Thermodynamics : Notes and Study Materials -pdf
- Concepts of Thermodynamics
- Thermodynamics Master File
- Thermodynamics Revision Notes
- Thermodynamics MindMap
- NCERT Solution Thermodynamics
- NCERT Exemplar Solution Thermodynamics
- Thermodynamics: Solved Example 1
- Thermodynamics: Solved Example 2
- Thermodynamics : Practice Paper 1
- Thermodynamics : Practice Paper 2
- Thermodynamics : Practice Paper 3
Subtopics included in Class 11 Chemistry Chapter 6 Chemical Thermodynamics
- Thermodynamic Terms
- The System and the Surroundings
- Types of Thermodynamic Systems
- State of the System
- Internal Energy as a State Function
- Enthalpy, H
- Measurement Of ΔU And ΔH: Calorimetry
- Enthalpy Change and Reaction Enthalpy
- Enthalpies For Different Types Of Reactions
- Gibbs Energy Change And Equilibrium
Thermodynamics Chemistry Chapter 6
• Important Terms and Definitions
System: Refers to the portion of universe which is under observation.
Surroundings: Everything else in the universe except system is called surroundings. The Universe = The System + The Surroundings.
Open System: In a system, when there is exchange of energy and matter taking place with
the surroundings, then it is called an open system.
For Example: Presence of reactants in an open beaker is an example of an open system. Closed System: A system is said to be a closed system when there is no exchange of matter ‘ but exchange of energy is possible.
For example: The presence of reactants in a closed vessel made of conducting material.
Isolated System: In a system, when no exchange of energy or matter takes place with the surroundings, is called isolated system.
For example: The presence of reactants in a thermoflask, or substance in an insulated closed vessel is an example of isolated system.
Homogeneous System: A system is said to be homogeneous when all the constituents present is in the same phase and is uniform throughout the system.
For example: A- mixture of two miscible liquids.
Heterogeneous system: A mixture is said to be heterogeneous when it consists of two or more phases and the composition is not uniform.
For example: A mixture of insoluble solid in water. ’
The state of the system: The state of a thermodynamic system means its macroscopic or bulk properties which can be described by state variables:
Pressure (P), volume (V), temperature (T) and amount (n) etc.
They are also known as state functions.
Isothermal process: When the operation is carried out at constant temperature, the process is said to be isothermal. For isothermal process, dT = 0 Where dT is the change in temperature.
Adiabatic process: It is a process in which no transfer of heat between system and surroundings, takes place.
Isobaric process: When the process is carried out at constant pressure, it is said to be isobaric. i.e. dP = 0
Isochoric process: A process when carried out at constant volume, it is known as isochoric in nature.
Cyclic process: If a system undergoes a series of changes and finally returns to its initial state, it is said to be cyclic process.
Reversible Process: When in a process, a change is brought in such a way that the process could, at any moment, be reversed by an infinitesimal change. The change r is called reversible.
• Internal Energy
It is the sum of all the forms of energies that a system can possess.
In thermodynamics, it is denoted by AM which may change, when
— Heat passes into or out of the system
— Work is done on or by the system
— Matter enters or leaves the system.
Change in Internal Energy by Doing Work
Let us bring the change in the internal energy by doing work.
Let the initial state of the system is state A and Temp. TA Internal energy = uA
On doing’some mechanical work the new state is called state B and the temp. TB. It is found to be
TB > TA
uB is the internal energy after change.
∴ Δu = uB – uA
Change in Internal Energy by Transfer of Heat
Internal energy of a system can be changed by the transfer of heat from the surroundings to the system without doing work.
Δu = q
Where q is the heat absorbed by the system. It can be measured in terms of temperature difference.
q is +ve when heat is transferred from the surroundings to the system. q is -ve when heat is transferred from system to surroundings.
When change of state is done both by doing work and transfer of heat.
Δu = q + w
First law of thermodynamics (Law of Conservation of Energy). It states that, energy can neither be created nor be destroyed. The energy of an isolated system is constant.
Δu = q + w.
• Work (Pressure-volume Work)
Let us consider a cylinder which contains one mole of an ideal gas in which a frictionless piston is fitted.
• Work Done in Isothermal and Reversible Expansion of Ideal Gas
• Isothermal and Free Expansion of an Ideal Gas
For isothermal expansion of an ideal gas into vacuum W = 0
• Enthalpy (H)
It is defined as total heat content of the system. It is equal to the sum of internal energy and pressure-volume work.
Mathematically, H = U + PV
Change in enthalpy: Change in enthalpy is the heat absorbed or evolved by the system at constant pressure.
ΔH = qp
For exothermic reaction (System loses energy to Surroundings),
ΔH and qp both are -Ve.
For endothermic reaction (System absorbs energy from the Surroundings).
ΔH and qp both are +Ve.
