CBSE Questions for Class 11 Medical Chemistry Thermodynamics Quiz 5 - MCQExams.com

In biological systems flow of energy occurs through :
  • Loss of electrons.
  • Gain of electrons.
  • Both A and B
  • Photons
In which of the processes, does the internal energy of the system remain constant?
  • Adiabatic
  • Isochoric
  • Isobaric
  • Isothermal
$$\Delta { U }^{ o }$$ of combustion of methane is $$-X$$ $$kJ$$ $${ mol }^{ -1 }$$. The value of $$\Delta { H }^{ o }$$ is:
  • $$=\Delta { U }^{ o }$$
  • $$>\Delta { U }^{ o }$$
  • $$<\Delta { U }^{ o }$$
  • $$=0$$
The molar enthalpy change for $$H_2O(1)\rightleftharpoons H2O(g)$$ at $$373$$K and $$1$$ atm is $$41$$kJ/mol. Assuming ideal behavior, the internal energy change for vaporization of $$1$$ mol of water at $$373$$K and $$1$$ atm in kJ $$mol^{-1}$$ is $$?$$
  • $$30.2$$
  • $$41.0$$
  • $$48.1$$
  • $$37.9$$
A mixture of two moles of carbon monoxide and one mole of oxygen, in a closed vessel is ignited to convert the carbon monoxide to carbon dioxide. If $$\Delta$$H is the enthalpy change and $$\Delta$$E is the change in internal energy, then:
  • $$\Delta H > \Delta E$$.
  • $$\Delta H < \Delta E$$.
  • $$\Delta H = \Delta E$$.
  • the relationship depends on the capacity of the vesssel.
"The change of enthalpy of a chemical reaction is the same whether the reaction takes place in one step or in several steps". This law was presented by:
  • Hess
  • La Chatelier
  • Kirchhoff
  • Lavoisier and Laplace
The enthalpies of combustion of $${C}_{(graphite)}$$ and $${C}_{(diamond)}$$ are $$-393.5$$ and $$-395.4kJ/mol$$ respectively. The enthalpy of conversion of $${C}_{(graphite)}$$ to $${C}_{(diamond)}$$ in $$kJ/mol$$ is:
  • $$-1.9$$
  • $$-788.9$$
  • $$1.9$$
  • $$788.9$$
For the reactions,
(i) $${H}_{2}(g)+{Cl}_{2}(g)=2HCl(g)+x$$ $$kJ$$
(ii) $$2HCl(g)={H}_{2}(g)+{Cl}_{2}(g)-y$$ $$kJ$$
which one of the following statements is correct?
  • $$x-y> 0$$
  • $$x-y< 0$$
  • $$x-y=0$$
  • None of these
From the thermochemical reactions,
$${C}_{(graphite)}+\cfrac{1}{2}{O}_{2}\rightarrow CO;\Delta H=-110.5kJ$$
$$CO+\cfrac{1}{2}{O}_{2}\rightarrow {CO}_{2};\Delta H=-283.2kJ$$
the heat of reaction of $${C}_{(graphite)}+{O}_{2}\rightarrow {CO}_{2}$$ is:
  • $$+393.7kJ$$
  • $$-393.7kJ$$
  • $$-172.7kJ$$
  • $$+172.7kJ$$
Spontaneous reactions are :
  • Endergonic
  • Exergonic
  • Energy neutral
  • Exer-endergonic reactions
 For the reactions,
(i) $${H}_{2}(g)+{Cl}_{2}(g)=2HCl(g)+x$$ $$kJ$$

