Thermodynamics

In an irreversible process taking place at constant T and P and in which only pressure-volume work is being done, the change in Gibbs free energy (dG) and change in entropy (dS), satisfy the criteria :

The correct relationship between free energy change in a reaction and the corresponding equilibrium constant K_c is :

In conversion of lime-stone to lime, CaCO₃(s) $$\to$$ CaO(s) + CO₂ (g) the vales of ∆H° and ∆S° are +179.1 kJ mol−1 and 160.2 J/K respectively at 298 K and 1 bar. Assuming that ∆H° do not change with temperature, temperature above which conversion of limestone to lime will be spontaneous is :

Standard entropy of X₂, Y₂ and XY₃ are 60, 40 and 50 JK^−1 mol^−1 , respectively. For the reaction, $${1 \over 2} X_2$$ + $${3 \over 2} Y_2 \to$$ XY₃, $$\Delta H$$ = -30 kJ, to be at equilibrium, the temperature will be :

For a particular reversible reaction at temperature T, ∆H and ∆S were found to be both +ve. If T_e is the temperature at equilibrium, the reaction would be spontaneous when :

The entropy change involved in the isothermal reversible expansion of 2 moles of an ideal gas from a volume of 10 dm³ to a volume of 100 dm³ at 27^oC is :

The incorrect expression among the following is :

The following reaction is performed at 298 K 2NO(g) + O₂ (g) $$\leftrightharpoons$$ 2NO₂ (g) The standard free energy of formation of NO(g) is 86.6 kJ/mol at 298 K. What is the standard free energy of formation of NO₂(g) at 298 K? (KP = 1.6 × 10¹²)

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 :

For the reaction, A(g) + B(g) $$\to$$ C(g) + D(g), $$\Delta$$H^o and $$\Delta$$S^o are, respectively, − 29.8 kJ mol^−1 and −0.100 kJ K^−1 mol^−1 at 298 K. The equilibrium constant for the reaction at 298 K is :

At 320 K, a gas A₂ is 20% dissociated to A(g). The standard free energy change at 320 K and 1 atm in J mol⁻¹ is approximately : (R = 8.314 JK⁻¹ mol⁻¹; ln2 = 0.693; ln 3 = 1.098)

For which of the following processes, $$\Delta$$S is negative ?

Given (i)   2Fe₂O₃(s) $$\to$$ 4Fe(s) + 3O₂(g); $$\Delta$$_rG^o = + 1487.0 kJ mol⁻¹ (ii)   2CO(g) + O₂(g) $$\to$$ 2CO₂(g); $$\Delta$$_rG^o = $$-$$ 514.4 kJ mol⁻¹ Free energy change, $$\Delta$$_rG^o for the reaction 2Fe₂O₃(s) + 6CO(g) $$\to$$ 4Fe(s) + 6CO₂(g) will be :

$$\Delta$$_fG^o at 500 K for substance 'S' in liquid state and gaseous state are +100.7 kcl mol-¹ and +103 kcal mol⁻¹, respectively. Vapour pressure of liquid 'S' at 500 K is approximately equal to : ( R = 2 cal K⁻¹ mol⁻¹ )

A process has $$\Delta$$H = 200 J mol^–1 and $$\Delta$$S = 40 JK^–1 mol^–1. Out of the values given below, choose the minimum temperature above which the process will be spontaneous :

For the chemical reaction X $$\rightleftharpoons$$ Y, the standard reaction Gibbs energy depends on temperature T (in K) as $$\Delta$$_rG^o (in kJ mol^–1) = 120 $$- {3 \over 8}$$ T. The major component of the reaction mixture at T is :

For the equilibrium, 2H₂O $$\rightleftharpoons$$ H₃O⁺ + OH^$$-$$, the value of $$\Delta$$G^o at 298 K is approximately :

The reaction, MgO(s) + C(s) $$\to$$ Mg(s) + CO(g), for which $$\Delta$$_rH^o + 491.1 kJ mol^–1 and $$\Delta$$_rS^o = 198.0 JK^–1 mol^–1, is not feasible at 298 K. Temperature above which reaciton will be feasible is :

The standard reaction Gibbs energy for a chemical reaction at an absolute temperature T is given by $$\Delta$$_rG^o = A – BT Where A and B are non-zero constants. Which of the following is TRUE about this reaction?