Relation between ΔH and Δu.
• Extensive property
An extensive property is a property whose value depends on the quantity or size of matter present in the system.
For example: Mass, volume, enthalpy etc. are known as extensive property.
• Intensive property
Intensive properties do not depend upon the size of the matter or quantity of the matter present in the system.
For example: temperature, density, pressure etc. are called intensive properties.
• Heat capacity
The increase in temperature is proportional to the heat transferred.
q = coeff. x ΔT
q = CΔT
Where, coefficient C is called the heat capacity.
C is directly proportional to the amount of substance.
Cm = C/n
It is the heat capacity for 1 mole of the substance.
• Molar heat capacity
It is defined as the quantity of heat required to raise the temperature of a substance by 1° (kelvin or Celsius).
• Specific Heat Capacity
It is defined as the heat required to raise the temperature of one unit mass of a substance by 1° (kelvin or Celsius).
q = C x m x ΔT
where m = mass of the substance
ΔT = rise in temperature.
• Relation Between Cp and Cv for an Ideal Gas
At constant volume heat capacity = Cv
At constant pressure heat capacity = Cp
At constant volume qv= CvΔT = ΔU
At constant pressure qp = Cp ΔT = ΔH
For one mole of an ideal gas
ΔH = ΔU + Δ (PV) = ΔU + Δ (RT)
ΔH = ΔU + RΔT
On substituting the values of ΔH and Δu, the equation is modified as
Cp ΔT = CvΔT + RΔT
or Cp-Cv = R
• Measurement of ΔU and ΔH—Calorimetry
Determination of ΔU: ΔU is measured in a special type of calorimeter, called bomb calorimeter.
Working with calorimeter. The calorimeter consists of a strong vessel called (bomb) which can withstand very high pressure. It is surrounded by a water bath to ensure that no heat is lost to the surroundings.
Procedure: A known mass of the combustible substance is burnt in the pressure of pure dioxygen in the steel bomb. Heat evolved during the reaction is transferred to the water and its temperature is monitored.
• Enthalpy Changes During Phase Transformation
Enthalpy of fusion: Enthalpy of fusion is the heat energy or change in enthalpy when one mole of a solid at its melting point is converted into liquid state.
Enthalpy of vaporisation: It is defined as the heat energy or change in enthalpy when one mole of a liquid at its boiling point changes to gaseous state.
Enthalpy of Sublimation: Enthalpy of sublimation is defined as the change in heat energy or change in enthalpy when one mole of solid directly changes into gaseous state at a temperature below its melting point.
• Standard Enthalpy of Formation
Enthalpy of formation is defined as the change in enthalpy in the formation of 1 mole of a substance from its constituting elements under standard conditions of temperature at 298K and 1 atm pressure.
Enthalpy of Combustion: It is defined as the heat energy or change in enthalpy that accompanies the combustion of 1 mole of a substance in excess of air or oxygen.
• Thermochemical Equation
A balanced chemical equation together with the value of ΔrH and the physical state of reactants and products is known as thermochemical equation.
Conventions regarding thermochemical equations
1. The coefficients in a balanced thermochemical equation refer to the number of moles of reactants and products involved in the reaction.
• Hess’s Law of Constant Heat Summation
The total amount of heat evolved or absorbed in a reaction is same whether the reaction takes place in one step or in number of steps.
• Born-Haber Cycle
It is not possible to determine the Lattice enthalpy of ionic compound by direct experiment. Thus, it can be calculated by following steps. The diagrams which show these steps is known as Born-Haber Cycle.
Spontaneous Process: A process which can take place by itself or has a tendency to take place is called spontaneous process.
Spontaneous process need not be instantaneous. Its actual speed can vary from very slow to quite fast.
A few examples of spontaneous process are:
(i) Common salt dissolves in water of its own.
(ii) Carbon monoxide is oxidised to carbon dioxide of its own.
• Entropy (S)
The entropy is a measure of degree of randomness or disorder of a system. Entropy of a substance is minimum in solid state while it is maximum in gaseous state.
The change in entropy in a spontaneous process is expressed as ΔS
• Gibbs Energy and Spontaneity
A new thermodynamic function, the Gibbs energy or Gibbs function G, can be defined as G = H-TS
ΔG = ΔH – TΔS
Gibbs energy change = enthalpy change – temperature x entropy change ΔG gives a criteria for spontaneity at constant pressure and temperature, (i) If ΔG is negative (< 0) the process is spontaneous.
(ii) If ΔG is positive (> 0) the process is non-spontaneous.
• Free Energy Change in Reversible Reaction
CBSE Class 11 Chemistry Chapter-6 Important
1 Marks Questions
1.Define a system.
Ans. A system in thermodynamics refers to that part of the universe in which observations are made.
Ans.The rest of the universe which might be in a position to exchange energy and matter with the system is called its surroundings.
