THERMOCHEMISTRY ENERGY CHANGES ASSOCIATED WITH CHEMICAL REACTION

ENERGY Capacity to do work or supply heat Kinetic Energy: E K = 1/2 mv 2 = energy due to motion, Joule Potential Energy: E P = 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 = Joule

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

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

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 = E final - E initial If ΔV = 0, then ΔE = q V ΔE < 0, energy lost by system ΔE > 0, energy gained by system

STATE FUNCTION PATH 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.

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 = q P ΔH = H final - H initial = H P - H R ΔH < 0 energy lost by system, exothermic ΔH > 0 energy gained by system, endothermic

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.

THERMOCHEMICAL EQUATION Balanced chemical equation at a specific T and P includes reactants, products, phases and ΔH. 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

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

PHYSICAL CHANGES Melting/freezingsolid  /  liquid Boiling/condensingliquid  /  vapor Subliming/condensingsolid  /  vapor The former changes are endothermic; the latter are exothermic. Note that these changes are reversible.

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 = T final - T initial. C is the heat capacity of the calorimeter; J/ o C

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

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.

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 ΔH o f (1 atm and 25 o C) values are tabulated in App. B; note elements have ΔH o f = 0. