**Revision Notes on Chemical Thermodynamics**

**Basic Terminology:**

**Heat, energy and work:**

**Heat (Q):**

**•**Energy is exchanged between system and surround in the form of heat when they are at different temperatures.

**•**Heat added to a system is given by a positive sign, whereas heat extracted from a system is given negative sign.

**•**It is an extensive property.

**•**It is not a state function.

**Energy:**

**•**It is the capacity for doing work.

**•**Energy is an extensive property.

**•**Unit : Joule.

**Work (W):**

**•**Work = Force × Displacement i.e. dW = Fdx

**•**Work done on the system is given by positive sigh while work done by the system is given negative sign.

**•**Mechanical Work or Pressure-Volume Work: work associated with change in volume of a system against an external pressure.

**•**Work done in reversible process: W=

_{2}/v

_{1}= –2.303 nRT log p

_{1}/p

_{2}

**•**Wok done in isothermal reversible contraction of an ideal gas:

_{2}/v

_{1}= –2.303 nRT log p

_{1}/p

_{2}

**•**Unit : Joule.

**Internal Energy (E or U):**

**•**Sum of all the possible types of energy present in the system.

**•**ΔE = heat change for a reaction taking place at constant temperature and volume.

**•**ΔE is a state function.

**•**It is an extensive property.

**•**Value of ΔE is -ve for exothermic reactions while it is +ve for endothermic reactions.

**First Law of Thermodynamics:**

**Mathematical Expression**

**•**For an isolated system, dq=0

**•**For system involving mechanical work only

**•**At constant volume i.e. isochoric process

**•**For Isothermal Process

**•**For adiabatic process

**Enthalpy (H):**

**•**For system involving mechanical work only

**•**For exothermic reactions:

**•**For endothermic reactions:

**Relation between dH and dE:**

_{g}RT

**Heat capacity:**

**•**Amount of heat required to rise temperature of the system by one degree.

**•**

**Specific heat capacity:**Heat required to raise the temperature of 1 g of a substance by one dgree.

_{s}= Heat capacity / Mass in grams

**• Molar heat capacity:**Heat required to raise the temperature of 1 g of a substance by one dgree.

_{m}= Heat capacity / Molar mass.

**• Heat capacity of system at constant volume:**

_{v}= (dE/dT)

_{v}

**• Heat capacity of system at constant pressure:**

_{p}= (dE/dT)

_{p}

_{p}– C

_{v}= R

**• Variation Of Heat Of Reaction With Temperature:**

_{P}= (dH

_{2}- dH

_{1})/(T

_{2}-T

_{1}) & dCV = (dE

_{2}- dE

_{1})/(T

_{2}-T

_{1}

**• Bomb Calorimeter:**

**Enthalpies of Reactions:**

**Hess’s Law of constant heat summation:**

_{1}+ ΔH

_{2}

**Born–Haber Cycle:**

_{1}+ 1/2 ΔH

_{2}+ ΔH

_{3}+ ΔH

_{4}+ ΔH

_{5}= ΔH

_{f}(MX) (Lattice energy)

**Lattice energy:**The change in enthalpy that occurs when 1 mole of a solid crystalline substance is formed from its gaseous ions.

**Second law of thermodynamics**

**Statement:**

**Mathematically:**

**•**Entropy is the degree of randomness thus it increases with increase in randomness of particles of the systemi.e. ΔS is positive for melting of ice.

**•**At equilibrium, ΔS = 0

**•**For a spontaneous process, ΔS > 0

**Gibbs free energy(ΔG):**

**•**ΔG = ΔH - TΔS

**•**ΔG = nRT ln Keq

**•**ΔG = nFE

_{cell}

**•**At equilibrium, ΔG = 0

**•**For spontaneous process, ΔG < 0

**Bond Energies:**

_{2}(g) + ½O

_{2}(g) ΔH = 430 Kj

_{O–H}= (498 + 430)/2 = 464 kJ mol

^{–1}

**•**Efficiency of a heat engine (carnot cycle):

_{2}– T

_{1}) ln v

_{2}/v

_{1}

_{2}= RT

_{2}ln v

_{2}/v

_{1}

_{2}

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