Assume each reaction is carried out in an open container. For which of the following reactions will $∆\mathrm{H\; be\; equal\; to }∆\mathrm{U}?$

1. ${\mathrm{PCl}}_3\left(\mathrm{g}\right)$ $\to {\mathrm{PCl}}_3\left(\mathrm{g}\right)$ $+{\mathrm{Cl}}_2\left(\mathrm{g}\right)$ $$

2. $2\mathrm{CO}\left(\mathrm{g}\right)$ $+{\mathrm{O}}_2\left(\mathrm{g}\right)$ $\to 2{\mathrm{CO}}_2\left(\mathrm{g}\right)$ $$

3. ${\mathrm{H}}_2\left(\mathrm{g}\right)$ $+{\mathrm{Br}}_2\left(\mathrm{g}\right)$ $\to 2\mathrm{HBr}\left(\mathrm{g}\right)$

4. $\mathrm{C}\left(\mathrm{s}\right)$ $+2{\mathrm{H}}_2\mathrm{O}\left(\mathrm{g}\right)$ $\to 2{\mathrm{H}}_2\left(\mathrm{g}\right)$ $+{\mathrm{CO}}_2\left(\mathrm{g}\right)\hspace{0.17em}$

Combustion of glucose takes place according to the equation,

${\mathrm{C}}_6{\mathrm{H}}_{12}{\mathrm{O}}_6+6{\mathrm{O}}_2\to 6{\mathrm{CO}}_2+6{\mathrm{H}}_2\mathrm{O}$, $∆ H = -72 K cal$

Energy required for the production of 1.6 g of glucose is -

(Molecular mass of glucose = 180 g)

1. 0.064 kcal

2. 0.64 kcal

3. 6.4 kcal

4. 64 kcal

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:

For the following given equations and $∆\mathrm{H}°$ values, determine the enthalpy of reaction at 298 K for the reaction-

C₂H₄(g) + 6F₂(g) $\to$ 2CF₄(g) + 4HF(g)

H₂(g) + F₂(g) $\to$ 2HF(g)       $∆{\mathrm{H}}_1^{°}$= -537 kJ

C(s) + 2F₂(g) $\to$CF₄(g)         $∆{\mathrm{H}}_2^{°}$=-680 kJ

2C(s) + 2H₂(g) $\to$C₂H₄(g)    $∆{\mathrm{H}}_3^{°}$= 52 kJ

1. -1165 

2. -2486

3. +1165

4. +2486

The bond energies of $\mathrm{C}\equiv \mathrm{C}$, C-H, H-H, and C=C are 198, 98, 103 and145 kcal respectively.

The enthalpy change of the reaction $\mathrm{HC}\equiv \mathrm{CH}+{\mathrm{H}}_2\to {\mathrm{C}}_2{\mathrm{H}}_4$ would be-

1. 48 kcal

2. 96 kcal

3. -40 kcal

4. -152 kcal

During complete combustion of one mole of butane, 2658 kJ of heat is released. The thermochemical reaction for this change is -

1. $2{\mathrm{C}}_4{\mathrm{H}}_{10}\left(\mathrm{g}\right)+13{\mathrm{O}}_2\left(\mathrm{g}\right)\to 8{\mathrm{CO}}_2\left(\mathrm{g}\right)+10{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right);$
${∆}_{\mathrm{c}}\mathrm{H}=-2658.0 \mathrm{kJ} {\mathrm{mol}}^{-1}$

2. ${\mathrm{C}}_4{\mathrm{H}}_{10}\left(\mathrm{g}\right)+\frac{13}{2}{\mathrm{O}}_2\left(\mathrm{g}\right)\to 4{\mathrm{CO}}_2\left(\mathrm{g}\right)+5{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right);$
${∆}_{\mathrm{c}}\mathrm{H}=-1329.0 \mathrm{kJ} {\mathrm{mol}}^{-1}$

3. ${\mathrm{C}}_4{\mathrm{H}}_{10}\left(\mathrm{g}\right)+\frac{13}{2}{\mathrm{O}}_2\left(\mathrm{g}\right)\to 4{\mathrm{CO}}_2\left(\mathrm{g}\right)+5{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right);$
${∆}_{\mathrm{c}}\mathrm{H}=-2658.0 \mathrm{kJ} {\mathrm{mol}}^{-1}$

4. ${\mathrm{C}}_4{\mathrm{H}}_{10}\left(\mathrm{g}\right)+\frac{13}{2}{\mathrm{O}}_2\left(\mathrm{g}\right)\to 4{\mathrm{CO}}_2\left(\mathrm{g}\right)+5{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right);$
${∆}_{\mathrm{c}}\mathrm{H}=+2658.0 \mathrm{kJ} {\mathrm{mol}}^{-1}$

Which plot represents an exothermic reaction?

1.   
2.   

3.    
4.   

Consider the following diagram for a reaction $\mathrm{A}\to \mathrm{C:}$

The nature of the reaction is-

As an isolated box, equally partitioned, contains two ideal gasses A and B as shown:

When the partition is removed, the gases mix. The changes in enthalpy $(∆\mathrm{H})$ and entropy $(∆\mathrm{S})$ in the process, respectively, are

1. Zero, positive

2. Zero, negative

3. Positive, zero

4. Negative, zero