CBSE Questions for Class 11 Medical Chemistry Thermodynamics Quiz 11 - MCQExams.com

Hess's law is used to calculate:
  • Enthalpy of reaction
  • Entropy of reaction
  • Work done in reaction
  • All the above
A sample of liquid in a thermally insulated container (a calorimeter) is stirred for 2 hr by a mechanical linkage to a motor in the surrounding. For this process:
  • $$w < 0; q = 0; \Delta U = 0$$
  • $$w > 0; q < 0; \Delta U > 0$$
  • $$w < 0; q > 0; \Delta U = 0$$
  • $$w > 0; q = 0; \Delta U > 0$$
The temperature in K at which $$\displaystyle \Delta G=0$$, for a given reaction with $$\displaystyle \Delta H=-20.5\ kJ{ mol }^{ -1 }$$ and $$\displaystyle \Delta S=-50.0\ J{ K }^{ -1 }{ mol }^{ -1 }$$ is:
  • -410
  • 410
  • 2.44
  • -2.44
Calculate $${ \Delta H }/{ kJ }$$ for the following reaction using the listed standard enthalpy of reaction data:
                       $$2{ N }_{ 2 }\left( g \right) +5{ O }_{ 2 }\left( g \right) \longrightarrow 2{ N }_{ 2 }{ O }_{ 5 }\left( s \right) $$

$${ N }_{ 2 }\left( g \right) +3{ O }_{ 2 }\left( g \right) +{ H }_{ 2 }\left( g \right) \longrightarrow 2HN{ O }_{ 3 }\left( aq \right)                       { \Delta H }/{ kJ }=-414.0$$

$${ N }_{ 2 }{ O }_{ 5 }\left( s \right) +{ H }_{ 2 }O\left( I \right) \longrightarrow 2HN{ O }_{ 3 }\left( aq \right)                             { \Delta H }/{ kJ }=-86.0$$

$$2{ H }_{ 2 }\left( g \right) +{ O }_{ 2 }\left( g \right) \longrightarrow 2{ H }_{ 2 }O\left( I \right)                               { \Delta H }/{ kJ }=-571.6$$
  • $$-84.4$$
  • $$-243.6$$
  • $$-71.2$$
  • $$-121.8$$
What can be used in combination with a calorimeter to compare the specific heats of two substances?
  • Thermometer
  • Conductivity tester
  • Graduated cylinder
  • Buret
  • Salt bridge