(ii) $${H}_{2}(g)+{Cl}_{2}(g)=2HCl(l)+y$$ $$kJ$$

which one of the following statements is correct?
  • $$x> y$$
  • $$x< y$$
  • $$x-y=0$$
  • $$x=y$$
If $$\Delta {H}_{f}^{o}$$ for $${H}_{2}{O}_{2}(l)$$ and $${H}_{2}O(l)$$ are $$-188kJ$$ $${mol}^{-1}$$ and $$-286kJ$$ $${mol}^{-1}$$, what will be the enthalpy change of the reaction?
$$2{H}_{2}{O}_{2}(l)\rightarrow 2{H}_{2}O(l)+{O}_{2}(g)$$
  • $$146kJ$$ $${mol}^{-1}$$
  • $$-196kJ$$ $${mol}^{-1}$$
  • $$-494kJ$$ $${mol}^{-1}$$
  • $$-98kJ$$ $${mol}^{-1}$$
The heats of combustion of yellow phosphorus and red phosphorus are $$-9.91kJ$$ and $$-8.78kJ$$ respectively. The heat of transition of yellow phosphorus to red phosphorus is:
  • $$-18.69kJ$$
  • $$+1.13kJ$$
  • $$+18.69kJ$$
  • $$-1.13kJ$$
Given:
$$C(s)+{O}_{2}(g)\rightarrow {CO}_{2}(g)+94.2kcal$$
$${H}_{2}(g)+\cfrac{1}{2}{O}_{2}(g)\rightarrow {H}_{2}O(l)+68.3kcal$$
$${CH}_{4}+2{O}_{2}(g)\rightarrow {CO}_{2}(g)+2{H}_{2}O(l)+210.8kcal$$
The heat of formation of methane in $$kcal$$ will be:
  • $$45.9$$
  • $$47.8$$
  • $$20.0$$
  • $$47.3$$
The heats of combustion of rhombic and monoclinic sulphur are $$-70960$$ and $$-71030$$ calorie respectively. What will be the heat of conversion of rhombic sulphur to monoclinic sulphur?
  • $$-70960\ cal$$
  • $$-71030\ cal$$
  • $$+70\ cal$$
  • $$-70\ cal$$
For an exothermic reaction to be spontaneous:
  • temperature must be high
  • temperature must be zero
  • temperature may have any magnitude
  • temperature must be low
For which reaction from the following, $$\Delta S$$ will be maximum?
  • $$Ca(s)+\cfrac{1}{2}{O}_{2}(g)\rightarrow CaO(s)$$
  • $$Ca{CO}_{3}(s)\rightarrow CaO(s)+{CO}_{2}(g)$$
  • $$C(s)+{O}_{2}(g)\rightarrow {CO}_{2}(g)$$
  • $${N}_{2}(g)+{O}_{2}(g)\rightarrow 2NO(g)$$
The value of $$\Delta H$$ and $$\Delta S$$ for a reaction are respectively $$30kJ$$ $${mol}^{-1}$$ and $$100J{K}^{-1}$$ $${mol}^{-1}$$. Then temperature above which the reaction will become spontaneous is:
  • $$300K$$
  • $$30K$$
  • $$100K$$
  • $${300}^{o}C$$
Which of the following thermodynamics relation is correct?
  • $$dG=VdP-SdT$$
  • $$dU=PdV+TdS$$
  • $$dH=-Vdp+TdS$$
  • $$dG=VdP+SdT$$
On combustion carbon forms two oxides $$CO$$ and $${CO}_{2}$$, heat of formation of $${CO}_{2}$$ is $$-94.3\ kcal$$ and that of $$CO$$ is $$-26.0 \ kcal$$. Heat of combustion of carbon is:
  • $$-26.0\ kcal$$
  • $$-68.3\ kcal$$
  • $$-94.3\ kcal$$
  • $$-120.3\ kcal$$
The heat of combustion of ethanol determined in a bomb calorimeter is $$-670.48kcal$$ $${mol}^{-1}$$ at $$298K$$. What is the $$\Delta U$$ at $$298K$$ for the reaction?
  • $$-760\ kcal$$ $${mol}^{-1}$$
  • $$-670.48\ kcal$$ $${mol}^{-1}$$
  • $$+760\ kcal$$ $${mol}^{-1}$$
  • $$+670.48\ kcal$$ $${mol}^{-1}$$
$$\Delta { G }^{ o }$$ for the reaction $$x+y\rightleftharpoons z$$ is $$-4.606kcal$$. The value of equilibrium constant of the reaction at $${227}^{o}C$$ is:
($$R=2.0cal.{mol}^{-1}{K}^{-1}$$)
  • $$100$$
  • $$10$$
  • $$2$$
  • $$1$$
In the conversion of limestone to lime,
$$Ca{CO}_{3}(s)\rightarrow CaO(s)+{CO}_{2}(g)$$
the values of $$\Delta {H}^{o}$$ and $$\Delta {S}^{o}$$ are $$+179.1kJ{ mol }^{ -1 }$$ and $$160.2kJ{ mol }^{ -1 }$$ respectively at $$298K$$ and $$1$$ bar. Assuming that, $$\Delta {H}^{o}$$ and $$\Delta {S}^{o}$$ do not change with temperature; temperature above which conversion of limestone to lime will be spontaneous is:
  • $$1118K$$
  • $$1008K$$
  • $$1200K$$
  • $$845K$$
Considering entropy(s) as a thermodynamic parameter, the criterion for the spontaneity of any process is:
  • $$\Delta {S}_{system}+\Delta {S}_{surroundings}> 0$$
  • $$\Delta {S}_{system}-\Delta {S}_{surroundings}> 0$$
  • $$\Delta {S}_{system}> 0$$ only
  • $$\Delta {S}_{surroundings}> 0 $$ only
For the process,
$${H}_{2}O(l) (1bar,373K)\rightarrow {H}_{2}O(g)(1bar;373K)$$
The correct set of thermodynamics parameters is:
  • $$\Delta G=0,\Delta S=+ve$$
  • $$\Delta G=0,\Delta S=-ve$$
  • $$\Delta G=+ve,\Delta S=0$$
  • $$\Delta G=-ve,\Delta S=+ve$$
Calculate $$\triangle G^{\circ}$$ for the following cell reaction:

$$Zn(s) + Ag_{2}O(s) + H_{2}O(l) \rightleftharpoons Zn^{2+}(aq) + 2Ag(s) + 2OH^{-}(aq)$$

$$E^{\circ}_{Ag^{+}/Ag} = + 0.80V$$ and $$E^{\circ}_{Zn^{2+}/ ZN} = -0.76V$$.
  • $$-305\ kJ/mol$$
  • $$-301\ kJ/mol$$
  • $$+305\ kJ/mol$$
  • $$+301\ kJ/mol$$
Hess law is based on: 
  • Law of conservation of mass
  • Law of conservation of energy
  • Enthalpy is a state function
  • None of these
Which of the following is true for spontaneous process?
  • $$\Delta G> 0$$
  • $$\Delta G< 0$$
  • $$\Delta G= 0$$
  • $$\Delta G=T\Delta S$$
The Haber's process for production of ammonia involves the equilibrium:
$${N}_{2}(g)+3{H}_{2}(g)\rightleftharpoons  2{NH}_{3}(g)$$
Assuming $$\Delta {H}^{o}$$ and $$\Delta {S}^{o}$$ for the reaction do not change with temperature, which of the statements is true?
($$\Delta {H}^{o}=-95kJ$$ and $$\Delta {S}^{o}=-190J{ K }^{ -1 }$$)
  • Ammonia dissociates spontaneously below $$500K$$
  • Ammonia dissociates spontaneously above $$500K$$
  • Ammonia dissociates at all temperatures
  • Ammonia does not dissociate at any temperature
A $$1g$$ sample of substance $$A$$ at $${100}^{o}C$$ is added to $$100mL$$ of $${H}_{2}O$$ at $${25}^{o}C$$. Using separate $$100mL$$ portion of $${H}_{2}O$$ the procedure is repeated with substance $$B$$ then with substance $$C$$. How will the final temperatures of the water compare?
SubstanceSpecific heat
$$A$$$$0.6{g}^{-1}$$ $$^{ o }{ { C }^{ -1 } }$$
$$B$$$$0.4{g}^{-1}$$ $$^{ o }{ { C }^{ -1 } }$$
$$C$$$$0.2{g}^{-1}$$ $$^{ o }{ { C }^{ -1 } }$$
  • $${T}_{C}>{T}_{B}> {T}_{A}$$
  • $${T}_{B}>{T}_{A}> {T}_{C}$$
  • $${T}_{A}>{T}_{B}> {T}_{C}$$
  • $${T}_{A}= {T}_{B}= {T}_{C}$$
$$\Delta H$$ and$$\Delta S$$ for the reaction
$${ Br }_{ 2 }(l)+{ Cl }_{ 2 }(l)\longrightarrow 2BrCl(g)$$
are $$29.37\ kJ$$ and $$104.0\ J{ K }^{ -1 }$$ respectively. Above that temperature will this reaction become spontaneous?
  • T > 177.8 K
  • T > 35.1 K
  • T > 28.4 K
  • T > 141.2 K
Following enthalpy changes are given:

$$\alpha -D\quad glucose(s)\longrightarrow \alpha -D\quad glucose(aq.);\quad \Delta H=10.72kJ$$

$$\beta -D\quad glucose(s)\longrightarrow \beta -D\quad glucose(aq.);\quad \Delta H=4.68kJ$$

$$\alpha -D\quad glucose(aq.)\longrightarrow \beta -D\quad glucose(aq.);\quad \Delta H=1.16kJ$$

Calculate the enthalpy change in,

$$\alpha -D\quad glucose(s)\longrightarrow \beta -D\quad glucose(s)$$
  • $$14.24kJ$$
  • $$16.56kJ$$
  • $$7.2kJ$$
  • $$4.88kJ$$
For a particular reaction, $$\Delta { H }^{ o }=-38.3\ kJ$$ and $$\Delta { S }^{ o }=-113\ J{ K }^{ -1 }{ mol }^{ -1 }$$. This reaction is:
  • spontaneous at all temperatures
  • non-spontaneous at all temperatures
  • spontaneous at temperatures below $${66}^{o}C$$
  • spontaneous at temperatures above $${66}^{o}C$$
Under which circumstances would the free energy change for a reaction be relatively temperature independent?
  • $$\Delta { H }^{ o }$$ is negative
  • $$\Delta { H }^{ o }$$ is positive
  • $$\Delta { S }^{ o }$$ has a large positive value
  • $$\Delta { S }^{ o }$$ has a small magnitude
Given that:
$$2C(s)+2{ O }_{ 2 }(g)\longrightarrow 2C{ O }_{ 2 }(g);\Delta H=-787kJ...(i)$$
$${ H }_{ 2 }(g)+\cfrac { 1 }{ 2 } { O }_{ 2 }(g)\longrightarrow { H }_{ 2 }O(l);\Delta H=-286kJ...(ii)$$
$${ C }_{ 2 }{ H }_{ 2 }(g)+\cfrac { 5 }{ 2 } { O }_{ 2 }(g)\longrightarrow 2C{ O }_{ 2 }(g)+3{ H }_{ 2 }O(l)....(iii)$$
$$\Delta H=-1301kJ$$
Heat formation of acetylene is:
  • $$-1802kJ$$
  • $$+1802kJ$$
  • $$-800kJ$$
  • $$+228kJ$$
Which is correct about $$\Delta G$$?
  • $$\Delta G=\Delta H-T\Delta S$$
  • At equilibrium $$\Delta { G }^{ o }=0$$
  • At eq. $$\Delta G=-RT\log{K}$$
  • $$\Delta G=\Delta { G }^{ o }+RT\log{K}$$
The enthalpy changes for two reactions are given by the equations:
$$2Cr(s)+1\cfrac { 1 }{ 2 } { O }_{ 2 }(g)\longrightarrow { Cr }_{ 2 }{ O }_{ 3 }(s);\Delta H=-1130kJ$$
$$C(s)+\cfrac { 1 }{ 2 } { O }_{ 2 }(g)\longrightarrow CO(g);\Delta H=-110kJ$$
What is the enthalpy change, in $$kJ$$, for the reaction?
$$3C(s)+{ Cr }_{ 2 }{ O }_{ 3 }(s)\longrightarrow 2Cr(s)+3CO(g)$$
  • $$-1460kJ$$
  • $$-800kJ$$
  • $$+800kJ$$
  • $$+1020kJ$$
  • $$+1460kJ$$
The standard Gibbs free energy $$\Delta { G }^{ o }$$ is related to equilibrium constant $${K}_{P}$$ as:
  • $${K}_{P}=-RT\log{\Delta { G }^{ o }}$$
  • $${K}_{P}={[e/RT]}^{\Delta { G }^{ o }}$$
  • $${K}_{P}=-\Delta { G }^{ o }/RT$$
  • $${K}_{P}={e}^{-\Delta { G }^{ o }/RT}$$
Hess's law of constant heat summation in based on: 
  • $$E=mc^{2}$$
  • Conservation of mass
  • First law of thermodynamics
  • None of the above
Hess law of heat summation includes 
  • Initial reactants only
  • Initial reactants and final products
  • Final products only
  • Intermediates only
The enthalpy change for two reactions are given by the equations:
$$2Cr(s)+\cfrac { 3 }{ 2 } { O }_{ 2 }(g)\longrightarrow { Cr }_{ 2 }{ O }_{ 3 }(s);\Delta H=-1130kJ$$
$$C(s)+\cfrac { 1 }{ 2 } { O }_{ 2 }(g)\longrightarrow CO(g);\Delta H=-110kJ$$
What is the enthalpy change in $$kJ$$ for the following reaction?
$$3C(s)+{ Cr }_{ 2 }{ O }_{ 3 }(s)\longrightarrow 2Cr(s)+3CO(g)$$
  • $$-1460kJ$$
  • $$-1800kJ$$
  • $$+800kJ$$
  • $$-1020kJ$$
  • $$+1460kJ$$
Given that:
$$2Fe(s)+\cfrac { 3 }{ 2 } { O }_{ 2 }(g)\longrightarrow { Fe }_{ 2 }{ O }_{ 3 }(s);\Delta H=-193.4kJ...(i)$$
$$Mg(s)+\cfrac { 1 }{ 2 } { O }_{ 2 }(g)\longrightarrow MgO(s);\Delta H=-140.2kJ....(ii)$$
What is $$\Delta H$$ of the reaction?
$$3Mg+{ Fe }_{ 2 }{ O }_{ 3 }\longrightarrow 3MgO+2Fe$$
  • $$-227.2kJ$$
  • $$-237.3kJ$$
  • $$2227.2kJ$$
  • $$-257.3kJ$$

 A 100.0g ice cube at 0.0°C is placed in 650g of water at 25°C. What is the final temperature of the mixture?
  • $$10^0C$$
  • $$11^0C$$
  • $$12^0C$$
  • $$13^0C$$
For spontaneous process:
  • $$\Delta { S }_{ total }=0$$
  • $$\Delta { S }_{ total }>0$$
  • $$\Delta { S }_{ total }<0$$
  • none of these
Reaction of silica with mineral acids may be given as:
$$Si{ O }_{ 2 }+4HF\longrightarrow Si{ F }_{ 4 }+2{ H }_{ 2 }O;\Delta H=-10.17kcal$$
$$Si{ O }_{ 2 }+4HCl\longrightarrow Si{ Cl }_{ 4 }+2{ H }_{ 2 }O;\Delta H=+36.7kcalHCl$$
Which among the following is correct?
  • $$HF$$ and $$HCl$$ both will react with silica
  • Only $$HF$$ will react with silica
  • Only $$HCl$$ will react with silica
  • Neither $$HF$$ nor $$HCl$$ will react with silica
Calculate the heat of formation of methane, given that
heat of formation of water $$=-286kJ\quad { mol }^{ -1 }$$
heat of combustion of methane $$=-890kJ\quad { mol }^{ -1 }$$
heat of combustion of carbon $$=-393.5kJ\quad { mol }^{ -1 }$$
  • $$75.5\ kJ\ { mol }^{ -1 }$$
  • $$-75.5\ kJ\ { mol }^{ -1 }$$
  • $$-55.5\ kJ\ { mol }^{ -1 }$$
  • $$55.5\ kJ\ { mol }^{ -1 }$$
From the following data of heats of combustion, find the heat of formation of $${ CH }_{ 3 }OH(l)$$:
$${ CH }_{ 3 }OH(l)+\cfrac { 3 }{ 2 } { O }_{ 2 }(g)\longrightarrow C{ O }_{ 2 }(g)+2{ H }_{ 2 }O(l);\Delta H=-726kJ$$
$$C(s)+{ O }_{ 2 }(g)\longrightarrow C{ O }_{ 2 }(g);\Delta H=-394kJ$$
$${ H }_{ 2 }(g)+\cfrac { 1 }{ 2 } { O }_{ 2 }(g)\longrightarrow { H }_{ 2 }O(l);\Delta H=-286kJ$$
  • $$240\ kJ\ { mol }^{ -1 }$$
  • $$-240\ kJ\ { mol }^{ -1 }$$
  • $$-140\ kJ\ { mol }^{ -1 }$$
  • $$140\ kJ\ { mol }^{ -1 }$$
The free energy for a reaction having $$\Delta H=31400\ cal,\Delta S=32\ cal\ { K }^{ -1 }{mol}^{-1}$$ at $${1000}^{o}C$$ is:
  • $$-9336\ cal$$
  • $$-7386\ cal$$
  • $$-1936\ cal$$
  • $$+9336\ cal$$
The most abundant element in the universe is thought to be:
  • Hydrogen
  • Carbon
  • Oxygen
  • Nitrogen
A spontaneous process may be defined as :
  • A process which is exothermic and evolves a lot of heat
  • A process which is slow and reversible
  • A process which takes place only in presence of a catalyst
  • A process that occurs without any input from the surroundings
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