Thermodynamics practice · Medium

If the enthalpy change for the reaction is $2 \mathrm{NH}_{3(\mathrm{~g})} \rightarrow \mathrm{N}_{2(\mathrm{~g})}+3 \mathrm{H}_{2(\mathrm{~g})}$ is $92 \mathrm{~kJ} \mathrm{~mol}^{-1}$ the enthalpy of formation of ammonia is

The change in internal energy equals

Select the correct statements.

The correct relation for $K_f$ is/are:

If at 298 K the bond energies of $C - H,\,C - C,\,C = C$ and $H - H$ are respectively, 414, 347, 615 and 435 ${\text{kJ}}\,{\text{mo}}{{\text{l}}^{ - 1}},$ the value of enthalpy change for the reaction ${H_2}C = C{H_2}\left( g \right) + {H_2}\left( g \right) \to {H_3}C - C{H_3}\left( g \right)$
at 298 K will be:

The temperature of $15 \mathrm{~ml}$ of a strong acid increases by $2^{\circ} \mathrm{C}$ when $15 \mathrm{~ml}$ of a strong base is added to it. If $5 \mathrm{~ml}$ of each are mixed, temperature should increase by

Temperature of 5 moles of a gas is decreased
by 2K at constant pressure. Indicate the correct
statement

The latent heat of fusion of ice at $0^{\circ} \mathrm{C}$ being $1.44 \mathrm{~k~cal} \mathrm{~mol}^{-1}$ and difference in heat capacity of ice and water being $9.3 \mathrm{~cal} \mathrm{~mol}^{-1}$, thus latent heat of fusion of ice at $-27^{\circ} \mathrm{C}$ is

Statement I: The final temperature of an ideal
gas in adiabatic compression is less in reversible
compression than in irreversible compression
against a constant external pressure.

Statement II: The magnitude of work done on
an ideal gas in adiabatic compression is more in
irreversible process than that in reversible process.

$\mathrm{H}_2 \mathrm{S}(\mathrm{g}) \rightarrow \mathrm{HS}^{-}(\mathrm{g})+\mathrm{H}^{+}(\mathrm{g}), \Delta \mathrm{H}^{\circ}=\mathrm{x}_1,$
$\Delta \mathrm{H_f^\circ}\left(\mathrm{H}_2 \mathrm{S}\right. (\mathrm{g})) =\mathrm{x}_2, \Delta \mathrm{H_f^\circ}\left(\mathrm{H}^{+}(\mathrm{g})\right)=\mathrm{x}_3$
Hence, $\Delta \mathrm{H_f^\circ}\left(\mathrm{HS}^{-}\right)$ is:

The heat of combustion of carbon monoxide at constant pressure and at constant volume at $27^{\circ} \mathrm{C}$ will differ from one another by:

A: Dissolution of sugar in water proceed via increases in entropy.
R: Entropy decreases when egg is hard boiled.

${\mathrm{C}_{\mathrm{diamond}}} + {\mathrm{O}_2} \to \mathrm{C}{\mathrm{O}_2};\Delta H$=-395.3 kJ/mol
${\mathrm{C}_{\mathrm{graphite}}} + {\mathrm{O}_2} \to \mathrm{C}{\mathrm{O}_2};\Delta H$=-393.4 kJ/mol
${\mathrm{C}_{\mathrm{graphite}}} \to {\mathrm{C}_{\mathrm{diamond}}};\Delta \mathrm{H}$=?

Given that $\mathrm{C}+\mathrm{O}_2 \rightarrow \mathrm{CO}_2, \Delta \mathrm{H}^{\circ}=-\mathrm{x} \mathrm{~kJ}$

$2 \mathrm{CO}+\mathrm{O}_2 \rightarrow 2 \mathrm{CO}_2, \Delta \mathrm{H}^{\circ}=-\mathrm{y} \mathrm{~kJ}$

What is heat of formation of $\mathrm{CO}$?