The equilibrium constant $$\left( K \right) $$ of a reaction may be written as:
  • $$K={ e }^{ { \Delta G }/{ RT } }$$
  • $$K={ e }^{ { -\Delta { G }^{ } }/{ RT } }$$
  • $$K={ e }^{ { \Delta H }/{ RT } }$$
  • $$K={ e }^{ -{ \Delta { H }^{ } }/{ RT } }$$
$$H_2(g) + \frac{ 1 }{ 2 }O_2(g)\longrightarrow 2H_2O(l);   \Delta H = -86  kJ$$
$$2H_2(g) + O_2(g)\longrightarrow 2H_2O(l)......... kJ(\pm?)$$
  • $$2H_2(g) + O_2(g)\longrightarrow 2H_2O(l) + 172 kJ$$.
  • $$2H_2(g) + O_2(g)\longrightarrow 2H_2O(l) - 172 kJ$$.
  • $$2H_2(g) + O_2(g)\longrightarrow 2H_2O(l) +0 kJ$$.
  • None of these
Based on the following thermochemical equations
$${H}_{2}(g)+C(s)\longrightarrow  CO(g)+{H}_{2}(g)$$; $$\Delta H=133kJ {mol}^{-1}$$
$$CO(g)+\cfrac{1}{2}{O}_{2}(g)\longrightarrow  {CO}_{2}(g)$$; $$\Delta H=-282kJ{mol}^{-1}$$
$${H}_{2}(g)+\cfrac{1}{2}{O}_{2}(g)\longrightarrow  {H}_{2}O(g)$$;  $$\Delta H=-242kJ{mol}^{-1}$$
$$C(s)+{O}_{2}(g)\longrightarrow  {CO}_{2}(g)$$; $$\Delta H=xkJ{mol}^{-1}$$
The value of $$x$$ will be:
  • $$393.0$$
  • $$655.0$$
  • $$-393.0$$
  • $$-655.0$$
I : $$\displaystyle { S }_{ 8 }\left( s \right) +8{ O }_{ 2 }\left( g \right) \rightarrow 8{ SO }_{ 2 }\left( g \right) \Delta H=-2374.6kJ$$
II : $$\displaystyle { S }_{ 8 }\left( s \right) +12{ O }_{ 2 }\left( g \right) \rightarrow 8{ SO }_{ 3 }\left( g \right) \Delta H=-3165.8.6kJ$$
From the information given above, determine the heat of reaction for the combustion of sulfur dioxide.
$$\displaystyle 2{ SO }_{ 2 }\left( g \right) +{ O }_{ 2 }\left( g \right) \rightarrow 2{ SO }_{ 3 }\left( g \right) $$
  • -5540.4 kJ
  • -1385.1 kJ
  • -791.2 kJ
  • -197.8 kJ
  • -899.2 kJ
Bond Average Bond Energy (kJ/mol)
 $$C \equiv O$$ 1075
 $$C=O$$ 728
 $$C-Cl$$ 326
 $$Cl-Cl$$ 243
From the above given data, calculate the heat of reaction for the following reaction :

