NCERT Solutions For Class 11 Chemistry Chapter 6

NCERT Solutions Class 11 Chemistry Chemical Thermodynamics

NCERT Solutions For Class 11 Chemistry Chapter 6 is provided here. This topic is an extremely important topic in chemistry and is important for class 11, Class 12 and for competitive exams like JEE and NEET. Students must have a good knowledge of the topic in order to excel in the examination. We at BYJU'S provide the NCERT Solutions For Class 11 Chemistry Chemical Thermodynamics PDF which students can download. Practicing all the questions will be very helpful for the students as many questions are framed from the topic.

NCERT Solutions for Class 11 Chemistry Chapter 6 includes the topic - Thermodynamic Terms: The system and the surroundings, Types of System: Open, Closed, Isolated. The State of the system. The Internal energy as a state Function: Work, Heat, The general case. Applications: Work, Isothermal and free expansion of an Ideal Gas. Enthalpy, Extensive and Intensive Properties, Heat capacity, The relation between Cp and Cv for an ideal gas. Measurement of ∆U and ∆H Calorimetry. Gibbs energy change and equilibrium, Enthalpy Change. Enthalpies for different types of reactions. Lattice enthalpy. Spontaneity. NCERT Solutions Class 11 Chemistry Chemical Thermodynamics PDF is provided here for better understanding and clarification of the chapter.

Thermodynamics is a branch of science. It deals with the relationship between heat and other forms of energy. A part of the universe where observations are made is called system. Surrounding is part of universe excluding system. Based on the exchange of energy and matter, system is classified in 3 types: Closed system, Open system, Isolated system. In closed system only energy can be exchanged with the surrounding. In open system energy as well as matter can be exchanged with the surrounding. In isolated system both energy and matter cannot be exchanged with the surrounding. This was a brief on Chemical Thermodynamics.

Q-1: A thermodynamic state function is _______.

1. A quantity which depends upon temperature only.

2. A quantity which determines pressure-volume work.

3. A quantity which is independent of path.

4. A quantity which determines heat changes.

Ans:

(3) A quantity which is independent of path.

Reason:

Functions like pressure, volume and temperature depends on the state of the system only and not on the path.

Q-2: Which of the following is a correct conditions for adiabatic condition to occur.

1. q = 0

2. w = 0

3. Δp=0$\Delta p = 0$

4. ΔT=0$\Delta T = 0$

Ans:

1: q = 0

Reason:

For an adiabatic process heat transfer is zero, i.e. q = 0.

Q-3: The value of enthalpy for all elements in standard state is _____.

(1) Zero

(2) < 0

(3) Different for every element

(4) Unity

Ans:

(1) Zero

Q-4: For combustion of methane ΔUΘ$\Delta U^{\Theta }$ is – Y kJmol1$kJ mol^{-1}$. Then value of ΔHΘ$\Delta H^{\Theta }$ is ____.

(a) > ΔUΘ$\Delta U^{\Theta }$

(b) = ΔUΘ$\Delta U^{\Theta }$

(c) = 0

(d) < ΔUΘ$\Delta U^{\Theta }$

Ans:

(d) < ΔUΘ$\Delta U^{\Theta }$

Reason:

ΔHΘ=ΔUΘ+ΔngRT$\Delta H^{\Theta } = \Delta U^{\Theta } + \Delta n_{g}RT$ ; ΔUΘ$\Delta U^{\Theta }$ = – Y kJmol1$kJ mol^{-1}$,

ΔHΘ=(Y)+ΔngRTΔHΘ<ΔUΘ$\Delta H^{\Theta } = ( – Y) + \Delta n_{g}RT \Rightarrow \Delta H^{\Theta } < \Delta U^{\Theta }$

Q-5: For, Methane, di-hydrogen and and graphite the enthalpy of combustion at 298K are given -890.3kJ mol1$mol^{-1}$, -285.8kJmol1$mol^{-1}$ and -393.5kJmol1$mol^{-1}$ respectively. Find the enthalpy of formation of Methane gas?