For an adiabatic expansion of a perfect gas ∆P/P is equal to:

In the following reaction,

$\begin{aligned} & \mathrm{HF}(a q)+\mathrm{NaOH}(a q) \longrightarrow \mathrm{NaF}(a q)+\mathrm{H}_2 \mathrm{O}(l) ; \Delta H^\circ=-68.6 \mathrm{~kJ}^3 \mathrm{mol}^{-1} \\ \end{aligned}$

This value is much higher than the heat of neutralisation of strong acid with strong base $\left(-57.3 \mathrm{~kJ} \mathrm{~mol}^{-1}\right)$. This is because

A quantity of $1 \text{ g}$ of water on evaporation at
atmospheric pressure forms $1671 \, \text{cm}^3$
of steam.
Heat of vaporization is $540 \text{ cal/g}.$ The approximate
increase in internal energy is (1 L – atom = 24 cal)

Calculate the number of $\mathrm{kJ}$ of heat necessary to raise the temperature of $60 \mathrm{~g}$ of aluminium from $35^{\circ} \mathrm{C}$ to $55^{\circ} \mathrm{C}$. Molar heat capacity of ${Al}$ is $24 \mathrm{~J} \mathrm{~mol}^{-1} \mathrm{~K}^{-1}$.

The enthalpy of formation of HI is 30.4 kJ. Which statement is false according to this observation?

The latent heat of vaporization of water is 540 cal $\mathrm{g}^{-1}$ at $100^{\circ} \mathrm{C}$. What will be the entropy increase when one mole of water is evaporated at $100^{\circ} \mathrm{C}$ ?

Find the approximate value of $(\Delta \mathrm{H}-\Delta \mathrm{U})$ in $\mathrm{J}$. $\mathrm{mol}^{-1}$, for the formation of $\mathrm{CO}$ from its elements at $298 \mathrm{~K} .\left(\mathrm{R}=8.314 . \mathrm{J}\mathrm{~K}^{-1} \cdot \mathrm{mol}^{-1}\right)$

$\Delta \mathrm{G}$ is equal to

Standard enthalpy (heat) of formation of liquid water at $25^{\circ} \mathrm{C}$ is around

$\mathrm{H}_2(\mathrm{~g})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{H}_2 \mathrm{O}(\mathrm{l})$

For the reaction:

$\begin{aligned} & A(\ell) \rightarrow 2 B(g) \\ & \Delta {U}=2.1 \mathrm{~kcal}, \Delta {S}=20 \mathrm{~cal} / \mathrm{K}, \text { at } {T}=300 \mathrm{~K}. \end{aligned}$

Hence $\Delta {G}$ (in $\mathrm{kcal})$ is $-X$. Find the value of $X$.

Work done in the reversible isothermal process is given by

∆G°  for the reaction X+Y⇌Z is -4.606 kcal. The value of equilibrium constant of the reaction at 227℃ is $\left( {R = 2.0\,\,\mathrm{cal}\,\,\mathrm{mo}{\mathrm{l}^{ - 1}}\,{\mathrm{K}^{ - 1}}} \right)$

Consider the reactions given below. On the basis of these reactions find out which of the algebraic relationship given in options is correct?

1. $\mathrm{C(g})+4 \mathrm{H(g}) \longrightarrow \mathrm{CH_4(g}) ; \Delta_r H=x \mathrm{~kJ~mol}^{-1}$ ,

2. $\mathrm{C}$ (graphite, $\mathrm{s})+2 \mathrm{H_2(g}) \longrightarrow \mathrm{CH_4(g}); \Delta_r H=y \mathrm{~kJ~mol}^{-1}$

Assertion(A): Work and internal energy are not state functions.
Reason(R): The sum of $q+w$ is path function.

Calculate the difference between $C_p$ and $C_V$ for 10 moles of an ideal gas.

For which reaction(s) is the $\Delta \mathrm{H}$ value equal to the $\Delta \mathrm{E}$ value?
(I) $\mathrm{CH}_4(\mathrm{g})+2 \mathrm{O}_2(\mathrm{g}) \rightarrow \mathrm{CO}(\mathrm{g})+2 \mathrm{H}_2 \mathrm{O} (l)$
(II) $\mathrm{C}_2 \mathrm{H}_4(\mathrm{g})+3 \mathrm{O}_2(\mathrm{g}) \rightarrow 2 \mathrm{CO}(\mathrm{g})+2 \mathrm{H}_2 \mathrm{O}(\mathrm{g})$