Thermodynamics practice · Mixed difficulty

The carbon dioxide exhaled in the breath of astronaut is often removed from the spacecrafts by reaction with lithium hydroxide.

$2 \mathrm{LiOH}(\mathrm{s})+\mathrm{CO}_2(\mathrm{g}) \rightarrow \mathrm{Li}_2 \mathrm{CO}_3(\mathrm{s})+\mathrm{H}_2 \mathrm{O}({l})$

Assume that each astronaut requires $2100 \mathrm{~kcal}$ of energy per day. Further assume that this energy is obtained only from the combustion of glucose into $\mathrm{CO}_2(\mathrm{g})$ and $\mathrm{H}_2 \mathrm{O}(l).$ The standard enthalpies of formation of glucose, $\mathrm{CO}_2$ and water are $-1280,-395$ and $-285 \mathrm{~kJ} / \mathrm{mol},$ respectively. The minimum mass of $\mathrm{LiOH}$ required per astronaut per day to react completely with all the $\mathrm{CO}_2$ produced is

Which of the following statement is incorrect?

In a chemical reaction ∆H=150 kJ and ∆S=100 $\mathrm{J}{\mathrm{K}^{ - 1}}$ at 300 K. The ∆G for the reaction is:

If one mole of an ideal gas $\left(C_{p, m}=\frac{5}{2} R\right)$ is expanded isothermally at $300 \mathrm{~K}$ until its volume is tripled then $\Delta_{\text {sys }} S$ is

5 moles of an ideal gas at 100 K are allowed to undergo reversible compression till its temperature becomes 200 K. If ${C_V} = 28\mathrm{ }\,\,J{K^{ - 1}}\mathrm{ }mo{l^{ - 1}},$ calculate $\Delta U\mathrm{ }$ and $\Delta pV$ for this process, $(R = 8.0\mathrm{ }\,J{K^{ - 1}}\mathrm{ }mo{l^{ - 1}})$

What is the final temperature (in kelvin) of $0.10$ mole monoatomic ideal gas that performs 75 cal of work adiabatically if the initial temperature is $227^{\circ} \mathrm{C}$? (use $\mathrm{R}=2 \mathrm{~cal} / \mathrm{K}-\mathrm{mol}$)

When a certain amount of ethylene was burnt 6226 kJ heat was evolved. If heat of combustion of ethylene is 1411 kJ, the volume of ${{\mathrm{O}_2}}$ (at NTP) that entered into the reaction is:

Using bond energy data, calculate heat of formation of isoprene. Given $\mathrm{BE}$ of $\mathrm{C}-\mathrm{H}=98.8 \mathrm{~kcal}, \mathrm{H}-\mathrm{H}=104 \mathrm{~kcal}, \mathrm{C}-\mathrm{C}=83 \mathrm{~kcal}, \mathrm{C}=\mathrm{C}=147 \mathrm{~kcal}$ and $\mathrm{C}(\mathrm{s}) \longrightarrow \mathrm{C}(\mathrm{g})=171 \mathrm{~kcal}$.

${FeCO}_3(\mathrm{s})$ decomposes at constant pressure as ${FeO}(\mathrm{s})+{CO}_2(\mathrm{g})$

${FeCO}_3(\mathrm{s}) \stackrel{\Delta}{\longrightarrow} {FeO}(\mathrm{s})+{CO}_2(\mathrm{g})$

at $25^{\circ} \mathrm{C}$, the heat absorbed during the reaction is $80 \mathrm{~kJ}$.

Calculate $\Delta H$ & $\Delta U$ for the reaction

Which of the following pairs of a chemical reaction is certain to result in a spontaneous reaction?

A sample of argon gas at $1 \mathrm{~atm}$ pressure and $27^{\circ} \mathrm{C}$ expands reversibly and adiabatically from $1.25 \mathrm{~dm}^3$ to $2.50 \mathrm{~dm}^3.$ Calculate the enthalpy change in this process. $\mathrm{C}_{\mathrm{v}, \mathrm{m}}$ for argon is $12.48 \mathrm{~JK}^{-1} \mathrm{~mol}^{-1}$

If $\Delta \mathrm{H}$ is the enthalpy change and $\Delta \mathrm{E}$ the change in internal energy accompanying a gaseous reaction then

The standard heat of combustion of a hydrocarbon compound is an/a-

$\Delta H$ for the reaction $2 \mathrm{C}(\mathrm{s})+3 \mathrm{H}_2(\mathrm{g})\rightarrow \mathrm{C}_2 \mathrm{H}_6(\mathrm{g})$ is $-20.24 \mathrm{~kcal} / \mathrm{mol}.$ The bond energies (in $\mathrm{kcal} / \mathrm{mol}$) of $\mathrm{C}-\mathrm{C}, \mathrm{C}-\mathrm{H}$ and $\mathrm{H}-\mathrm{H}$ are $63 , 85.6$ and $102.6 ,$ respectively. The enthalpy of sublimation of $\mathrm{C}(\mathrm{s})$ is

On increasing temperature, the equilibrium constant of exothermic and endothermic reactions repectively

