Thermodynamics Chemistry notes of all topics with short explanation. You can revision your Board examination from here because The notes is written by our Expert Theacher with lots of experience. We provide Solution of NCERT question with important intext question
System : Specific part of universe in which thermodynamic observations are made.
Surroundings : Everything which surrounds the system.
Types of the System
- Open System : Exchange both matter and energy with the surroundings. For example : Reactants in an open test tube.
- Closed System : Exchange energy but no matter with the surroundings. For example : Reactants in a closed vessel.
- Isolated System : Neither exchange energy nor matter with the surroundings. For example : Reactants in a thermos flask. No system is perfectly isolated.
Thermodynamic Processes :
- Isothermal process : ∆T = 0
- Adiabatic process : ∆q = 0
- Isobaric process : ∆P = 0
- Isochoric process : ∆V = 0
- Cyclic process : ∆U = 0
- Reversible process : Process which proceeds infinitely slowly by a series of equilibrium steps.
- Irreversible process : Process which proceeds rapidly and the system does not have chance to achieve equilibrium.
Extensive Properties : Properties which depend upon the quantity or size of matter present in the system. For example : mass, volume, internal energy, enthalpy, heat capacity, work etc
Intensive Properties : Properties which do not depend upon the quantity or size of matter present in the system. For example : temperature, density, pressure, surface tension, viscosity, refractive index, boiling point, melting point etc.
State Functions : The variables of functions whose value depend only on the state of a system or they are path independent. For example : pressure (P), volume (V), temperature (T), enthalpy (H), free energy (G), internal energy (U), entropy (S), etc.
Internal Energy (U) : It is the sum of all kind of energies possessed by the system.
First Law of Thermodynamics : “The energy of an isolated system is constant.”
Mathematical Form : ∆U = q + w
Sign Conventions for Heat (q) and Work (w) :
- W = + ve, if work is done on system
- W = – ve, if work is done by system
- q = + ve, if heat is absorbed by the system
- q = – ve, if heat is evolved by the system
Work of Expansion/compression : w = – Pext (Vf — Vi )
Work done in Isothermal Reversible Expansion of an Ideal Gas :
wrev = – 2.303 nRT log Vf/Vi Or,
wrev = – 2.303 nRT log Pf/Pi
Significance of ∆H and ∆U : ∆H = qp and ∆U = qv
Relation between ∆H and ∆U
∆H = ∆U + (nl p – nr )RT for gaseous reaction.
- ∆H = ∆U if (n p – nr ) is zero; e.g., H2(g) + I 2(g) → 2HI(g)
- ∆H > ∆U if (n p – nr ) is positive; e.g., PC15(g) → PCl3(g) + C12(g)
- ∆H < ∆U if (n p – nr ) is negative; e.g., N2(g) + 3H2(g) → 2NH3(g)
Heat capacity (C) : Amount of heat required to raise the temperature ofl a substance by 1°C or 1 K.
q = C∆T
Specific heat capacity (Cs ) : Amount of heat required to raise the temperature of 1g of a substance by 1°C or 1K.
q = Cs × m × ∆T
Molar Heat Capacity (Cm) : Amount of heat required to raise thel temperature of 1 mole of a substance by 1°C or 1K.
q = Cm × n × ∆T
Standard State of a Substance : The standard state of a substance at a specified temperature is its, pure form at 1 bar.
Standard Enthalpy of Formation (∆f Hθ )
Enthalpy change accompanying the formation of one mole of a substance from its constituent elements under standard condition of temperature (normally 298 K) and pressure (1 bar).
- ∆f Hθ of an element in standard state is taken as zero.
- Compounds with – ve value of ∆f Hθ are more stable than their constituents.
- ∆r H° = Σi ai ∆f Hθ (products) – Σi bi ∆f Hθ (reactants) : Where ‘a’ and ‘b’ are coefficients of products and reactants in balanced equation.
Standard Enthalpy of Combustion (∆cHθ)
Enthalpy change accompanying the complete combustion of one mole of a substance under standard conditions (298 K, 1 bar) Hess’s Law of Constant Heat Summation : The total enthalpy changel of a reaction remains same whether it takes place in one step or in several steps.
Bond Dissociation Enthalpy : Enthalpy change when one mole of a gaseous covalent bond is broken to form products in gas phase. For example : Cl2(g) → 2Cl(g); ∆Cl-Cl Hθ = 242k/mol–1.
- For diatomic gaseous molecules; Bond enthalpy = Bond dissociation Enthalpy = Atomization Enthalpy.
- For Polyatomic gaseous molecules; Bond Enthalpy = Average of the bond dissociation enthalpies of the bonds of the same type. ∆rl Hθ = Σ∆bondHθ (Reactants) –– Σ∆bondHθ (Products).
Spontaneous Reaction : A reaction which can take place either of its ownl or under some initiation.
Entropy (S) : It is measure of degree of randomness or disorder of a system.
∆S sys = (qrev)sys/∆T = (∆H)sys/∆T Unit of Entropy = JK–1 mol–1
Second Law of Thermodynamics : For all the spontaneous processesl totally entropy change must be positive.
∆Stotal = ∆S sys + ∆Ssurr > 0
Gibbs Helmholtz Equation for determination of Spontaneity :
∆G = ∆H – T∆S
(i) If ∆G = – ve, the process is spontaneous
(ii) If ∆G = + ve, the process is non-spontaneous
(iii) If ∆G = 0, the process is in equilibrium Relation between Gibbs Energy Change and Equilibrium Constant:l ∆Gθ = – 2.303 RT log Kc .
Third law of thermodynamic : The entropy of a perfectly crystalline solid at absolute zero (0 K) is taken to be zero.