RBSE Class 11 Chemistry Notes Chapter 6 Thermodynamics

These comprehensive RBSE Class 11 Chemistry Notes Chapter 6 Thermodynamics will give a brief overview of all the concepts.

RBSE Class 11 Chemistry Chapter 6 Notes Thermodynamics

→ Thermodynamics: It deals with energy change in chemical or physical process and enable us to study these changes quantitatively and to make useful predictions.

→ System: It is the specific part of universe in which transformation of heat/energy takes place and selected for studying the thermodynamic observations.

→ Surroundings: The remaining part of universe which surrounds the system is called surrounding.

→ Open System: In an open system, there is exchange of energy and matter between system and surroundings.

→ Closed System: In closed system, there is no exchange of matter but exchange of energy is possible between the system and surroundings.

→ Isolated System: A system in which no exchange of energy and matter between system and surroundings takes place.

→ Homogeneous System: In this system, all matters are in same phase.

→ Heterogeneous System: In this system, the matter is present in two or more phases.

→ State Function: Those thermodynamic properties of a system which depend only upon the state of system and is independent of the path by which state of system is changed.

→ State Variable: These are variables like temperature, pressure, volume and composition which represent the state of any system.

→ Path Function: Those thermodynamic properties of-any system whose values depend upon path.

→ Isothermal Process: It is the process in which temperature remains constant throughout the process.

→ Adiabatic Process: When a process is carried out in such a way that no heat can flow from the system to surroundings or vice versa.

→ Isobaric Process: It is the process in which pressure remains constant throughout the process.

→ Isochoric Process: A process in which the volume remains constant throughout the process.

→ Cyclic Process: In this process, system reaches to its initial state passing through different states there is no change throughout the process.

→ Reversible Process: A process in which direction may be reversed due to its very slow speed and system remains virtually in the state of equilibrium with the surroundings.

→ Irreversible Process: Process which is not reversible and this takes place rapidly so that successive steps of the direct process can not be retracted and any change in the external condition disturb the equilibrium.

→ Macroscopic System: A system which is composed of chemical species such as atoms, ions or molecules, is known as macroscopic system.

→ Macroscopic Properties: The properties of system which are produced by combined behaviour of chemical species are known as macroscopic properties.

→ Extensive Properties: Those properties whose magnitude depends upon amount of substance present in system.

→ Intensive Properties: Properties which do not depend upon the amount of substances present in the system.

→ Internal Energy: Every system has some quantity of matter associated with a definite amount of energy, this is known as internal energy. ,

→ Mechanical Work: This work is related to the volume of any system which is against external pressure, this work is done against mechanical work.

→ Electrical Work: Work which is used to carry a charged mass from one potential to other, it is against electrical force.

→ Gravitational Force: Work which is done to carry a mass to certain height against gravity, it is against gravitational force.

→ Free Expansion: The expansion of gaseous system in vaccum (P_ex = 0) is called free expansion.

→ Zeroth Law of Thermodynamics: If two objects are in thermal equilibrium then their hotness will be equal.

→ First Law of Thermodynamics: Energy can neither be created nor destroyed although it can converted from one form into another.

→ Enthalpy: At constant pressure, the total heat absorbed of a system is equal to internal energy and work PV.

→ Exothermic Reactions: Those chemical reactions in which a certain amount of heat is released.

→ Endothermic Reactions: Those chemical reactions in which a certain amount of heat is absorbed.

→ Heat Capacity: Amount of heat required to raise the temperature by 1°C or 1 K of any substance.

→ Molar Heat Capacity: Amount of heat required to raise the temperature by 1°C of 1 K of 1 mole of substance is known as molar heat capacity.

→ Specific Heat Capacity: Amount of heat required to raise the temperature by 1°C or 1 K of 1 g of substance.

→ Heat Capacity at constant Volume: It is the amount of heat required to raise the temperature by 1°C or 1 K of any substance at constant volume.

→ Heat Capacity at Constant Pressure: At constant pressure, the amount of heat required to raise the' temperature by 1°C or IK of any substance.

→ Heat of Reaction: The enthalpy change in a reaction when the number of moles of reactants react completely to give the product, at constant temperature and pressure as given by balanced chemical reaction is called enthalpy of reaction.

→ Enthalpy of Fusion: The enthalpy change when 1 mole of solid substance is converted into its liquid state at its melting point, it is also called molar enthalpy of fusion.

→ Enthalpy of Vaporisation: The enthalpy change when 1 mole of a liquid is converted into gaseous state at its boiling point.

→ Enthalpy of Sublimation: The enthalpy change when 1 mole of a solid is directly converted into its gaseous state at a temperature below its melting point.

→ Enthalpy of Transition: The enthalpy change when 1 mole of solid substance is converted from a crystalline state to another crystalline state.

→ Enthalpy of Formation: The standard enthalpy change for the formation of one mole of a compound from its elements in their most stable state of aggregations is called standard enthalpy of formation.

→ Enthalpy of Combustion: The enthalpy change when one mole of a substance undergoes complete oxidation is called enthalpy of combustion.

→ Enthalpy of Neutralization: One gram equivalent of an acid is neutralised with one gram equivalent of a base in dilute solution. The heat evolved in this process is called enthalpy of neutralisation. .

→ Enthalpy of Solution: The enthalpy change when 1 mole of substance dissolves in a specified amount of solvent.

→ Enthalpy of Hydration: The enthalpy change due to released or absorbed amount of heat when 1 mole of dry substance is completely hydrated.

→ Lavosier and Laplace Law: According to this law, as much amount of energy is released in formation of compound from its constituents, same amount of energy is required to decompose its constituents.

→ 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 into which the overall reaction may be divided at the same temperature.

→ Enthalpy of Atomsiation: The enthalpy change on breaking of 1 mole of bonds completely to obtain atoms in the gaseous phase is called enthalpy of atomisation.

→ Lattice Enthalpy: Lattice enthalpy of an ionic compound is the enthalpy change which occurs when 1 mole of ionic compound dissociates into its ions in gaseous state.

→ Born-Haber Cycle: It is a cycle which is used to determine the lattice energy of any ionic compound.

→ Bond Dissociation Enthalpy: The change in enthalpy when 1 mole of gaseous covalent compound converts into gaseous product.

→ Spontaneous Process: The process which takes place without the help of any external agent is called spontaneous process.

→ Driving Force: The force which is responsible for spontaneity of process.

→ Entropy: It is the measurement of state of any system.

→ Molar Entropy of Fusion: When 1 mole of solid changes into liquid at its melting point then change in entropy is called molar entropy of fusion.

→ Molar Entropy of Vaporisation: When 1 mole of liquid converts into vapour at a temperature lower than its boiling point then change in entropy during this process is known as molar entropy of vaporisation.

→ Molar Entropy of Sublimation: When 1 mole of solid converts into vapour at a temperature lower than its melting point then change in entropy is called molar entropy of vaporisation.

→ Second Law of Thermodynamics: The entropy of universe always increases in the course of every spontaneous change.

→ Gibb’s Free Energy: It is the amount of energy required to work done at constant temperature and pressure.

→ Third law of Thermodynamics: The entropy of any solid or liquid approaches to zero at absolute zero temperature.

→ Lewis and Randall’s Statement: Each substance has positive entropy but at absolute temperature it may be zero and for purely crystalline solid it is zero.

→ Nernst’s Statement: At absolute zero, for reaction of purely crystalline substances change in entropy is zero.

