1. Chemical Thermodynamics Lecture 1. Chemical Thermodynamics

2. Chemical Thermodynamics Thermochemistry is the study of heat change in chemical reactions. The system is the specific part of the universe that is of interest in the study. open mass & energy Exchange: closed energy isolated nothing

3. Chemical Thermodynamics First Law of Thermodynamics Q= ΔU + W The first law of thermodynamics is an expression of the principle of conservation of energy. The law states that energy can be transformed, i.e. changed from one form to another, but cannot be created nor destroyed.conservation of energy

4. Chemical Thermodynamics Energy Internal energy = kinetic + potential energy of the molecules or atoms of a body. Kinetic energy = translation, rotation, vibration of the molecules or atoms Units of Energy SI Unit for energy is the joule, J. Sometimes the calorie is used instead of the joule: 1 cal = 4.184 J (exactly)

5. Chemical Thermodynamics Isochoric Process (courtesy F. Remer) V= const, W=0 Q v = U 2 -U 1 =  U p= const Q p =  U + p  V =  H Isobaric Process Isothermal Process T= const Adiabatic Process Q= const

6. Chemical Thermodynamics Exothermic and Endothermic Processes Endothermic: absorbs heat from the surroundings. An endothermic reaction feels cold,. H > 0 Exothermic: transfers heat to the surroundings. An exothermic reaction feels hot, H < 0 (combustion). The First Law of Thermodynamics

7. Chemical Thermodynamics Heat effect of reaction is a state function: depends only on the initial and final states of system, not on how the internal energy is used. The First Law of Thermodynamics

8. Chemical Thermodynamics Hess’s law: if a reaction is carried out in a number of steps,  H for the overall reaction is the sum of  H for each individual step. For example: CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g)  H = -802 kJ 2H 2 O(g)  2H 2 O(l)  H= - 88 kJ CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l)  H = -890 kJ Hess’s Law

9. Chemical Thermodynamics f Standard enthalpy of formation (  H 0 ) is the heat change that results when one mole of a compound is formed from its elements at a pressure of 1 atm. f The standard enthalpy of formation of any element in its most stable form is zero.  H 0 (O 2 ) = 0 f  H 0 (O 3 ) = 42 kJ/mol f  H 0 (C, graphite) = 0 f  H 0 (C, diamond) = 1.90 kJ/mol f 6.5

10. Chemical Thermodynamics The standard enthalpy of reaction (  H 0 ) rxn aA + bB cC + dD H0H0 rxn d  H 0 (D) f c  H 0 (C) f = [+] - b  H 0 (B) f a  H 0 (A) f [+] H0H0 rxn n  H 0 (products) f =  m  H 0 (reactants) f  -

11. Chemical Thermodynamics Spontaneous Processes Spontaneous processes are those that can proceed without any outside intervention.

12. Chemical Thermodynamics Entropy S = k lnW where k is the Boltzmann constant, 1.38  10  23 J/K. W- number of microstates. Entropy can be thought of as a measure of the randomness of a system. Entropy increases with the freedom of motion of molecules. Therefore,S(g) > S(l) > S(s)

13. Chemical Thermodynamics Second Law of Thermodynamics Entropy is a state function. Therefore,  S = S final  S initial The second law of thermodynamics: The entropy of the universe does not change for reversible processes and increases for spontaneous processes. Irreversible (real, spontaneous):

14. Chemical Thermodynamics Third Law of Thermodynamics The entropy of a pure crystalline substance at absolute zero is 0.

15. Chemical Thermodynamics Gibbs Free Energy

16. Chemical Thermodynamics Gibbs Free Energy 1.If  G is negative, the forward reaction is spontaneous. 2.If  G is 0, the system is at equilibrium. 3.If  G is positive, the forward reaction is nonspontaneous.

17. Chemical Thermodynamics Free Energy and Equilibrium Remember from above: If  G is 0, the system is at equilibrium. So  G must be related to the equilibrium constant, K.The standard free energy,  G°, is directly linked to K eq by: