Tamilnadu Samacheer Kalvi 11th Chemistry Notes Chapter 7 Thermodynamics Notes
Thermodynamics – It means flow of heat. It is a branch of science which deals the relation between energy, heat, work and accompanying changes in the nature and behaviour of various substances around us.
Macroscopic properties – Heat and work.
First law of Thermodynamics – It states that “The total energy of an isolated system remains constant though it may change from one form to another”.
ΔU = q + w
ΔU – Change in the internal energy
q – Amount of heat supplied to the system
w – Amount of work done on the system
System – It is defined as any portion of matter under thermodynamic consideration, which is separated from the rest of the universe by real or imaginary boundaries.
Homogeneous system – If the physical states of all its matter are uniform, it is called homogeneous system.
Heterogeneous system – If the physical states of all its matter are not uniform, it is called heterogeneous system.
Surroundings – Everything in the universe that is not the part of system is called surroundings.
Boundary – Anything which separates the system from its surrounding is called boundary.
Isolated system – A system which can exchange neither matter nor energy with its surroundings is called an isolated system, e.g., Hot water contained in a thermos flask.
Closed system – A system which can exchange only energy but not matter with its surroundings is called a closed system, e.g., Hot water contained in a closed beaker.
Open system – A system which can exchange both matter and energy with its surroundings is called an open system, e.g., Hot water contained in an open beaker.
Intensive properties – The properties that are independent of the mass or size of the system is called as intensive property, e.g., Refractive index, surface tension, density, temperature, boiling point and freezing point etc.
Extensive properties – The properties that are depended on the mass or size Of the system is called as extensive property, e.g., Volume, number of moles, mass and internal energy etc.
Thermodynamic process – The method of operation which can bring about the change in the system is called thermodynamic process, e.g., Heating, cooling and fusion.
Reversible process – The process in which the system and surroundings can be restored to the initial state from the final state without producing any changes in the thermodynamic properties of the universe is called a reversible process.
Irreversible process – The process in which the system and surroundings cannot be restored to the initial state from the final state is called an irreversible process.
Adiabatic process – It is defined as one in which there is no exchange of heat (q) between the system and surrounding during process, [q = 0]
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Isothermal process – It is defined as one in which the temperature of the system remains constant, during the change from its initial to final state. [dT = 0]
Isobaric process – It is defined as one in which the pressure of the system remains constant during its change from the initial to final state. [dP = 0]
Isochoric process – It is defined as one in which the volume of system remains constant during its change from initial to final state. [dV= 0]
Cyclic process – When a system returns to its original state after completing a series of changes the process is known as a cyclic process. For a cyclic process dU = 0, dH = 0, dP = 0, dV= 0 and dT= 0.
State functions – The properties of P, V, T and ‘n’ that are used to describe the state of a system are called as state functions.
Path functions – A path function is a thermodynamic property of the system whose value depends on the path by which the system goes from its initial to final states, e.g., Work and heat.
Internal energy – Internal energy of a system is equal to the energy possessed by all its constituents namely atoms, ions and molecules.
U = Internal energy = Ut + Uv + Ur + Ub + Ue + Ui.
ΔU – Change in internal energy = Uf– Ui.
Heat (q) – It is regarded as an energy in transit across the boundary separating a system from its surrounding. It is a path function.
Units of heat – SI unit of heat is joule (J).The other unit is calories (cal).
Calories (cal) – It is defined as the quantity of heat required to raise the temperature of 1 gram of water by 1 °C in the vicinity of 15°C.
Work (w) – It is defined as the force (F) multiplied by the displacement (x). [- w = F. x ]
Unit of work – SI unit of work is joule (J) or kilojoule (kJ).
Joule – It is defined as the work done by a force of one Newton through a displacement of one meter (J = Nm).
Gravitational work – When an object is raised to a certain height against the gravitational field, gravitational work is done on the object, [w = m.g.h]
Electrical work – When a charged body moves from one potential region to another, electrical work is done, [w = QV]
Mechanical work – It is defined as force multiplied by the displacement through which the force acts, [w = F. x]
Pressure – Volume work – The work done in expansion or compression of a gas. PV work (or) expansion work.
Work involved in expansion and compression process – [w = -PΔV]
Expansion – During expansion, work is done by the system. Since Vf > Vi., the sign obtained for work will be negative.