$$CO+{ Cl }_{ 2 }\rightarrow CO{ Cl }_{ 2 }$$
  • $$+62$$ kJ
  • $$-62$$ kJ
  • $$-409$$ kJ
  • $$+706$$ kJ
Name the apparatus used to measure the heat absorbed or released by a reaction.
  • Centrifuge
  • Barometer
  • Balance
  • Calorimeter
  • Battery
Statement 1: The $$\Delta H_{reaction}$$ of a particular reaction can be arrived at by the summation of the $$\Delta H_{reaction}$$ values of two or more reactions that, added together, will give the $$\Delta H_{reaction}$$ of the particular reaction.
Statement 2: Hess's Law conforms to the First Law of Thermodynamics, which states that the total energy of the universe is a constant.

  • Both Statement 1 and Statement 2 are correct and Statement 2 is the correct explanation of Statement 1.

  • Both Statement 1 and Statement 2 are correct and Statement 2 is not the correct explanation of Statement 1.

  • Statement 1 is correct but Statement 2 is not correct.

  • Statement 1 is not correct but Statement 2 is correct.

  • Both the Statement 1 and Statement 2 are not correct.

Calculate the enthalpy , $$\Delta  H$$, for the given reaction using the the given bond energies ?
$$C_{2}H_{4} + Cl_{2}\rightarrow ClH_{2}C-CH_{2}Cl$$
Bond energieskJ/ mol
$$C-C$$$$347$$
$$C = C$$$$612$$
$$C-Cl$$$$341$$
$$C-H$$$$414$$
$$Cl-Cl$$$$243$$
  • $$\Delta H = -800 kJ$$
  • $$\Delta H = -680 kJ$$
  • $$\Delta H = -174 kJ$$
  • $$\Delta H = +174 kJ$$
  • $$\Delta H = +200 kJ$$
The spontaneity of a reaction is indicated by:
  • enthalpy change
  • entropy change
  • gibbs free energy change
  • activation energy
  • specific heat capacity
At the standard temperature and pressure (STP) of $$0$$ $$^oC$$ and $$1$$ atmosphere, the Gibbs free energy change is  ______.
  • $$-247 kJ/mol.$$
  • $$-257 kJ/mol.$$
  • $$-237 kJ/mol.$$
  • $$-227 kJ/mol.$$
The reverse of a spontaneous reaction is ......... .
  • always spontaneous
  • always non spontaneous
  • sometimes spontaneous
  • sometimes non spontaneous
  • There is no way of telling
The important considerations in deciding if a reaction will be spontaneous are :
  • stability & state of reactants
  • energy gained & heat evolved
  • exothermic energy & randomness of the products
  • endothermic energy & randomness of the products
  • endothermic energy & structure of the products
What is the value of $$\Delta H$$ for the reaction $$X + 2Y \rightarrow 2Z$$?
$$ W+X\rightarrow  2Y$$ ;                                  $$\Delta H=-  400 $$ kcal/mol
$$ 2W + 3X \rightarrow  2Z + 2Y$$  ;                  $$\Delta H=- 150 $$ kcal/mol 
  • $$- 550$$ kcal/mol
  • $$+50$$ kcal/mol
  • $$- 50$$ kcal/mol
  • $$+ 650$$ kcal/mol
  • $$+250$$ kcal/mol
If $$\Delta G$$ standard is zero, this means :
  • the reaction is spontaneous at standard conditions
  • the reaction is non spontaneous at standard conditions
  • the system is at equilibrium at standard conditions
  • the reaction is both non spontaneous and at equilibrium
  • the reaction is both spontaneous and at equilibrium
Which of the following drives spontaneous reactions?
  • Low enthalpy values and high entropy values.
  • Low enthalpy values and low entropy values.
  • High enthalpy values and low entropy values.
  • High enthalpy values and high entropy values.
  • High temperatures and low pressures.
From the heats of reaction of these individual reactions:
$$A + B \rightarrow 2C$$      $$\triangle H = -500 kJ$$
$$D + 2B \rightarrow E$$      $$\triangle H = -700 kJ$$
$$2D + 2A \rightarrow F$$    $$\triangle H = +50 kJ$$
Find the heat of reaction for $$F + 6B$$
$$\rightarrow 2E + 4C$$
  • +450 kJ
  • -1100 kJ
  • +2350 kJ
  • -350 kJ
  • -2450 kJ
For a reaction to be spontaneous, the sign on delta G should be :
  • positive
  • there should be no sign
  • negative
  • spontaneity is not related to Gibbs Free Energy
  • positive or Negative
Which of the following is most likely to produce a spontaneous reaction?
  • Negative Enthalpy
  • Positive Enthalpy
  • Negative Entropy
  • Positive Entropy
  • Negative Enthalpy and positive Entropy
If the gibbs free energy is negative than reaction will be?
  • always positive
  • sometimes negative
  • non-spontaneous 
  • spontaneous 
  • None of the above
Determine $$\Delta G^{o}$$  for the following reaction.