Using the data provided, calculate the multiple bond energy $(kJ\,mo{l^{ - 1}})$of a C≡C bond in ${\mathrm{C}_2}{\mathrm{H}_2}.$ Given that the heat of formation of ${\mathrm{C}_2}{\mathrm{H}_2}$ is $225kJmo{l^{ - 1}}$ (take the bond energy of C─H bond as $350\,\,kJ\,\,mo{l^{ - 1}}.):$
$2\mathrm{C}\left( s \right) \to 2\mathrm{C}\left( \mathrm{g} \right)\Delta H = 1410\,\,\mathrm{kJ\,\,mo}{\mathrm{l}^{ - 1}}$
$2\mathrm{C}\left( s \right) \to 2\mathrm{C}\left( \mathrm{g} \right)\Delta H = 1410\,\,\mathrm{kJ\,\,mo}{\mathrm{l}^{ - 1}}$
${\mathrm{H}_2}\left( \mathrm{g} \right) \to 2\mathrm{H}\left( \mathrm{g} \right)\Delta H = 330\,\,\mathrm{kJ\,\,mo}{\mathrm{l}^{ - 1}}$

1 Mole of ideal gas expanded reversibly and isothermally from 1 lit to 10 lit.
then find out the entropy change in cal/Kelvin.

For a certain reaction the change in enthalpy and change in entropy are $40.63 \mathrm{~kJ} \mathrm{~mol}^{-1}$ and $100 \mathrm{~JK}^{-1}$ respectively. What is the value of $\Delta G$ at $27^{\circ} \mathrm{C}?$

The criterion for a spontaneous process is

When one mole of monoatomic ideal gas at T temperature undergoes adiabatic change under a constant external pressure of 1 atm change in volume is from 1 L to 2 L, the final temperature in Kelvin would be

Two moles of an ideal gas expand spontaneously into vacuum. The work is :-

The value of $\Delta H^\circ$ for the reaction $2 A_{(g)}+B_{(g)} \rightleftarrows A_2 B_{(g)}$ for which $K_p=1.0 \times 10^{-10} \mathrm{~atm}^{-2}$ and $\Delta S^\circ=5 \text{ JK}^{-1}$ and $T=300 \text{ K}$ is

In a fuel cell methanol is used as fuel and oxygen gas is used as an oxidizer. The reaction is $C{H_3}OH\left( l \right) + \frac{3}{2}{O_2}\left( g \right) \to C{O_2}\left( g \right) + 2{H_2}O\left( l \right)$ At 298 K, standard Gibb’s energies of formation for $C{H_3}OH\left( l \right),\,{H_2}O\left( l \right)\,\,and\,\,C{O_2}\left( g \right)$ are -166.2, -237.2 and $- 394.4\mathrm{ }kJ\mathrm{ }mo{l^{ - 1}}$ respectively. If standard enthalpy of combustion of methanol is $- 726kJ\mathrm{ }mo{l^{ - 1}}$ efficiency of the fuel cell will be:

$2.56 \mathrm{~g}$ of sulphur is burnt in a constant volume calorimeter with excess of $\mathrm{O}_2(g)$. Temperature increases from $21.25^{\circ} \mathrm{C}$ to $26.72^{\circ} \mathrm{C}$. The bomb has a heat capacity of $923 \mathrm{~J} \mathrm{~K}^{-1}$. Calorimeter contains $815 \mathrm{~g}$ of water. Thus, change in internal energy per mole of $\mathrm{SO}_2$ formed for the reaction is

$\mathrm{S}_8(s)+8 \mathrm{O}_2(g) \longrightarrow 8 \mathrm{SO}_2(g)$

(Specific heat of water is $4.184 \mathrm{~J} \mathrm{~K}^{-1} \mathrm{~g}^{-1}$ )

Which statement(s) is/are correct?

Decomposition of ozone is exothermic reaction

Standard molar enthalpy of formation at $298 \mathrm{~K}$ is zero for

Given,

$\begin{aligned} & 2 {Fe}_2 {O}_{3(s)} \longrightarrow 4 {Fe}_{(s)}+3 {O}_{2(g)} ; \\ & \Delta_r G_1^{\circ}=+1487 {~kJ} {~mol}^{-1} \\ & 6 {CO}_{(g)}+3 {O}_{2(g)} \longrightarrow {CO}_{2(g)} ; \\ & \Delta_r G_2^{\circ}=-1543.2 {~kJ} {~mol}^{-1} \end{aligned}$

Select the correct statement :

$1 \mathrm{~g}$ ice absorbs $335 \mathrm{~J}$ of heat to melt at $0^{\circ} \mathrm{C}$. The entropy change will be

For the reaction, $2 A(g)+B(g) \longrightarrow 2 D(g) ;$
$\Delta U^{\circ}=-10.5 \mathrm{~kJ} \text { and } \Delta S^{\circ}=-44.1 \mathrm{~J~K}^{-1} .$

The $\Delta G^{\circ}$ for the reaction will be, also predict whether the reaction may occur spontaneously.

$\text {(} R=8.314 \times 10^{-3} \mathrm{~kJ} \mathrm{~K}^{-1} \mathrm{~mol}^{-1} \text { and } T=298 \mathrm{~K} \text {)}$