(a) -52.27kJmol1$mol^{-1}$

(b) 52kJmol1$mol^{-1}$

(c) +74.8kJmol1$mol^{-1}$

(d) -74.8kJmol1$mol^{-1}$

Ans:

(d) -74.8kJmol1$mol^{-1}$

à CH4(g)  + 2O2(g) $\rightarrow$ CO2(g) + 2H2O(g)

ΔH=890.3kJmol1$\Delta H = -890.3kJmol^{-1}$

à C(s) + O2(g)  $\rightarrow$ CO2(g)

ΔH=393.5kJmol1$\Delta H = -393.5kJmol^{-1}$

à 2H2(g) + O2(g) $\rightarrow$ 2H2O(g)

ΔH=285.8kJmol1$\Delta H = -285.8kJmol^{-1}$

à C(s) + 2H2(g) $\rightarrow$ CH4(g)

ΔfHCH4$\Delta _{f}H_{CH_{4}}$  = ΔcHc$\Delta _{c}H_{c}$ + 2ΔfHH2$2\Delta _{f}H_{H_{2}}$ΔfHCO2$\Delta _{f}H_{CO_{2}}$

= [ -393.5 +2(-285.8) – (-890.3)] kJmol1$mol^{-1}$

= -74.8kJmol1$mol^{-1}$

Q-6: A reaction, X + Y à U + V + q is having a +ve entropy change. Then the reaction ____.

(a) will be possible at low temperature only

(b) will be possible at high temperature only

(c) will be possible at any temperature

(d) won’t be possible at any temperature

Ans:

(c) will be possible at any temperature

ΔG$\Delta G$ should be –ve, for spontaneous reaction to occur

ΔG$\Delta G$ = ΔH$\Delta H$ – TΔS$\Delta S$

As per given in question,

ΔH$\Delta H$ is –ve ( as heat is evolved)

ΔS$\Delta S$ is +ve

Therefore, ΔG$\Delta G$ is negative

So, the reaction will be possible at any temperature.

Q-7: In the process, system absorbs 801 J and work done by the system is 594 J. Find ΔU$\Delta U$ for the given process.

Ans:

As per Thermodynamics 1st law,

ΔU$\Delta U$ = q + W(i);

ΔU$\Delta U$ internal energy = heat

W = work done

W = -594 J (work done by system)

q = +801 J (+ve as heat is absorbed)

Now,

ΔU$\Delta U$ = 801 + (-594)

ΔU$\Delta U$ = 207 J

Q-8: The reaction given below was done in bomb calorimeter, and at 298K we get, ΔU$\Delta U$ = -753.7 kJ mol1$mol^{-1}$. Find ΔH$\Delta H$ at 298K.

NH2CN(g) +3/2O2(g) à N2(g) + CO2(g) + H2O(l)

Ans:

ΔH$\Delta H$ is given by,

ΔH=ΔU+ΔngRT$\Delta H = \Delta U + \Delta n_{g}RT$………………(1)

Δng$\Delta n_{g}$ = change in number of moles

ΔU$\Delta U$ = change in internal energy

Here,

Δng=ng(product)ng(reactant)$\Delta n_{g} = \sum n_{g}(product) – \sum n_{g}(reactant)$

= (2 – 2.5) moles

Δng$\Delta n_{g}$ = -0.5 moles

Here,

T =298K

ΔU$\Delta U$ = -753.7 kJmol1$kJmol^{-1}$

R  = 8.314×103kJmol1K1$8.314\times 10^{-3}kJmol^{-1}K^{-1}$

Now, from (1)

ΔH=(753.7kJmol1)+(0.5mol)(298K)(8.314×103kJmol1K1)$\Delta H = (-753.7 kJmol^{-1}) + (-0.5mol)(298K)( 8.314\times 10^{-3}kJmol^{-1}K^{-1})$

= -753.7 – 1.2

ΔH$\Delta H$ = -754.9 kJmol1$kJmol^{-1}$

Q-9: Calculate the heat (in kJ) required for 50.0 g aluminium to raise the temperature from 45Cto65C$45^{\circ}C\; to\; 65^{\circ}C$. For aluminium molar haet capacity is 24 Jmol1K1$J mol^{-1}K^{-1}$

Ans:

Expression of heat(q),

q=mCPΔT$q = mCP\Delta T$;………………….(a)

ΔT$\Delta T$ = Change in temperature

c = molar heat capacity

m = mass of substance

From (a)

q=(5027mol)(24mol1K1)(20K)$q = ( \frac{50}{27}mol )(24mol^{-1}K^{-1})(20K)$

q = 888.88 J q

Q-10: Calculate ΔH$\Delta H$ for transformation of 1 mole of water into ice from10Cto(10)C$10^{\circ}C\;to\;(-10)^{\circ}C$.. ΔfusH=6.03kJmol1at10C$\Delta _{fus}H = 6.03 kJmol^{-1} \;at\;10^{\circ}C$.

Cp[H2O(l)]=75.3Jmol1K1$C_{p}[H_{2}O_{(l)}] = 75.3 J\;mol^{-1}K^{-1}$

Cp[H2O(s)]=36.8Jmol1K1$C_{p}[H_{2}O_{(s)}] = 36.8 J\;mol^{-1}K^{-1}$

Ans:

ΔHtotal$\Delta H_{total}$ = sum of the changes given below:

(a) Energy change that occurs during transformation of 1 mole of water from 10Cto0C$10^{\circ}C\;to\;0^{\circ}C$.

(b) Energy change that occurs during transformation of 1 mole of water at 0C$0^{\circ}C$  to  1 mole of ice at 0C$0^{\circ}C$.

(c) Energy change that occurs during transformation of 1 mole of ice from 0Cto(10)C$0^{\circ}C\;to\;(-10)^{\circ}C$.

ΔHtotal=Cp[H2OCl]ΔT+ΔHfreezingCp[H2Os]ΔT$\Delta H_{total} = C_{p}[H_{2}OCl]\Delta T + \Delta H_{freezing} C_{p}[H_{2}O_{s}]\Delta T$

= (75.3 Jmol1K1$J mol^{-1}K^{-1}$)(0 – 10)K + (-6.03*1000 Jmol1$J mol^{-1}$(-10-0)K

= -753 Jmol1$J mol^{-1}$ – 6030Jmol1$J mol^{-1}$ – 368Jmol1$J mol^{-1}$

= -7151 Jmol1$J mol^{-1}$

= -7.151kJmol1$kJ mol^{-1}$

Thus, the required change in enthalpy for given transformation is -7.151kJmol1$kJ mol^{-1}$.

Q-11: Enthalpies of formation for CO2(g), CO(g), N2O4(g), N2O(g) are -393kJmol1$kJ mol^{-1}$,-110kJmol1$kJ mol^{-1}$, 9.7kJmol1$kJ mol^{-1}$ and 81kJmol1$kJ mol^{-1}$ respectively. Then, ΔrH$\Delta _{r}H$ = _____.

N2O4(g) + 3CO(g) à  N2O(g) + 3 CO2(g)

Ans:

ΔrH$\Delta _{r}H$ for any reaction is defined as the fifference between ΔfH$\Delta _{f}H$ value of products and ΔfH$\Delta _{f}H$ value of reactants.”

ΔrH=ΔfH(products)ΔfH(reactants)$\Delta _{r}H = \sum \Delta _{f}H (products) – \sum \Delta _{f}H (reactants)$

Now, for

N2O4(g) + 3CO(g) à  N2O(g) + 3 CO2(g)

ΔrH=[(ΔfH(N2O)+(3ΔfH(CO2))(ΔfH(N2O4)+3ΔfH(CO))]$\Delta _{r}H = [(\Delta _{f}H(N_{2}O) + (3\Delta _{f}H(CO_{2})) – (\Delta _{f}H(N_{2}O_{4}) + 3\Delta _{f}H(CO))]$

Now, substituting the given values in the above equation, we get:

ΔrH$\Delta _{r}H$ = [{81kJmol1$kJ mol^{-1}$ + 3(-393) kJmol1$kJ mol^{-1}$} – {9.7kJmol1$kJ mol^{-1}$ + 3(-110) kJmol1$kJ mol^{-1}$}]

ΔrH$\Delta _{r}H$ = -777.7 kJmol1$kJ mol^{-1}$

Q-12 Enthalpy of combustion of C to CO2 is -393.5 kJmol1$kJ mol^{-1}$. Determine the heat released on the formation of 37.2g of CO2 from dioxygen and carbon.

Ans:

Formation of carbon dioxide from di-oxygen and carbon gas is given as:

C(s) + O2(g) à CO2(g); ΔfH$\Delta _{f}H$ = -393.5 kJmol1$kJ mol^{-1}$

1 mole CO2 = 44g

Heat released during formation of 44g CO2 = -393.5 kJmol1$kJ mol^{-1}$

Therefore, heat released during formation of 37.2g of CO2  can be calculated as

= 393.5kJmol144g×37.2g$\frac{-393.5kJmol^{-1}}{44g}\times 37.2g$

= -332.69 kJmol1$kJ mol^{-1}$

Q-13: N2(g) + 3H2(g) à 2NH3(g) ; ΔrHΘ$\Delta _{r}H^{\Theta }$ = -92.4 kJmol1$kJ mol^{-1}$

Standard Enthalpy for formation of ammonia gas is _____.

Ans:

“Standard enthalpy of formation of a compound is the enthalpy that takes place during the formation of 1 mole of a substance in its standard form, from its constituent elements in their standard state.”

Dividing the chemical equation given in the question by 2, we get

(0.5)N2(g) + (1.5)H2(g) à 2NH3(g)

Therefore, Standard Enthalpy for formation of ammonia gas

= (0.5) ΔrHΘ$\Delta _{r}H^{\Theta }$

= (0.5)(-92.4kJmol1$kJ mol^{-1}$)

= -46.2kJmol1$kJ mol^{-1}$

Q-14: Determine Standard Enthalpy of formation for CH3OH(l) from the data given below:

CH3OH(l) + (3/2)O2(g) à CO2(g) + 2H2O(l); ΔrHΘ$\Delta _{r}H^{\Theta }$ = -726 kJmol1$kJ mol^{-1}$

C(g) + O2(g) à CO2(g); ΔcHΘ$\Delta _{c}H_{\Theta }$ = -393 kJmol1$kJ mol^{-1}$

H2(g) + (1/2)O2(g) à H2O(l); ΔfHΘ$\Delta _{f}H^{\Theta }$ = -286 kJmol1$kJ mol^{-1}$

Ans:

C(s) + 2H2O(g) + (1/2)O2(g) à CH3OH(l) …………………………(i)

CH3OH(l) can be obtained as follows,

ΔfHΘ$\Delta _{f}H_{\Theta }$ [CH3OH(l)] = ΔcHΘ$\Delta _{c}H_{\Theta }$

2ΔfHΘ$\Delta _{f}H_{\Theta }$ΔrHΘ$\Delta _{r}H_{\Theta }$

= (-393 kJmol1$kJ mol^{-1}$) +2(-286kJmol1$kJ mol^{-1}$) – (-726kJmol1$kJ mol^{-1}$)

= (-393 – 572 + 726) kJmol1$kJ mol^{-1}$

= -239kJmol1$kJ mol^{-1}$

Thus, ΔfHΘ$\Delta _{f}H_{\Theta }$ [CH3OH(l)] = -239kJmol1$kJ mol^{-1}$

Q-15: Calculate ΔH$\Delta H$ for the following process

CCl4(g) à C(g) + 4Cl(g) and determine the value of bond enthalpy for C-Cl in CCl4(g).

ΔvapHΘ$\Delta _{vap}H^{\Theta }$ (CCl4) = 30.5 kJmol1$kJ mol^{-1}$.

ΔfHΘ$\Delta _{f}H^{\Theta }$ (CCl4) = -135.5 kJmol1$kJ mol^{-1}$.

