THERMOCHEMISTRY

1. THERMOCHEMISTRY ENERGY CHANGES ASSOCIATED WITH CHEMICAL REACTION

2. ENERGY • Capacity to do work or supply heat • Kinetic Energy: EK = 1/2 mv2 = energy due to motion, Joule • Potential Energy: EP = stored energy due to position, energy in a chemical bond, Joule • Energy is conserved (Fig 8.1) • SI unit: Joule = kg (m/s)2; 1 calorie = 4.184 Joule

3. HEAT • Energy transfer between system (chem rxn of reactants and products) and surroundings (everything else) due to temperature difference, Joule • q > 0 if heat absorbed by chem rxn; endothermic • q < 0 if heat given off by chem rxn; exothermic

4. WORK • Energy transfer between system and surroundings, Joule • w = F · d = force that moves object a distance d • w = -P ΔV where P = external pressure • If w < 0, gas expands, system loses energy • If w > 0, gas is compressed, system gains energy

5. FIRST LAW OF THERMODYNAMICS • Total energy of an isolated system is constant; in a phys. or chem. change, energy is exchanged between system and surroundings, but not created nor destroyed. • ΔE = internal energy = q + w = Efinal - Einitial • If ΔV = 0, then ΔE = qV • ΔE < 0, energy lost by system • ΔE > 0, energy gained by system

6. STATE FUNCTIONPATH FUNCTION • State Function: A property of the system which depends only on the present state of the system and not the path used to get there; E, V, T • Path Function; a property that depends on path taken during the change; w and q. • Note ΔE = w + q is a constant for specific initial and final states even though q and w are path functions.

7. ENTHALPY • If the reaction occurs at constant pressure, heat associated with rxn = enthalpy, Joule • H = state function, tabulated in B1, B2 • H = E + PV; ΔH = ΔE + P ΔV = qP • ΔH = Hfinal - Hinitial = HP - HR • ΔH < 0 energy lost by system, exothermic • ΔH > 0 energy gained by system, endothermic

8. ENTHALPY (2) • Enthalpy depends on amount of substance (I.e. #mol, #g); extensive property. • Chemical rxns are accompanied by enthalpy changes that are measurable and unique.

9. THERMOCHEMICAL EQUATION • Balanced chemical equation at a specific T and P includes reactants, products, phases andΔH . • Basis for stoichiometric problems that focus on ΔH associated with the chemical rxn. • ΔH for reverse rxn =- ΔH for forward rxn • If amount of reactants or products changes, then ΔH changes

10. THERMODYNAMIC STANDARD STATE • We define the standard state of a substance as its most stable state at T = 25oC, P = 1 atm (or 1 bar) and concentration = 1 M. • ΔHo = standard enthalpy of rxn or heat of rxn when products and reactants are in their standard states.

11. PHYSICAL CHANGES • Melting/freezing solid / liquid • Boiling/condensing liquid / vapor • Subliming/condensing solid / vapor • The former changes are endothermic; the latter are exothermic. • Note that these changes are reversible.

12. CALORIMETRY • Experimental method of determining heat (q) absorbed or released during a chem. rxn. at constant P (ΔH) or constant V (ΔE). • This heat is proportional to the temp. change during the rxn: q = C ΔT where C is a constant and ΔT = Tfinal - Tinitial. • C is the heat capacity of the calorimeter; J/oC

13. CALORIMETRY (2) • s = specific heat capacity = amount of energy needed to raise the temp. of 1 g of material 1 oC; (s has units of J/oC-g) T 8.1 • Cm = Molar Heat Capacity = amt of energy needed to raise temp. of 1 mol of sample 1 oC ( J/mol-oC) • q = s m ΔT or q = Cm n ΔT • If ΔP = 0, then ΔH = q; if ΔV = 0, then ΔE = q

14. HESS’S LAW: Law of Heat Summation • Given a specific chem rxn at a stated T and P values, ΔH for the chem rxn is • constant and not dependent on intermediate chem rxns. • the sum of the enthalpy changes for the intermediate rxns. (Chem eqns are additive and their associated rxn ΔH values are additive). • Hess’s Law facilitates the determination of rxn enthalpies for numerous rxns.

15. STANDARD HEAT OF FORMATION • Enthalpy change for the formation of one mole of a substance in its standard state from its elements in their standard states • ΔHof (1 atm and 25 oC) values are tabulated in App. B; note elements have ΔHof = 0. • Combine ΔHof to calculate heat of rxn. • ΔHorxn = ∑ΔHof (prod.) - ∑ ΔHof (react.)

16. BOND DISSOCIATION ENERGY • We can use bond dissociation energies to approximate heat of rxn (recall prob 7.110) • ΔHo = D(react bonds) – D(prod bonds) • D values in T7.1 • D values are positive (bond breaking requires energy and bond formation releases energy) and are given in kJ/mol

17. COMBUSTION • Type of reaction in which substance burns in oxygen.

18. ENTROPY • Achieving stability has been related to minimizing energy; e.g. molecular geometry. But there is another important property called entropy (S) that affects the direction of chem rxn. • Entropy is a measure of disorder. Processes proceed spontaneously in the direction that increases disorder.

19. ENTROPY (2) • Therefore, spontaneous processes are favored by a decrease in enthalpy (energy when pressure is constant) and an increase in entropy. • Mathematically, this means ΔH < 0 (exothermic) and ΔS > 0 favor R  P • Units of J/oC-mol

20. FREE ENERGY • A new property, Gibbs Free Energy, combines the contributions of ΔH and ΔS. • ΔG = ΔH - T ΔS • ΔG determines direction of rxn • Units of kJ/mol