Class 11 Chemistry Chapter 6 Thermodynamics
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NCERT Notes For Class 11 Chemistry Chapter 6 Thermodynamics
Class 11 Chemistry Chapter 6 Thermodynamics
- It is a branch of science that deals with the relationship between heat and work.
- Chemical thermodynamics is a branch of chemistry that deals with the heat changes associated with chemical reactions.
System and Surroundings
- System is the part of the universe which is under observation or investigation.
- The part of the universe except system is called surroundings.
- The system and surroundings are separated by a boundary which may be real or imaginary.
System + surroundings → Universe
Types of systems
Depending on the ability to exchange energy and matter with the surroundings, systems are classified into three types:
- Open system: It is a system that can exchange both energy and matter with the surroundings. E.g. Hot water taken in an open vessel.
- Closed system: It is a system that can exchange only energy and not matter with the surroundings. E.g. Hot water taken in a closed vessel.
- Isolated system: It is a system that cannot exchange both energy and matter with the surroundings. E.g. Hot water taken in a thermo flask.
Depending on the number of particles, systems are of two types:
- Microscopic system: It is a system containing a few number of particles.
- Macroscopic system: it is a system contains a large number of particles.
- The properties of a macroscopic system are called macroscopic properties.
- The important macroscopic properties are:
- Temperature (T), Pressure (P), Volume (V), length (l), breadth (b), height (h), internal energy (U), enthalpy (H), entropy (S), Gibb’s energy (G) etc.
Macroscopic properties can be divided into two:
- Extensive properties: These are properties which depend on the amount of matter present in the system.
- Or, these are the properties which change when a system is divided.
- E.g.: Volume (v), length (l), breadth (b), height (h), internal energy (U), enthalpy (H), entropy (S), Gibb’s energy (G), heat capacity etc.
- Intensive properties: These are properties which are independent of the amount of matter present in the system.
- Or, these are the properties which do not change when a system is divided.
- E.g. : Temperature (T), Pressure (P), Volume (V), density, refractive index, molar heat capacity, viscosity, surface tension etc.
State and Path functions
- A function or a property that depends only on the initial and final state of a system and not on the path followed is called a state function.
- E.g. for state functions: T, P, V, U, H, S, G etc.
- Path functions: These are properties which depend on the path followed also.
- E.g. heat (q) and work (w)
A process is the method (path) by which a state change occurs in a system. The different types of thermodynamic process are:
- Isothermal process: It is a process that occurs at constant temperature. For this process ΔT = 0 but Δq≠ 0
- Isobaric process: It is a process that occurs at constant Pressure. For such a process ΔP = 0
- Isochoric process: It is a process that occurs at constant volume. For such a process ΔV = 0
- Adiabatic process: It is a process that occurs at constant heat energy. Here no heat enters into or leaves from the system. For such a process Δq= 0 but ΔT ≠ 0
- Cyclic process: It is a process that takes place in a cyclic manner.
- Here the system undergoes a series of changes and finally returns to its initial state.
- For such a process ΔU=0 and ΔH = 0
- Reversible process: Every process is associated with two types of forces – driving force and opposing force.
- Driving force favours the process while opposing force opposes it.
- If the driving and opposing forces are differed by an infinitesimally small quantity the process takes place in both directions.
- Such a process is called reversible process.
- Irreversible process: If the driving and opposing forces are differed by a large quantity, then the process takes place in only one direction.
- Such a process is called irreversible process.
Heat and Work
- Heat: It is a form of energy.
- Heat flows from a hot body to a cold body, when there is a thermal contact between the two.
- When a body absorbs heat, its energy increases and when it evolves heat, its energy decreases.
- Thus by international convention, when heat is absorbed by a system, q becomes +ve and when heat is evolved by a system q becomes –ve.
Work: In thermodynamics, there are two types of work – expansion work and non-expansion work.
- Expansion work (Wexp): It is related to gaseous systems. It is the product of pressure and change in volume. i.e., expansion work (Wexp) = -PΔV
- For irreversible process, (Wexp) = -PΔV and
- For reversible process, (Wexp) = -2.303nRT log(V2/V1)
- Non-expansion work (Wnon-exp): It is related to electrochemical cells. It is the product of potential difference (E) and charge (Q). i.e., Wnon-exp = E x Q
By international convention, w becomes +ve, when work is done on the system and w becomes –ve, when work is done by the system.
Free expansion: Expansion of a gas into vacuum (external pressure = 0) is called free expansion. No work is done during free expansion of an ideal gas whether the process is reversible or irreversible.
Internal Energy (U)
- Every body is associated with a definite amount of energy.
- This energy is called internal or intrinsic energy.
- It is the energy possessed by a body.
- It is the sum of different types of molecular energies like translational kinetic energy, rotational kinetic energy, vibrational kinetic energy, electronic energy, nuclear energy etc.
- Internal energy of a body is an extensive property and state function.
- We cannot calculate the exact value of internal energy, but we can calculate the change in internal energy (ΔU) during a process by using an apparatus called Bomb Calorimeter.
- The change in internal energy (ΔU) = U2-U1, where U1 and U2 are the initial and final internal energies respectively. The unit of internal energy is kJ/mol.
The internal energy of a system can change in mainly by two ways:
- By the transfer of heat
- By doing work
First law of thermodynamics: It is same as law of conservation of energy.
- It states that energy can neither be created nor be destroyed.
- Or, the total energy in the universe is always a constant.
- Or, the total energy of an isolated system is always a constant.
Mathematically ΔU = q + w ………………. (1)
Where q is the amount of heat absorbed by the system and w is the amount of work done on the system.
If there is only expansion work, the above equation becomes ΔU = q – PΔV ………………(2)
(since, Wexp = -PΔV)
For a system containing only solids or liquids, ΔV = 0. So ΔU = q
For an isothermal reversible process, ΔU = 0. So, q = -w = 2.303nRT log V2/V1 For an adiabatic process, q =0, so ΔU = w.
Significance of ΔU
We know that ΔU = q – PΔV
For a process taking place at constant volume, ΔV = 0. So ΔU = qv
i.e., ΔU gives the amount of heat absorbed or evolved by a system at constant volume.
- It is the total heat content of a system. It is the sum of internal energy and pressure-volume energy of a system.
i.e. H = U + PV
- It is a state function and an extensive property. i.e, whenever there is a change in enthalpy, it depends only on the initial and final values and not on the path followed.
- If H1 is the enthalpy of a system in the initial state and H2 is that in the final state, then the change in enthalpy, ΔH = H2 – H1.
- The unit of enthalpy is kJ/mol. ΔH is determined by using an apparatus called calorimeter.
- For a chemical reaction, if HP is the enthalpy of products and HR is that of reactants, then
ΔH = HP – HR.
- If HP > HR, ΔH is positive and heat is absorbed during the process. Such a process is called endothermic process.
- If HP < HR, ΔH is negative and heat is evolved during the process. Such a process is called exothermic process.
Relation between ΔH and ΔU
Consider a gaseous reaction taking place at constant pressure (P) and temperature (T). Let H1, U1 and V1 be the initial enthalpy, internal energy and volume respectively. Let these values changes to H2, U2 and V2 respectively. Then,
H1 = U1 + PV1 ……………… (1)
And H2 = U2 + PV2 ……………… (2)
(2) – (1) gives: H2 –H1 = (U2 + PV2) – (U1 + PV1)
Or, ΔH = (U2 – U1) + (PV2 – PV1) Or, ΔH = ΔU + P (V2 – V1)
Or, ΔH = ΔU + PΔV …… (3)
From ideal gas equation, PV = nRT
If n1 and n2 are the total no. of moles of reactants and products respectively, then
PV1 = n1RT and PV2 = n2 RT
PV2 – PV1 = n2 RT – n1RT
Or, P(V2 –V1) = (n2 – n1)RT
Or, PΔV = ΔnRT
So equation (3) becomes, ΔH = ΔU + ΔnRT
Where Δn = np(g) – nR(g)
i.e., Δn = no. of moles of gaseous products – no. of moles of gaseous reactants If Δn =0, then ΔH = ΔU If Δn > 0, then ΔH > ΔU And if Δn < 0, then ΔH < ΔU.
Significance of ΔH
We know that ΔH = ΔU + PΔV (at constant pressure)
From first law of thermodynamics, ΔU = q – PΔV
So, q = ΔU + PΔV
From the above equations, we can write, ΔH = qp
i.e. enthalpy change in a chemical reaction is equal to the amount of heat evolved or absorbed at constant pressure. This is the significance of ΔH.
- It is the amount of heat required to raise the temperature of a body through 10C or 1K. if q amount of heat is required to raise the temperature of a body through ΔT, then heat capacity (C) = q/ΔT.
- Specific heat capacity is the amount of heat required to raise the temperature of unit mass of a body through 10C or 1K.
- Molar heat capacity is the amount of heat required to raise the temperature of 1 mole of a substance through 10C or 1K.
Relation between Cp and Cv
At constant volume, the heat capacity, C is denoted by Cv and at constant pressure, it is denoted by Cp .
We know that heat, q = C ΔT
At constant volume, the heat qV = ΔU = Cv.ΔT
At constant pressure, qp = ΔH = Cp.Δ T
We know that H = U +PV
For 1 mole of an ideal gas, ΔH = ΔU + Δ(pV )
= ΔU + Δ(RT) [Since PV = RT for one mole of ideal gas] = ΔU + RΔT ∴ ΔH = ΔU+R ΔT
On putting the values of ΔH and ΔU, we get
CpΔT = CvΔT + RΔT Or, Cp = Cv +R
Or, Cp– Cv = R
ENTHALPY CHANGE OF A REACTION – REACTION ENTHALPY (ΔrH)
In a chemical reaction, reactants are converted into products and is represented by, Reactants → Products
The enthalpy change during a chemical reaction is called the reaction enthalpy. It is given by the symbol ΔrH.
ΔrH = (sum of enthalpies of products) – (sum of enthalpies of reactants) i.e. ΔrH = ∑ HP– ∑ HR
Standard enthalpy of reactions (ΔrH0)
- The standard enthalpy of reaction is the enthalpy change for a reaction when all the substances are in their standard states.
- In thermodynamics, the standard state refers to the existence of a substance in its pure form at 1 bar pressure and at 298 K temperature.
Standard enthalpy of formation (ΔfH0)
- It is the standard enthalpy change for the formation of one mole of a compound from its elements in their most stable state of aggregation (reference state). It is denoted by the symbol ΔfH0. For e.g. the std. enthalpy of formation of CO2 is the enthalpy change when 1 mole of CO2 is formed from C and O2 at 298K temperature, 1 bar pressure and all the substances are in their stable state.
- The reference state of an element is its most stable state at 25°C and 1 bar pressure. The reference state of dihydrogen is H2 gas and those of dioxygen, carbon and sulphur are O2 gas, Cgraphite and SRhombic respectively. The enthalpy of formation may be positive or negative.
- By convention, standard enthalpy of formation (ΔfH0) of an element in reference state ( i.e., its most stable state of aggregation) is taken as zero.
- By knowing the std. enthalpies of formation of reactants and products, we can determine the std. enthalpy of reaction by the equation:
ΔrH0 = ∑ ΔfH0(P) – ∑ ΔfH0(R)
i.e. Std. enthalpy of reaction = Sum of the standard enthalpies of formation of products – Sum of the std. enthalpies of formation of reactants.
Enthalpies of phase transition
It is the enthalpy change when one mole of a substance changes from one phase to another phase at a particular temperature. The important types of Enthalpies of phase transition are:
- Enthalpy of fusion (ΔfusH0): It is the enthalpy change when one mole of a solid substance changes to liquid
state at its melting point. Melting of a solid is endothermic, so all enthalpies of fusion are positive.
- Enthalpy of vaporization (ΔvapH0):It is the enthalpy change when one mole of a liquid substance changes to
its vapour state at its boiling point.
- Enthalpy of sublimation (ΔsubH0): It is the enthalpy change when one mole of a solid substance is directly converted to gaseous state at a particular temperature below its m.p.
Hess’s Law of Constant Heat Summation
- The law states that the total enthalpy change for a physical or chemical process is the same whether the reaction taking place in a single step or in several steps. Or, the total enthalpy change for a process is independent of the path followed.
- Thus according to Hess’s law, if a reaction takes place in several steps then its standard reaction enthalpy is the sum of the standard enthalpies of the intermediate reactions.
Consider a process in which the reactant A is converted to product B in a single step by involving heat change ΔH. Let the same reactant A is first converted to C, then to D and finally to B involving heat changes ΔH1, ΔH2 and ΔH3 respectively.
Then according to Hess’s law: ΔH = ΔH1 + ΔH2 + ΔH3
A balanced chemical equation together with the value of enthalpy of reaction is called a thermochemical equation. Here the physical state of the each of the reactants and products is also shown in bracket.
e.g. H2(g) + ½ O2(g) → H2O(g), ΔrH0 = -285.83 kJ/mol
or, H2(g) + ½ O2(g) → H2O(g) + 285.83 kJ/mol
Standard enthalpy of combustion (ΔcH0)
It is defined as the enthalpy change when 1mole of a substance is completely burnt in presence of excess of air or oxygen and all the reactants and products being in their standard states. It is always negative since heat is always evolved during combustion.
e.g. CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g)
Enthalpy of atomization (ΔaH0)
- It is the enthalpy change on breaking one mole of bonds completely to obtain atoms in the gas phase.
- In the case of diatomic molecules the enthalpy of atomization is same as the bond dissociation enthalpy.
e.g. H2(g) → 2H(g); ΔaH0 = 435.0 kJ mol–1
Bond Enthalpy (ΔbondH0)
- Chemical reactions involve the breaking and making of chemical bonds. Energy is required to break a bond and energy is released when a bond is formed. So enthalpy of bond dissociation is positive and enthalpy of bond formation is negative.
- The amount of heat required to break one mole of covalent bonds to form products in the gas phase is called the bond dissociation enthalpy. For diatomic molecules, bond dissociation enthalpy and bond enthalpy are same.
e.g. H2(g) → 2H(g); ΔbondH0 = 435.0 kJ mol–1
- In the case of polyatomic molecules, bond dissociation enthalpy is different for different bonds within the same molecule. So here bond enthalpy is the average of bond dissociation enthalpies of various similar bonds.
For e.g. for water, the bond enthalpy is calculated as follows:
H2O(g) → H(g) + OH(g); ΔaH0 = 502 kJ/mol
OH(g) → O(g) + H(g); ΔaH0 = 427 kJ/mol
Here the bond enthalpy of O-H bond is taken as the average of the two bond dissociation enthalpies. i.e. bond enthalpy of O-H bond = 502 + 427 / 2 = 464.5kJ/mol
The standard enthalpy of reaction, ΔrH0 is related to bond enthalpies of the reactants and products in gas phase reactions as:
ΔrH0 = ∑ bond enthalpies of reactants – ∑ bond enthalpies of products
This equation is valid when all the reactants and products in the reaction are in gaseous state.
Enthalpy of Solution (ΔsolH0)
Enthalpy of solution is the enthalpy change when one mole of a substance is dissolved in a specified amount of solvent.
The enthalpy of solution in water is determined by the values of the lattice enthalpy, ΔlatticeH0 and enthalpy of hydration of ions, ΔhydH0.
i.e. ΔsolH0 = ΔlatticeH0 + ΔhydH0
For most of the ionic compounds, ΔsolH0 is positive and the dissociation process is endothermic. Therefore the solubility of most salts in water increases with rise of temperature.
Lattice Enthalpy (ΔlatticeH0)
The lattice enthalpy of an ionic compound is the enthalpy change when one mole of an ionic compound dissociates
into gaseous ions.
We cannot determine lattice enthalpies directly by experiment. So an indirect method called a Born-Haber Cycle is
used to calculate it.
This can be explained by consider the formation of NaCl. Na(s) + ½ Cl2(g) → NaCl(s)
This involves the following steps:
- Conversion of solid sodium atom to gaseous sodium atom. The energy change involved in this process is called sublimation energy.
Na(s) → Na(g); ΔsubH0
- Conversion gaseous sodium atom to gaseous sodium ion. The energy change in this process is called ionisation enthalpy.
Na(g)→ Na+(g); ΔiH0
- Conversion of gaseous chlorine molecule to gaseous chlorine atom. The energy change during this process is called bond dissociation enthalpy.
Cl2(g)→ 2Cl(g); ΔbondH0
½ Cl2(g) → Cl(g); ½ ΔbondH0
- Conversion of chlorine atom to chloride ion. The energy change involved in this process is called electron gain enthalpy (ΔegH0).
Cl(g) → Cl-(g); ΔegH0
- Packing of Na+(g)and Cl–(g) to form NaCl(s). the energy change in this process is called lattice enthalpy (ΔlatticeH0)
Na+(g) + Cl–(g) → NaCl(s); ΔlatticeH0
The different steps can be represented in a cyclic for as follows:
By applying Hess’s law we can write: ΔfH0 = ΔsubH0 + ΔiH0 + ½ ΔbondH0 + ΔegH0 + ΔlatticeH0 Form this lattice enthalpy can be determined as:
ΔlatticeH0 = ΔfH0 – [ΔsubH0 + ΔiH0 + ½ ΔbondH0 + ΔegH0]
Enthalpy of Dilution (ΔdilH0):
It is the enthalpy change when a solution is diluted (i.e. more solvent is added to the solution. It depends on the original concentration of the solution and the amount of solvent added.
- It is a process that takes place without the help of any external agency.
- All natural processes are spontaneous. E.g. flow of water from high level to low level, flow of heat from hot body to cold body, inter mixing of gases, burning of fuels, melting of ice, evaporation of water etc.
- A spontaneous process cannot reverse its direction by its own.
- Spontaneous chemical reactions are also called feasible or probable or irreversible reactions.
A process that takes place with the help of an external agency is called non-spontaneous process. E.g. flow of water from low level to high level.
Criteria for spontaneity
- During spontaneous processes like burning of fuels, flow of heat from hot body to cold body, flow of water from high level to low level etc, energy of the system decreases. So one of the criteria for spontaneity is decrease in energy.
- But for some spontaneous processes like melting of ice, evaporation of water etc, heat is absorbed. i.e. the energy of the system increases during the process.
- The above processes are accompanied with increase in disorder (entropy) of the system. Thus another criterion for spontaneity is increase in disorderness or randomness of the system.
- It is a measure of degree of disorderness or randomness of a system. As the disorderness increases, entropy also increases. It is an extensive property and state function.
- If S1 is the initial entropy of a system and S2 is its final value, then the change in entropy ∆S = S2 –S1.
For a given substance, the solid state has the least entropy and the gaseous state has the most. When temperature increases entropy also increases.The unit of entropy and entropy change is J/K/mol.
If a system absorbs ‘q’ amount of heat reversibly at a temperature T, then the change in entropy,
∆S = qrev./T
Where qrev. is the amount of heat absorbed reversibly.
Entropy and spontaneity
- During a spontaneous process, disorderness of the system increases. Thus entropy increases and hence ∆S becomes positive.
- The total entropy change for the system and surroundings is given by ∆STotal = ∆Ssyst. + ∆Ssurr. For a spontaneous process, ∆STotal > 0
- When a system attains equilibrium, the entropy becomes maximum and there is no further change in entropy. So ∆S Total = 0
If ∆S Total < 0, the process is non-spontaneous.
Second Law of Thermodynamics:
It can be stated as the entropy of the universe always increases during every spontaneous process.
Third law of Thermodynamics:
It states that the entropy of any perfectly crystalline substance becomes zero at absolute zero of temperature. The importance of this law is that by using this, we can calculate the absolute value of entropy of pure substances from thermodynamical data alone.
Gibb’s energy (G)
- It is defined as the maximum amount of available energy that can be converted to useful work. It is given by the equation, G = H – TS
- It is an extensive property and a state function. If G1 is the initial Gibb’s energy and G2 is its final value, then the change in Gibb’s energy (∆G) = G2 – G1
- The unit of Gibb’s energy is kJ/mol.
Relation between ∆G, ∆H and ∆S (Gibb’s Equation)
We know that G = H- TS
So ∆G = ∆H – ∆(TS)
= ∆H – (T∆S + S∆T)
At constant temperature, ∆T = 0. So, ∆G = ∆H – T∆S
The change in Gibb’s energy for a system can be written as
∆Gsyst. = ∆Hsyst. – T∆Ssyst. This equation is known as Gibb’s equation.
Gibb’s energy and Spontaneity
We know that ∆STotal = ∆Ssyst. + ∆Ssurr. ……….. (1)
Consider a system which absorbs ‘q’ amount of heat reversibly at a constant temperature T and constant pressure P.
On multiplying by –T, we get
-T∆STotal = -T∆Ssyst. + ∆Hsyst.
Or, -T∆STotal = ∆Gsyst. (Since, ∆H syst. – T∆Ssyst. = ∆Gsyst.)
For a spontaneous process, ∆STotal is positive. So ∆Gsyst. is negative. (or, ∆Gsyst. < 0)
For a non-spontaneous process, ∆STotal is negative, so ∆Gsyst. is positive. (or, ∆Gsyst. > 0)
For a process at equilibrium, ∆STotal = 0. So ∆Gsyst. = 0
Conditions for ∆G to be negative
We know that for a spontaneous process, ∆Gsyst. is negative. Also ∆G = ∆H – T∆S.
- If ∆H is negative and ∆S is +ve, ∆G is always –ve and the process is always spontaneous.
- If both ∆H and ∆S are positive, ∆G will be –ve when T∆S > ∆H. This is possible at high temperature.
- If both ∆H and ∆S are negative, ∆G will be –ve when T∆S < ∆H. This is possible at low temperature. Note: If ∆H is +ve and ∆S -ve, then ∆G will be always +ve and the process will be always non-spontaneous.
Gibb’s energy change (∆G) and Equilibrium constant (K)
The standard Gibb’s energy change of a reaction (∆rG0) is related to equilibrium constant (K) by the equation:
∆rG0 = -RT lnK
Or, ∆rG0 = -2.303RT logK
Where R is the universal gas constant and T is the absolute temperature.
So for a spontaneous process, the value of K should be +ve.