3.State the first law of thermodynamics.
Ans.The first law of thermodynamics stales that ‘the energy of an isolated system is constant’.
4.What kind of system is the coffee held in a cup?
Ans .Coffee held in a cup is an open system because it can exchange matter (water vapors) and energy (heat) with the surroundings.
5.Give an example of an isolated system.
Ans.Coffee held in a thermos flask is an isolated system because it can neither exchange energy nor matter with the surroundings.
6.Name the different types of the system.
Ans.There are three types of system –
7.What will happen to internal energy if work is done by the system?
Ans.The internal energy of the system will decrease if work is done by the system.
8.From thermodynamic point of view, to which system the animals and plants belong?
Ans. Open system.
9.How may the state of thermodynamic system be defined?
Ans.The state of thermodynamic system may be defined by specifying values of state variables like temperature, pressure, volume.
Ans. It is defined as total heat content of the system.
11.Give the mathematical expression of enthalpy.
H = U + pv where U is internal energy.
12.When is enthalpy change–
Ans. (i) is positive for endothermic reaction which absorbs heat from the surroundings.
(ii) is negative for exothermic reactions which evolve heat to the surroundings.
13.Give the expression for
(i)isothermal irreversible change, and
isothermal reversible change.
Ans.(i) For isothermal irreversible change Q = -w = pex (vf – vi)
(ii)For isothermal reversible changeq = -w = nRT ln
= 2.303 nRT log
14.Define Heat capacity
Ans.The heat capacity for one mole of the substance is the quantity of heat needed fo raise the temperature of one mole by one degree Celsius.
15.Define specific heat.
Ans. Specific heat /specific heat capacity is the quantity of heat required to raise the temperature of one unit mass of a substance by one degree Celsius (or one Kelvin).
16.Give the mathematical expression of heat capacity.
Ans.The mathematical expression of heat capacity
q = c x m x (c = heat capacity) when m = 1
where C = specific heat
m = mass
= temperature change.
17.Define reaction enthalpy.
Ans. The enthalpy change accompanying a reaction is called the reaction enthalpy
18.Define standard enthalpy.
Ans. The standard enthalpy of reaction is the enthalpy change for a reaction is the enthalpy change for a reaction when all the participating substances are in their standard states.
19.The standard heat of formation of Fe2O3 (s) is 824.2kJ mol-1 Calculate heat change for the reaction.
4Fe(s) + 302 (g) 2Fe2O3(s)
= [2 X Hfo Fe2O3(s) ] – [4Hf oFe (s) + 3Hf oO2(g)]
= 2(-824.2kJ) – [ 4 x o + 3 x o ]
20.Define spontaneous process.
Ans. A spontaneous process is an irreversible process and may only be reversed by some external agency.
21.Define non-spontaneres process.
Ans. A process is said to be non-spontaneous if it does not occur of its own under given condition and occur only when an external force is continuously applied.
22.What is the sign of enthalpy of formation of a highly stable compound?
23.Predict the sign of for the following reaction
Ans. is positive.
24.Two ideal gases under same pressure and temperature are allowed to mix in an isolated system – what will be sign of entropy change?
Ans. Entropy change is positive. It is because disorder or degree of freedom increase on mixing.
2 Marks Questions
1.Change in internal energy is a state function while work is not, why?
Ans. The change in internal energy during a process depends only upon the initial and final state of the system. Therefore it is a state function. But the work is related the path followed. Therefore, it is not a state function.
2.With the help of first law of thermodynamics and H = U + pv, prove = qp
Ans.The enthalpy is defined as
H = U + pv
For a change in the stales of system,
The first law of thermodynamics states that –
From (i) and (ii),
When the pressure is constant,
3.Why is the difference between and not significant for solids or liquids?
Ans. The difference between and is not usually significant for systems consisting of only solids and / or liquids because they do not suffer any significant volume changes upon heating.
4.What is an extensive and intensive property?
Ans. Extensive property is a property whose value depends on the quantity or size of matter present in the system.
Intensive property is a property which do not depend upon the quantity or size of matter present.
5.Show that for an ideal gas Cp- Cv = R
Ans.When a gas is heated under constant pressure, the heat is required for raising the temperature of the gas and also for doing mechanical work against the external pressure during expansion.
At constant volume, the heat capacity, C is written as Cv and at constant pressure this is denoted by Cp.
we write heat q
at constant volume as qv = Cv
at constant pressure as qp =
The difference between Cp and Cv can be derived for an ideal gas as :
For a mole of an ideal gas,
On putting the values of
Cp = Cv +R
6.Show that for an ideal gas, the molar heat capacity under constant volume conditions is equal to 3/2 R.