 $$CO(g)+\cfrac{1}{2}O_2(g)\rightarrow CO_2(g)$$; $$\Delta H^{o} = -282.84 kJ$$

[Given: $$S^{o} _{CO_2}=213.8$$ JK$$^{-1}$$mol$$^{-1}$$, $$S^{o}\ _{ CO(g)} = 197.9$$ J K$$^{-1}$$mol$$^{-1}$$ $$S^{o} \ _{O_2}=205.8$$ J K$$^{-1}$$mol$$^{-1}$$]
  • $$-157.33$$ kJ
  • $$+201.033$$ kJ
  • $$-256.91$$ kJ
  • $$+257.033$$ kJ
An ideal mono-atomic gas of given mass is heated at constant pressure. In this process, the fraction of supplied energy used for the increase of the internal energy of the gas is
  • $$\dfrac{3}{8}$$
  • $$\dfrac{3}{5}$$
  • $$\dfrac{3}{4}$$
  • $$\dfrac{2}{5}$$
How much energy is needed to convert $$100\ g$$ of ice at $$263\ K$$ to liquid water at a temperature of $$283\ K$$?
$${C}_{ice}=0.49\ cal/({g}^{o}C)$$
$${C}_{water}=1.00\ cal/({g}^{o}C)$$
$$\Delta {H}_{fus}= 79.8\ cal/{g}$$
$$\Delta {H}_{vap}=540\ cal/g$$
  • $$9470\ cal$$
  • $$3288\ cal$$
  • $$2288\ cal$$
  • $$727.3\ cal$$
$$H_2(g)+Cl_2(g)=2HCl(g)$$;
$$\Delta H(298K)=-22.06$$kcal. For this reaction, $$\Delta U$$ is equal to:
  • $$-22.06+2\times 10^{-3}\times 298\times 2$$kcal
  • $$-22.06+2\times 298$$ kcal
  • $$-22.06-2\times 298\times 4$$kcal
  • $$-22.06$$ kcal
$$Mg(s) \rightarrow \rightarrow \rightarrow Mg^{2+}$$ ; Energy = A KJ/mol.
$$Cl_{2}(g) \rightarrow \rightarrow \rightarrow 2Cl^{-}$$ ; Energy = B KJ/mol.
What is the $$\triangle H_{f}$$ of $$MgCl_{2}$$ if lattice enthalpy involved is C KJ/mol ?
  • $$A+B+C$$ KJ/mol.
  • $$A-B+C$$ KJ/mol.
  • $$A+B-C$$ KJ/mol.
  • $$-(A+B+C)$$ KJ/mol.
What is the necessary condition for the sponataneity of a process?
  • $$\Delta S>0$$
  • $$\Delta E<0$$
  • $$\Delta H<0$$
  • $$\Delta G<0$$
The equilibrium constant of a reaction is $$0.008$$ at $$298$$ K. The standard free energy change of the reaction at the same temperature is :
  • $$+11.96 \ kJ$$
  • $$-11.96\ kJ$$
  • $$-5.43 \ kJ$$
  • $$-8.46 \ kJ$$
Which of the following conditions are true about spontaneous process?
  • $$\Delta (G_{system})_{T,P}<0$$
  • $$\Delta S_{system} + \Delta S_{Surroundings} >0$$
  • $$\Delta (G_{system})_{T,P}=0$$
  • $$\Delta S_{system} + \Delta S_{Surroundings} <0$$
The process is spontaneous at the given temperature, if:
  • $$\Delta H$$ is $$+ve$$ and $$\Delta S$$ is $$-ve$$
  • $$\Delta H$$ is $$-ve$$ and $$\Delta S$$ is $$+ve$$
  • $$\Delta H$$ is $$+ve$$ and $$\Delta S$$ is $$+ve$$
  • $$\Delta H$$ is $$+ve$$ and $$\Delta S$$ is equal to zero
Calculate enthalpy for formation of ethylene from the following data:

(I) $$C(graphite) + O_2 (g) \rightarrow CO_2 (g); \ \ \ \Delta H = -393.5 kJ$$
(II) $$H_2(g) + \dfrac{1}{2} O_2 (g) \rightarrow H_2O(l); \ \ \ \ \Delta H = - 286.2 kJ$$
(III) $$C_2H_4(g) + 3O_2(g) \rightarrow 2CO_2(g)  + 2H_2 O(l); \ \ \ \   \Delta H = - 1410.8 kJ$$
  • 54.1 kJ
  • 44.8 kJ
  • 51.4 kJ
  • 48.4 kJ
At $$1000\ K$$, from the data :
$$N_{2}(g) + 3H_{2}(g) \rightarrow 2NH_{3}(g)$$; $$\triangle H = -123.77\ kJ\ mol^{-1}$$