Compression – During compression, work is done on the system. Since Vf < Vi the sign obtained for work will be positive.
For Reveresible process wrev = 2.303 nRT log \(\left(\frac{\mathrm{V}_{f}}{\mathrm{~V}_{i}}\right)\)
Zeroth law of thermodynamics – This law states that ‘If two systems are separately in thermal equilibrium with a third one, then they tend to be in thermal equilibrium with themselves’.
First law of thermodynamics – “Energy can neither be created nor destroyed, but may be converted from one form to another”.
(or)
The total energy of a system and surroundings remains constant (or conserved).
ΔU = q + w
Mathematical statement of first law –
| ΔU = q + w | ||
| Cyclic process | ΔU = 0 | ∴ q = -w |
| Isochoric process | ΔV = 0 w = 0 |
∴ ΔU = qv |
| Adiabatic process | Q = 0 | ∴ ΔU = w |
| Isobaric process | ΔU = q + w | ΔU = q – p ΔV |
Enthalpy (H) – It is defined as the sum of the internal energy (U) of a system and the product of pressure and volume of the system.
H = U + PV
ΔH = ΔU + PΔV
Relation between enthalpy ‘H’ and internal energy ‘U’ –
ΔH — ΔU + Δn(g)RT
Standard heat of formation – It is defined as “the change in enthalpy that takes place when one mole of a compound is formed from its elements, present in their standard states (298K and 1 bar pressure).
Thermochemical equations – It is a balanced stoichiometric chemical equation that includes the enthalpy change (ΔH).
Standard enthalpy of reaction (ΔH0r) – It is the enthalpy change for a reaction when all the reactants and products are present in their standard states.
\(\Delta \mathrm{H}_{r}^{0}=\Sigma \Delta \mathrm{H}_{r}^{0} \text { (products) }-\Sigma \Delta \mathrm{H}_{r}^{\theta} \text { (reactants) }\)
Heat of combustion (ΔHc) – It is defined as the change in enthalpy of a system when one mole of the substance is completely burnt in excess of air or oxygen.
Specific heat capacity (C) – It is defined as the heat absorbed by one kilogram of a substance in raising its temperature by one Kelvin at a specified temperature.
\(\mathrm{C}=\left[\frac{q}{m\left(\mathrm{~T}_{2}-\mathrm{T}_{1}\right)}\right]\)
Molar heat capacity – It is defined as the amount of heat absorbed by one mole of the substance to raise its temperature by 1 Kelvin.
cm = q/ΔT
Unit of heat capacity – JK-1 mol-1
Cv : (Molar heat capacity at constant volume) – It is defined as the rate of change of internal energy with respect to temperature at constant volume.
\(C_{V}=\left(\frac{d U}{d T}\right)_{V}\)
Cp: (Molar heat capacity at constant pressure) – It is defined as the rate of change of enthalpy with respect to temperature at constant pressure.
\(C_{P}=\left(\frac{d H}{d T}\right)_{P}\)
Relation between Cp and Cv for an ideal gas –
Cp – Cv = nR
Calorimetry – It is the science associated with determining the changes in energy of a system by measuring the heat exchanged with the surroundings. ,
Exothermic reaction – The reaction in which heat is evolved and system loses heat to the surroundings. [ΔHr = -ve]
Endothermic reaction – The reaction in which heat is absorbed and system gains heat from the surroundings. [ΔHr= +ve] .
Calorific value – It is defined as the amount of heaf produced in calories (or joules) when one gram of a substance is completely burnt. The SI unit is J kg-1.
Heat of solution – It is defined as the change in enthalpy of the system when one mole of a substance is dissolved in a specified quantity of a solvent at a given temperature.
Heat of neutralisation – It is defined as the change in enthalpy of the system when one gram equivalent of an acid is completely neutralised by one gram equivalent of a base in dilute solution.
Molar heat of fusion – It is defined as the change in enthalpy when one mole of a solid substance is converted into the liquid state at its melting point.
Molar heat of vapourisation – It is defined as the change in enthalpy when one mole of a liquid is converted into vapour or gaseous state at its boiling point.
Molar heat of sublimation – It is defined as the change in enthalpy when one mole of a solid is directly converted into the gaseous state at its sublimation temperature.
Heat of transition – It is defined as the change in enthalpy when one mole of an element changes from one allotropic form to another.
Hess’s law of constant heat summation – It states that “the enthalpy change of a reaction either at constant volume or constant pressure is the same whether it takes place in a single step or multiple steps provided the initial and final states are same”.
Lattice energy – It is defined as the amount of energy required to completely remove the constituent ions from its crystal lattice to an infinite distance.
Born Haber’s cycle – It applies Hess’s law to calculate lattice enthalpy.
Second law of thermodynamics –
(i) Entropy statement: “Whenever a spontaneous process takes place, it is accompanied by an increase in the total entropy of the Universe”.
ΔSUniverse > ΔSsystem + ΔSsurroundings
(ii) Kelvin-Planck statement – It is impossible to take heat from a hotter reservoir and convert it completely into work by a cyclic process without transferring a part of heat to a cooler reservoir.
% Efficiency – \(\left[\frac{\mathrm{T}_{1}-\mathrm{T}_{2}}{\mathrm{~T}_{1}}\right] \times 100\)
Clausius statement – It states that heat flows spontaneously from hot objects to cold objects and to get it flow in the opposite direction, we have to expend some work.
Spontaneous process – A reaction that does occur under the given set of conditions without any external driving force is called a spontaneous reaction.
Non-spontaneous process – A reaction that does not occur under given set of conditions without any external driving force is called a non-spontaneous reaction.
Entropy (S) – It is defined as “for a reversible change taking place at a constant temperature (T), the change in entropy (AS) of the system is equal to heat energy absorbed or evolved (q) by the system divided by the constant temperature (T)”. \(\Delta S_{\text {sys }}=\frac{q_{\text {rev }}}{T}\)
Unit of entropy – JK-1(SI unit).
Standard entropy of formation – It is defined as the entropy of formation of 1 mole of a compound from the elements under standard conditions. It is denoted as \(\Delta \mathrm{S}_{f}^{0}\)
Standard entropy change AS0 –
\(\Delta \mathrm{S}_{r}^{0}=\Sigma \mathrm{S}_{\text {products }}^{0}-\Sigma \mathrm{S}_{\text {reactants }}^{0}\)
Entropy of fusion – When one mole of the solid melts at its melting point reversibly, the heat absorbed is called molar heat of fusion.
\(\Delta \mathrm{S}_{f}=\frac{\Delta \mathrm{H}_{f}}{\mathrm{~T}_{f}}\)
ΔHf = Molar heat of fusion
Tf= Melting point
Entropy of vapourisation – When one mole of liquid is boiled at its boiling point reversibly, the heat absorbed is called molar heat of vapourisation.
\(\Delta \mathrm{S}_{v}=\frac{\Delta \mathrm{H}_{v}}{\mathrm{~T}_{b}}\)
ΔHv = Molar heat of vapourisation
Tb = Boiling point
Entropy of transition – The heat change, when one mole of a solid changes reversibly from one allotropic form to another at its transition temperature.
\(\Delta \mathrm{S}_{t}=\frac{\Delta \mathrm{H}_{t}}{\mathrm{~T}_{t}}\)
ΔHt = Molar heat of transition .
Tt = Transition temperature
Gibbs free energy – It is the available energy to do the work. It is defined as the part of total energy of a system that can be converted (or) available for conversion into work.
G = H – TS
ΔG = ΔH – TΔS
ΔG = -w- PΔV
| Spontaneous (irreversible) |
Equilibrium (reversible) |
Non-spontaneous |
| ΔH < 0 | ΔH = TΔS | ΔH > 0 |
| ΔS > 0 | ΔS = ΔH/T | ΔS < 0 |
| ΔG < 0 | ΔG = 0 | ΔG > 0 |
Reaction quotient (Q) – It is defined as the ratio of concentration of the products to the concentrations of the reactants under non equilibrium conditions.
ΔG° = Standard free energy change of the reaction.
ΔG = ΔG° + RT In Q
ΔG0 = – RT In Keq
ΔG0 = – 2.303 RT log Keq
ΔG0 = ΔH0 – TΔS0 = – RT In Keq
Third law of thermodynamics – It states that the entropy of pure crystalline substance at absolute zero is zero (or) it is impossible to lower the temperature of an object to absolute zero in a finite number of steps, \(\lim _{T \rightarrow 0} S=0\)