Ans.For an ideal gas, from kinetic theory of gases, the average kinetic energy per mole (Ek) of the gas at any temperature Tk is given by
At (T+1)k, the kinetic energy per mole (Ek1) is Ek1 =
Therefore increase in the average kinetic energy of the gas for 10C (or 1K) rise in temperature is
by definition is to the molar heat capacity of a gas at constant volume, Cv.
7.A 1.25g sample of octane (C18 H18) is burnt in excess of oxygen in a bomb calorimeter. The temperature of the calorimeter rises from 294.05 to 300.78K. If heat capacity of the calorimeter is 8.93 KJ/K. find the heat transferred to calorimeter.
Ans .Mass of octane,
M = 1.250g.
Heat capacity, c = 8.93 kJ/k
Rise in temp,
Heat transferred to calorimeter
= 0.00125 x 8.93 x 6.73
= 0.075 kJ
8.Calculate the heat of combustion of ethylene (gas) to from CO2 (gas) and H2O (gas) at 298k and 1 atmospheric pressure. The heats of formation of CO2, H2O and C2H4 are – 393.7, – 241.8, + 52.3 kJ per mole respectively.
Ans. C2H4 (g) + 302(g) 2CO2(g) + 2H2O (g)
= [2 x (CO2) + 2 x ] –
= 2 x[(-393.7)m+2x (-241.8)] – [(523.0) + 0)]
= [-787.4 – 483.6 ] -53.3
= – 1323.3 kJ.
9.Give two examples of reactions which are driven by enthalpy change.
Ans. Examples of reactions driven by enthalpy change:
The process which is highly exothermic, i.e. enthalpy change is negative and has large value but entropy change is negative is said to be driven by enthalpy change, eg.
10.Will the heat released in the following two reactions be equal? Give reasons in support of your answer.
(i)H2 (g) +
Ans. No, the heats released in the two reactions are not equal. The heat released in any reaction depends upon the reactants, products and their physical states. Here in reaction (i), the water produced is in the gaseous state whereas in reaction (ii) liquid is formed. As we know, that when water vapors condensed to from water, heat equal to the latent heat of vaporization is released. Thus, more heat is released in reaction (ii).
11.What is the relation between the enthalpy of reaction and bond enthalpy?
Ans .A chemical reaction involves the breaking of bonds in reactants and formation of new bonds in products. The heat of reaction (enthalpy change) depends on the values of the heat needed to break the bond formation .Thus
(Heat of reaction = (Heat needed to break the bonds in reactants – Heat liberated to from bonds in products).
= Bond energy in (to break the bonds) – Bond energy out (to form the bonds)
= Bond energy of reactants – Bond energy of products.
12.The reaction C (graphite) + O2 (g) CO2(g) + 393.5 kJ mol-1 represents the formation of CO2 and also combustion of carbon. Write the values of the two processes.
Ans.(i) The standard enthalpy of formation of CO2 is -393.5 kJ per mole of CO2.
(ii) The standard enthalpy of combustion of carbon is – 393.5 kJ per mole of carbon i.e.
13.Explain how is enthalpy related to spontaneity of a reaction?
Ans.Majority of the exothermic reactions are spontaneous because there is decrease in energy.
Burning of a substance is a spontaneous process.
Neutralisation of an acid with a base is a spontaneous reaction.
Many spontaneous reactions proceed with the absorption of heat. Conversion of water into water vapour is an endothermic spontaneous change. Therefore change in enthalpy is not the only criterion for deciding the spontaneity of a reaction.
14.The are given + 61.17kJ mol-1 and + 132 Jk-1mol-1 respectively. Above what temperature will the reaction be spontaneous?
Ans. The reaction
Will be spontaneous when is negative.
Shows that would be –ve when,
Or T >
The process will be spontaneous above a temperature of .
3 Marks Questions
1.Give the relationship between for gases.
Ans.For gases the volume change is appreciable.
let VA be the total volume of gaseous reactants, and
VB be the total volume of gaseous product.
nA be the number of moles of the reactant and
nB be the number of moles of the product,
Then at constant pressure and temperature,
p VA = nA RT
p VB = nB RT
or p VB – pVA = (nB – nA) RT
where and is equal to the difference between the number of moles of gaseous products and gaseous reactants.
Substituting the value of p awe get.
(heat change under constant pressure)
(heat change under constant volume)
for gaseous system.
2.It has been found that 221.4J is needed to heat 30g of ethanol from 150C to 180C. calculate (a) specific heat capacity, and (b) molar heat capacity of ethanol.
Ans.(a) Specific heat capacity
Since 10C is equal to 1k, the specific heat capacity of ethanol = 2.46Jg-1 0c-1.
(b) Molar heat capacity, Cm = specific heat x molar mass.
Therefore, Cm (ethanol) = 2.46 x 46
= 113.2 Jmol-1 0c-1
The molar heat capacity of ethanol is 113.2 J mol-1 0c-1.