Substance$$N_{2}$$$$H_{2}$$$$NH_{3}$$
$$P/R$$$$3.5$$$$3.5$$$$4$$

Calculate the heat of formation of $$NH_{3}$$ at $$300\ K$$.
  • $$-44.42\ kJ\ mol^{-1}$$
  • $$-88.85\ kJ\ mol^{-1}$$
  • $$+ 44.42\ kJ\ mol^{-1}$$
  • $$+ 88.85\ kJ\ mol^{-1}$$
For $$A\rightarrow B,\, \Delta H=4\,\text{kcal mol}^{-1}$$, $$\Delta S=10 \,\text{cal mol}^{-1}K^{-1}$$. Reaction is spontaneous when temperature is:
  • $$400$$K
  • $$300$$K
  • $$500$$K
  • None of these
For spontaneity of cell, which is correct?
  • $$\Delta G=0, \ \Delta H=0$$
  • $$\Delta G=-ve, \ \Delta H=0$$
  • $$\Delta G=+ve,\ \Delta H=0$$
  • $$\Delta G=-ve$$
Given that the bond energies of $$:N\equiv N$$ is $$946$$ kJ $$mol^{-1}$$ $$H-H$$ is $$435 $$ kJ $$mol^{-1}$$, $$N-N$$ is $$159$$ kJ $$mol^{-1}$$, and $$N-H$$ is $$389$$ kJ $$mol^{-1}$$, the heat of formation of hydrazine in the gas phase in kJ $$mol^{-1}$$ is:
  • $$833$$
  • $$101$$
  • $$334$$
  • $$1264$$
A heat engine operating between 227 deg C and 27 deg C absorbs 1 kcal of heat from the 227 deg C reservoir per cycle. Calculate
(1) the amount of heat discharged into the low temperature reservoir.
(2) the amount of work done per cycle.
(3) the efficiency of cycle.
  • 0.4 kcal, 0.6 kcal, 40%
  • 0.6 kcal, 0.4 kcal, 40%
  • 0.4 kcal, 0.6 kcal, 60%
  • 0.7 kcal, 0.4 kcal, 40%
For a system in equilibrium, $$\Delta G = 0$$, under conditions of constant ________.
  • temperature and pressure
  • temperature and volume
  • pressure and volume
  • energy and volume
Given, 
$$C(s)+{O}_{2}(g)\rightarrow {CO}_{2}(g); \Delta H=-395\ kJ$$
$$S(s)+{O}_{2}(g)\rightarrow {SO}_{2}(g);\Delta H=-295\ kJ$$
$$C{S}_{2}(l)+3{O}_{2}(g)\rightarrow {CO}_{2}(g)+2{SO}_{2}(g);\Delta H=-1110\ kJ$$

The heat of formation of $$C{S}_{2}(l)$$ is:
  • $$250\ kJ$$
  • $$62.5\ kJ$$
  • $$31.25\ kJ$$
  • $$125\ kJ$$
What will be the heat formation of methane; if the heat of combustion of carbon is '$$-x$$' $$kJ$$, heat of formation of water is  '$$-y$$' $$kJ$$ and heat of combustion of methane is  '$$-z$$' $$kJ$$?
  • $$(-x-y-z)kJ$$
  • $$(-z-x+2y)kJ$$
  • $$(-x-2y-z)kJ$$
  • $$(-x-2y+z)kJ$$
Which one of the following is not applicable for a thermochemical equation?
  • It tell about physical state of reactants and products
  • It tells whether the reaction is spontaneous
  • It tells whether the reaction is exothermic or endothermic
  • It tells about the allotropic form (if any) of the reactants
Super cooled water is liquid water that has been cooled below its normal freezing point. This state is thermodynamically :
  • unstable and tends to freeze into ice spontaneously
  • stable and tends to freeze Into ice spontaneously
  • stable and tends to fuse into liquid spontaneously
  • unstable and tends to fuse into liquid spontaneously
If $${H}_{2}+\cfrac{1}{2}{O}_{2}\rightarrow {H}_{2}O;\Delta H=-68.09kcal$$
$$K+{H}_{2}O+water \rightarrow KOH(aq.)+\cfrac{1}{2}{H}_{2};\Delta H=-48.0kcal$$
$$KOH+water\rightarrow KOH(aq); \Delta H=-14.0kcal$$
the heat of formation of $$KOH$$ is:
  • $$-68.39+48-14.0$$
  • $$-68.39-48.0+14.0$$
  • $$+68.39-48.0+14.0$$
  • $$+68.39+48.0-14.0$$
From the following reactions at $$298\ K$$.

