2}$; 1 Pa = 1 kg m$^{-1}$ s$^{-2}$; 1 bar = $10^5$ Pa]

For a reaction taking place in a container in equilibrium with its surroundings, the effect of temperature on its equilibrium constant $K$ in terms of change in entropy is described by

Given

$\ce{C_{(graphite)} + O2(g) -> CO2(g)}$ ; $\Delta_rH^\circ=-393.5$ kJ mol$^{-1}$

$\ce{H2(g) + 1/2O2(g) -> H2O(l)}$ ; $\Delta_rH^\circ=-285.8$ kJ mol$^{-1}$

$\ce{CO2(g) + 2H2O(l) -> CH4(g) + 2O2(g)}$ ; $\Delta_rH^\circ=+890.3$ kJ mol$^{-1}$

Based on the above thermochemical equations, the value of $\Delta_rH^\circ$ at 298 K for the reaction $\ce{C_{(graphite)} + 2H2(g) -> CH4(g)}$ will be :

ΔU is equal to :

For a reaction, $\ce{A(g) -> A(l)}$; ΔH = −3RT. The correct statement for the reaction is :

The enthalpy change on freezing of 1 mol of water at 5°C to ice at −5°C is :

(Given $\Delta_{fus}H=6$ kJ mol$^{-1}$ at 0°C, $C_p(\ce{H2O}, l)=75.3$ J mol$^{-1}$ K$^{-1}$, $C_p(\ce{H2O}, s)=36.8$ J mol$^{-1}$ K$^{-1}$)

An ideal gas undergoes isothermal expansion at constant pressure. During the process :

A gas undergoes change from state A to state B. In this process, the heat absorbed and work done by the gas is 5 J and 8 J, respectively. Now gas is brought back to A by another process during which 3 J of heat is evolved. In this reverse process of B to A :

One mole of an ideal gas at 300 K in thermal contact with surroundings expands isothermally from 1.0 L to 2.0 L against a constant pressure of 3.0 atm. In this process, the change in entropy of surroundings $(\Delta S_{surr})$ in $\mathrm{J\,K^{-1}}$ is

(1 L atm = 101.3 J)

The heats of combustion of carbon and carbon monoxide are −393.5 and −283.5 kJ mol$^{-1}$, respectively. The heat of formation (in kJ) of carbon monoxide per mole is :

A reaction at 1 bar is non-spontaneous at low temperature but becomes spontaneous at high temperature. Identify the correct statement about the reaction among the following :

If 100 mole of $\ce{H2O2}$ decompose at 1 bar and 300 K, the work done (kJ) by one mole of $\ce{O2(g)}$ as it expands against 1 bar pressure is :
$$\ce{2H2O2(l) <=> 2H2O(l) + O2(g)}$$
(R = 8.3 J K$^{-1}$ mol$^{-1}$)

Match the thermodynamic processes given under Column I with the expressions given under Column II.

Enthalpy of dissociation (in kJ mol⁻¹) of acetic acid obtained from the Expt. 2 is

The heat of atomization of methane and ethane are 360 kJ/mol and 620 kJ/mol, respectively. The longest wavelength of light capable of breaking the C - C bond is (Avogadro number = $6.02\times10^{23}$, $h=6.62\times10^{-34}$ J s) :

An ideal gas in a thermally insulated vessel at internal pressure $=P_1$, volume $=V_1$ and absolute temperature $=T_1$ expands irreversibly against zero external pressure, as shown in the diagram. The final internal pressure, volume and absolute temperature of the gas are $P_2$, $V_2$ and $T_2$, respectively. For this expansion,

For the process
$$\ce{H2O(l) -> H2O(g)}$$
at $T=100\ ^\circ$C and 1 atmosphere pressure, the correct choice is

For complete combustion of ethanol,
$$\ce{C2H5OH(l) + 3O2(g) -> 2CO2(g) + 3H2O(l)},$$
the amount of heat produced as measured in bomb calorimeter, is 1364.47 kJ mol$^{-1}$ at 25°C. Assuming ideality the Enthalpy of combustion, $\Delta_cH$, for the reaction will be :

($R=8.314$ kJ mol$^{-1}$)

The standard enthalpy of formation of $\ce{NH3}$ is $-46.0$ kJ/mol. If the enthalpy of formation of $\ce{H2}$ from its atoms is $-436$ kJ/mol and that of $\ce{N2}$ is $-712$ kJ/mol, the average bond enthalpy of N-H bond in $\ce{NH3}$ is :

The molar heat capacity ($C_p$) of $\ce{CD2O}$ is 10 cals at 1000 K. The change in entropy associated with cooling of 32 g of $\ce{CD2O}$ vapour from 1000 K to 100 K at constant pressure will be :

(D = deuterium, at. mass = 2 u)