→ Other Important Facts and Formulae

• Heat and work are the main forms of energy which exchange between system and surroundings,
• According to first law of thermodynamics
  ΔU = q + W
  or ΔU = q - PΔV
• Work done by system is -ve while work done on the system is +ve.
• The heat taken by system is +ve while heat given by system is -ve.
• Relation between internal energy and change in enthalpy is:
  (a) ΔH = ΔU + PΔV
  (b) ΔH = ΔU + Δn_gRT
• Standard enthalpy change (ΔH⁰) is measured at 298 K and 1 atm.
• In exothermic reactions, energy is released and AH is negative.
• In endothermic reactions, energy is absorbed and AH is positive.
• ΔrH = ΣΔ_fH⁰_(P) - ΣΔ_fH⁰_(R)
• Energy is released during combustion.
• ΔH = Bond energy of reactants - Bond energy of products.
• Entropy ΔS = +ve (Spontaneous process)
  ΔS = -ve (Non-spontaneous process)
  ΔS = 0 (Equilibrium state)
• Δ_fS⁰ = ΣΔ_fS⁰_(P) - ΣΔ_fS⁰_(R)
• ΔS = \(\frac{q_{\mathrm{rev}}}{T}\)

→ Entropy (S) determines the quantities of randomness.

• ΔS_(Fusion) = S_(liq) - S_(solid) = \(\frac{\Delta H_{\text {fus }}}{T_f}\)
• ΔS_vap = S_vap - S_liq = \(\frac{\Delta H_{\mathrm{vap}}}{T_b}\)
• ΔS_sub = S_vap - S_sol = \(\frac{\Delta H_{\mathrm{sub}}}{T}\)

→ Increasing order of entropy is
Crystalline solid < Amorphous solid < Liquid < Gas < Gaseous mixture

→ According to Gibb’s-Helmoltz equation
ΔG = ΔH - TΔS

→ According to third law of Thermodynamics
∴ ΔS = 2.303 C_p logT

→ For isothermal process
-ΔS_T = 2.303 nR log \(\frac{V_2}{V_1}\)
ΔS_T = 2.303 nR log \(\frac{P_2}{P_1}\)

→ For isobaric processs
ΔS_p = 2.303 ⁿC_p log\(\frac{T_2}{T_1}\)

→ For isochoric process,
ΔS_v = 2.303 ⁿC_v log\(\frac{T_2}{T_1}\)

→ Relation between Gibb’s free energy and work
ΔS_v = W_{non expand}

→ If applied work is electrical then
ΔG⁰ = - nFE_cell⁰
ΔG⁰ = ΣΔ_fG_(P)⁰ - ΣΔ_fG_(R)⁰

→ ΔG⁰ = - 2.303 ETlogK (K = equilibrium constant),

→ For reversible process, at constant temperature
w_rev = -2.303 nRT log\(\frac{V_2}{V_1}\)

→ For pressure-volume work, W = - PΔV.

→ Free expansion, P_ex = 0.

→ For isothermal irreversible process,
q = -W = P_ex(V₂ - V₁).
= 2.303 nRT log \(\frac{P_2}{P_1}\)

→ For isothermal reversible process,

→ For adiabatic process, q = 0, ΔU = W_ad.

→ For isothermal process, ΔT = 0, for isobaric process ΔP = 0 and for isochoric process ΔV = 0.
At constant pressure, C_p = \(\left(\frac{\Delta H}{\Delta T}\right)_P\)
At constant volume C_v = \(\left(\frac{\Delta U}{\Delta T}\right)_V\)

→ C_p - C_v = R (gas constant).

→ ΔU is determined by Bomb calorimeter,

→ Hess’s law, ΔH = ΔH₁ + ΔH₂ + ΔH₃ ....

→ Efficiency of Fuel cell = \(\frac{\Delta G}{\Delta H}\).

→ Nature of gas can be determined with the help of \(\frac{C_p}{C_v}\) ratio \(\frac{C_p}{C_v}\)= γ (constant).
For monoatomic gas γ = 1.66.
For diatomic gas γ = 1.40.

→ q = m × c × ΔT.

→ For cyclic process, ΔU = 0, ΔH = 0.

→ Heat at constant volume is known as internal energy
ΔU = q_v.

→ Heat at constant pressure is called enthalpy ΔH = q_p.