(A) $$CaC_{2}(s) + 2H_{2}O(l) \rightarrow Ca(OH)_{2}(s) + C_{2}H_{2} (g);\ \Delta H^{\circ}= - 127.9\ kJ\ mol^{-1}$$

(B) $$Ca(s) + \dfrac {1}{2} O_{2}(g) \rightarrow CaO(s) ;\ \Delta H=- 635.1kJ\ mol^{-1}$$

(C) $$CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(s);\ \Delta H=- 65.2\ kJ\ mol^{-1}$$

(D) $$C(s) + O_{2}(s) \rightarrow CO_{2}(s) ;\ \Delta H=- 393.5\ kJ\ mol^{-1}$$

(E) $$C_{2}H_{2}(g) + \dfrac {5}{2}O_{2}(g) \rightarrow 2CO_{2}(g) + H_{2}O(l);\ \Delta H= - 1299.58\ kJ\ mol^{-1}$$

Calculate the heat of formation of $$CaC_{2}(s)$$ at $$298\ K$$.
  • $$-59.82\ kJ\ mol^{-1}$$
  • $$+59.82\ kJ\ mol^{-1}$$
  • $$-190.22\ kJ\ mol^{-1}$$
  • $$+190.22\ kJ\ mol^{-1}$$
Standard heats of formation for $$C{Cl}_{4},{H}_{2}O,{CO}_{2}$$ and $$HCl$$ at $$298K$$ are $$-25.5,-57.8,-94.1$$ and $$-22.1kJ/mol$$ respectively.
For the reaction, what will be $$\Delta H$$?
$$C{Cl}_{4}+2{H}_{2}O\rightarrow {CO}_{2}+4HCl$$
  • $$36.4\ kJ$$
  • $$20.7\ kJ$$
  • $$-20.7\ kJ$$
  • $$-41.4\ kJ$$
$$S+{O}_{2}\rightarrow {SO}_{2}; \Delta H=-298.2kJ$$
$${ SO }_{ 2 }+\cfrac { 1 }{ 2 }{O}_{2} \rightarrow { SO_3 }; \Delta H=-98.7kJ$$
$${SO}_{3}+{H}_{2}O\rightarrow {H}_{2}{SO}_{4};\Delta H=-130.2kJ$$
$${H}_{2}+\cfrac { 1 }{ 2 } { O }_{ 2 }\rightarrow {H}_{2}O;\Delta H=-227.3kJ$$
The heat of formation of $${H}_{2}{SO}_{4}$$ will be:
  • $$-754.4kJ$$
  • $$+320.5kJ$$
  • $$-650.3kJ$$
  • $$-433.7kJ$$
$$\Delta {H}_{f(x)},\Delta {H}_{f(y)},\Delta {H}_{f(R)}$$ and $$\Delta {H}_{f(S)}$$ denote the enthalpies of formation of $$x,y,R$$ and $$S$$ respectively. The  enthalpy of the reaction $$x+y\rightarrow R+S$$ is:
  • $$\Delta {H}_{f(x)}+\Delta {H}_{f(y)}$$
  • $$\Delta {H}_{f(R)}+\Delta {H}_{f(S)}$$
  • $$\Delta {H}_{f(x)}+\Delta {H}_{f(y)}-\Delta {H}_{f(R)}-\Delta {H}_{f(S)}$$
  • $$\Delta {H}_{f(R)}+\Delta {H}_{f(S)}-\Delta {H}_{f(x)}-\Delta {H}_{f(y)}$$
What are the most favourable conditions for the reaction;
$${ SO }_{ 2 }(g)+\cfrac { 1 }{ 2 } { O }_{ 2 }(g)\leftrightharpoons { SO }_{ 3 }(g);\Delta { H }^{ o }=-ve$$ to occur?
  • Low temp and high press
  • Low temp and low press
  • High temp and low press
  • High temp and high press
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