For irreversible expansion of an ideal gas under isothermal condition, the correct option is :

1. $∆\mathrm{U}=0, ∆{\mathrm{S}}_{\mathrm{total}}\ne 0$

2. $∆\mathrm{U}\ne 0, ∆{\mathrm{S}}_{\mathrm{total}}=0$

3. $∆\mathrm{U}=0, ∆{\mathrm{S}}_{\mathrm{total}}=0$

4. $∆\mathrm{U}\ne 0, ∆{\mathrm{S}}_{\mathrm{total}}\ne 0$

Which of the following options correctly represents the relationship between \(C_p \text { and } C_V\) for one mole of an ideal gas?

1. \(C_P=R C_V \)
2. \(C_V=RC_P \)
3. \(C_P+C_V=R \)
4. \(C_{\mathrm{P}}-\mathrm{C}_{\mathrm{V}}=\mathrm{R}\)

The molar heat capacity of water at constant pressure, C, is 75 JK^–1 mol^–1. When 1.0 kJ of heat is supplied to 100 g of water which is free to expand, the increase in temperature of the water is:

For which one of the following equations is $∆\mathrm{H}{°}_{\mathrm{react}}$ equal to $∆\mathrm{H}{°}_{\mathrm{f}}$ for the product:

1. N₂(g) + O₃(g) → N₂O₃(g)

2. CH₄(g) + 2Cl₂(g) → CH₂Cl₂(l) + 2HCl(g)

3. Xe(g) + 2F₂(g) → XeF₄(g)

4. 2CO(g) + O₂(g) → 2CO₂(g)

The formation of a solution from two components can be considered as:

The solution so formed will be ideal if:
1. ∆H_Soln = ∆H₁ + ∆H₂ + ∆H₃

2. ∆H_Soln = ∆H₁ + ∆H₂ – ∆H₃

3. ∆H_Soln = ∆H₁ – ∆H₂ – ∆H₃

4. ∆H_Soln = ∆H₃ – ∆H₁ – ∆H₂

What is the entropy change (in JK^–1 mol^–1) when one mole of ice is converted into water at 0 ºC? (The enthalpy change for the conversion of ice to liquid water is 6.0 KJ mol^–1 at 0 ºC)

For the reaction:

\(C_{3} H_{8} \left(\right. g \left.\right) + 5 O_{2} \left(\right. g \left.\right) \rightarrow 3 CO_{2} \left(\right. g \left.\right) + 4 H_{2} O \left(\right. l \left.\right)\) at constant temperature, ∆H – ∆E is:

Enthalpy  of $C{H}_4$ $+$ $\frac{1}{2}{O}_2$ $\to$ $C{H}_{3}OH$ is negative. If the enthalpy of combustion of $C{H}_4$ $and$ $C{H}_{3}OH$ are x and y respectively, then which relation is correct:

When 1 mol gas is heated at constant volume, the temperature is raised from 298 to 308 K. Heat supplied to the gas is 500 J. The correct statement among the following is:
1.  q = w = 500 J, ∆U = 0
2.  q = ∆U = 500 J, w = 0
3.  q = w = 500 J, ∆U = 0
4.  ∆U = 0, q = w = – 500 J

For the given reaction 

$2H_2{O}_2\left(l\right)$ $\to$ $2H_{2}O\left(l\right)$ $+$ $O_2\left(g\right)$, the heat of formations of ${\mathrm{H}}_2{\mathrm{O}}_2\left(\mathrm{l}\right)$ $$ $\mathrm{and}$ ${\mathrm{H}}_2\mathrm{O}$ $\left(\mathrm{l}\right)$ $$are -188 kJ/mol & -286 KJ/mol respectively. The change in the enthalpy of the reaction will be:

1.  – 196 kJ/mol

2.  + 196 kJ/mol

3.  + 948 kJ/mol

4.  – 948 kJ/mol