ΔaHΘ$\Delta _{a}H^{\Theta }$ (C) = 715 kJmol1$kJ mol^{-1}$,

ΔaHΘ$\Delta _{a}H^{\Theta }$  is a enthalpy of atomisation

ΔaHΘ$\Delta _{a}H^{\Theta }$ (Cl2) = 242 kJmol1$kJ mol^{-1}$.

Ans:

“ The chemical equations implying to the given values of enthalpies” are:

(1) CCl4(l) à CCl4(g) ; ΔvapHΘ$\Delta _{vap}H^{\Theta }$ = 30.5 kJmol1$kJ mol^{-1}$

(2) C(s) à C(g) ΔaHΘ$\Delta _{a}H^{\Theta }$ = 715 kJmol1$kJ mol^{-1}$

(3) Cl2(g) à 2Cl(g) ; ΔaHΘ$\Delta _{a}H^{\Theta }$ = 242 kJmol1$kJ mol^{-1}$

(4) C(g) + 4Cl(g) à CCl4(g); ΔfHΘ$\Delta _{f}H^{\Theta }$ = -135.5 kJmol1$kJ mol^{-1}$

ΔH$\Delta H$ for the process CCl4(g) à C(g) + 4Cl(g) can be measured as:

ΔH=ΔaHΘ(C)+2ΔaHΘ(Cl2)ΔvapHΘΔfH$\Delta H = \Delta _{a}H^{\Theta }(C) + 2\Delta _{a}H^{\Theta }(Cl_{2}) – \Delta _{vap}H^{\Theta } – \Delta _{f}H$

= (715kJmol1$kJ mol^{-1}$) + 2(kJmol1$kJ mol^{-1}$) – (30.5kJmol1$kJ mol^{-1}$) – (-135.5kJmol1$kJ mol^{-1}$)

Therefore, H=1304kJmol1$H = 1304kJmol^{-1}$

The value of bond enthalpy for C-Cl in CCl4(g)

= 13044kJmol1$\frac{1304}{4}kJmol^{-1}$

= 326 kJmol1$kJ mol^{-1}$

Q-16: ΔU$\Delta U$ = 0 for isolated system, then what will be ΔU$\Delta U$?

Ans:

ΔU$\Delta U$ is positive ; ΔU$\Delta U$ > 0.

As, ΔU$\Delta U$ = 0 thenΔS$\Delta S$ will be +ve, as a result reaction will be spontaneous.

Q-17:

Following reaction takes place at 298K,

2X + Y à Z

ΔH$\Delta H$ = 400 kJmol1$kJ mol^{-1}$

ΔH$\Delta H$ = 0.2kJmol1K1$kJ mol^{-1}K^{-1}$

Find the temperature at which the reaction become spontaneous considering ΔS$\Delta S$ and ΔH$\Delta H$ to be constant over the entire temperature range?

Ans:

Now,

ΔG=ΔHTΔS$\Delta G = \Delta H – T\Delta S$

Let, the given reaction is at equilibrium, then ΔT$\Delta T$ will be:

T = (ΔHΔG)1ΔS$(\Delta H – \Delta G)\frac{1}{\Delta S}$

ΔHΔS$\frac{\Delta H}{\Delta S}$; (ΔG$\Delta G$ = 0 at equilibrium)

= 400kJmol1$kJ mol^{-1}$/0.2kJmol1K1$kJ mol^{-1}K^{-1}$

Therefore, T = 2000K

Thus, for the spontaneous, ΔG$\Delta G$ must be –ve and T > 2000K.

Q-18: 2Cl(g) à Cl2(g)

In above reaction what can be the sign for ΔS$\Delta S$ and ΔH$\Delta H$?

Ans:

ΔS$\Delta S$ and ΔH$\Delta H$ are having negative sign.

The reaction given in the question represents the formation of Cl molecule from Cl atoms. As the formation of bond takes place in the given reaction. So, energy is released. So,  ΔH$\Delta H$ is negative.

Also, 2 moles of Chlorine atoms is having more randomness than 1 mole of chlorine molecule. So, the spontaneity is decreased. Thus, ΔS$\Delta S$ is negative.

Q-19: 2X(g) + Y(g) à 2D(g)

ΔUΘ$\Delta U^{\Theta }$ = -10.5 kJ and ΔSΘ$\Delta S^{\Theta }$ = -44.1JK1$JK^{-1}$

Determine ΔGΘ$\Delta G^{\Theta }$ for the given reaction, and predict that whether given reaction can occur spontaneously or not.

Ans:

2X(g) + Y(g) à 2D(g)

Δng$\Delta n_{g}$ = 2 – 3

= -1 mole

Putting value of ΔUΘ$\Delta U^{\Theta }$ in expression of ΔH$\Delta H$:

ΔHΘ=ΔUΘ+ΔngRT$\Delta H^{\Theta } = \Delta U^{\Theta } + \Delta n_{g}RT$

= (-10.5KJ) – (-1)( 8.314×103kJK1mol1$8.314\times 10^{-3}kJK^{-1}mol^{-1}$)(298K)

= -10.5kJ -2.48kJ

ΔHΘ$\Delta H^{\Theta }$ = -12.98kJ

Putting value of ΔSΘ$\Delta S^{\Theta }$ and ΔHΘ$\Delta H^{\Theta }$ in expression of ΔGΘ$\Delta G^{\Theta }$:

ΔGΘ=ΔHΘTΔSΘ$\Delta G^{\Theta } = \Delta H^{\Theta } – T\Delta S^{\Theta }$

= -12.98kJ –(298K)(-44.1JK1$JK^{-1}$)

= -12.98kJ +13.14kJ

ΔGΘ$\Delta G^{\Theta }$ = 0.16kJ

As, ΔGΘ$\Delta G^{\Theta }$ is positive, the reaction won’t occur spontaneously.

Q-20: Find the value of ΔGΘ$\Delta G^{\Theta }$ for the reaction, if equilibrium is given 10.given that T = 300K and R =  8.314×103kJK1mol1$8.314\times 10^{-3}kJK^{-1}mol^{-1}$.

Ans:

Now,

ΔGΘ$\Delta G^{\Theta }$ = 2.303RTlogeq$-2.303RT\log eq$

= (2.303)( 8.314×103kJK1mol1$8.314\times 10^{-3}kJK^{-1}mol^{-1}$)(300K) log10$\log 10$

= -5744.14Jmol1$Jmol^{-1}$

-5.744kJmol1$kJmol^{-1}$

Q-21: What can be said about the thermodynamic stability of NO(g), given

(1/2)N2(g) + (1/2)O2(g); ΔrHΘ=90kJmol1$\Delta _{r}H^{\Theta } = 90kJmol^{-1}$

NO(g) + (1/2)O2(g) à NO2(g) ; ΔrHΘ=74kJmol1$\Delta _{r}H^{\Theta } = -74kJmol^{-1}$

Ans:

The +ve value of ΔrH$\Delta _{r}H$ represents that during NO(g) formation from O2 and N2,  heat is absorbed. The obtained product, NO(g) is having more energy than reactants. Thus, NO(g) is unstable.

The -ve value of ΔrH$\Delta _{r}H$ represents that during NO2(g) formation from O2(g) and NO(g),  heat is evolved. The obtained product, NO2(g) gets stabilized with minimum energy.

Thus, unstable NO(g) converts into unstable  NO2(g).

Q-22: Determine ΔS$\Delta S$ in surrounding given that  mole of H2O(l) is formed I standard condition. ΔrHΘ=286kJmol1$\Delta _{r}H^{\Theta } = -286kJmol^{-1}$.

Ans:

ΔrHΘ=286kJmol1$\Delta _{r}H^{\Theta } = -286kJmol^{-1}$ is given so that amount of heat is evolved during the formation of 1 mole of H2O(l).

Thus, the same heat will be absorbed by surrounding. Qsurr = +286kJmol1$kJmol^{-1}$.

Now, ΔSsurr$\Delta S_{surr}$ = Qsurr/7

= 286kJmol1298K$\frac{286kJmol^{-1}}{298K}$

Therefore, ΔSsurr=959.73Jmol1K1$\Delta S_{surr} = 959.73Jmol^{-1}K^{-1}$

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