Choose the correct answer. A thermodynamic state function is a quantity

The enthalpy of combustion of methane, graphite and dihydrogen at 298 K are, $-890.3 \text{ kJ mol}^{-1}$, $-393.5 \text{ kJ mol}^{-1}$, and $-285.8 \text{ kJ mol}^{-1}$ respectively. Enthalpy of formation of $CH_4(g)$ will be

A reaction, $A + B \rightarrow C + D + q$ is found to have a positive entropy change. The reaction will be

For the process to occur under adiabatic conditions, the correct condition is:

In a process, 701 J of heat is absorbed by a system and 394 J of work is done by the system. What is the change in internal energy for the process?

The enthalpies of all elements in their standard states are:

The reaction of cyanamide, $NH_2CN (s)$, with dioxygen was carried out in a bomb calorimeter, and $ΔU$ was found to be $-742.7 \text{ kJ mol}^{-1}$ at 298 K. Calculate enthalpy change for the reaction at 298 K.
$NH_2CN(g) + \frac{3}{2}O_2(g) \rightarrow N_2(g) + CO_2(g) + H_2O(l)$

The equilibrium constant for a reaction is 10. What will be the value of $ΔG^o$ ? $R = 8.314 \text{ JK}^{-1} \text{ mol}^{-1}$, $T = 300 K$.

For the reaction
$2 A(g) + B(g) \rightarrow 2D(g)$
$ΔU^o = -10.5 \text{ kJ}$ and $ΔS^o = -44.1 \text{ JK}^{-1}$.
Calculate $ΔG^o$ for the reaction, and predict whether the reaction may occur spontaneously.

Given
$N_2(g) + 3H_2(g) \rightarrow 2NH_3(g)$; $Δ_rH^o = -92.4 \text{ kJ mol}^{-1}$
What is the standard enthalpy of formation of $NH_3$ gas?

For the reaction,
$2 Cl(g) \rightarrow Cl_2(g)$, what are the signs of $ΔH$ and $ΔS$ ?

Calculate the standard enthalpy of formation of $CH_3OH(l)$ from the following data:
$CH_3OH (l) + \frac{3}{2} O_2(g) \rightarrow CO_2(g) + 2H_2O(l)$ ; $Δ_rH^o = -726 \text{ kJ mol}^{-1}$
$C(graphite) + O_2(g) \rightarrow CO_2(g)$ ; $Δ_cH^o = -393 \text{ kJ mol}^{-1}$
$H_2(g) + \frac{1}{2} O_2(g) \rightarrow H_2O(l)$; $Δ_fH^o = -286 \text{ kJ mol}^{-1}$.

For the reaction at 298 K,
$2A + B \rightarrow C$
$ΔH = 400 \text{ kJ mol}^{-1}$ and $ΔS = 0.2 \text{ kJ K}^{-1} \text{ mol}^{-1}$
At what temperature will the reaction become spontaneous considering $ΔH$ and $ΔS$ to be constant over the temperature range.

Calculate the enthalpy change for the process
$CCl_4(g) \rightarrow C(g) + 4 Cl(g)$
and calculate bond enthalpy of C – Cl in $CCl_4(g)$.
$Δ_{vap}H^o(CCl_4) = 30.5 \text{ kJ mol}^{-1}$.
$Δ_fH^o (CCl_4) = -135.5 \text{ kJ mol}^{-1}$.
$Δ_aH^o (C) = 715.0 \text{ kJ mol}^{-1}$, where $Δ_aH^o$ is enthalpy of atomisation
$Δ_aH^o (Cl_2) = 242 \text{ kJ mol}^{-1}$

For an isolated system, $ΔU = 0$, what will be $ΔS$ ?

Calculate the number of kJ of heat necessary to raise the temperature of 60.0 g of aluminium from $35^\circ C$ to $55^\circ C$. Molar heat capacity of Al is $24 \text{ J mol}^{-1} K^{-1}$.

Calculate the entropy change in surroundings when 1.00 mol of $H_2O(l)$ is formed under standard conditions. $Δ_fH^o = -286 \text{ kJ mol}^{-1}$.

Comment on the thermodynamic stability of $NO(g)$, given
$\frac{1}{2} N_2(g) + \frac{1}{2}O_2(g) \rightarrow NO(g)$; $Δ_rH^o = 90 \text{ kJ mol}^{-1}$
$NO(g) + \frac{1}{2}O_2(g) \rightarrow NO_2(g)$: $Δ_rH^o= -74 \text{ kJ mol}^{-1}$

$ΔU^o$ of combustion of methane is $-X \text{ kJ mol}^{-1}$. The value of $ΔH^o$ is

Calculate the enthalpy change on freezing of 1.0 mol of water at $10.0^\circ C$ to ice at $-10.0^\circ C$. $Δ_{fus}H = 6.03 \text{ kJ mol}^{-1}$ at $0^\circ C$.
$C_p [H_2O(l)] = 75.3 \text{ J mol}^{-1} K^{-1}$
$C_p [H_2O(s)] = 36.8 \text{ J mol}^{-1} K^{-1}$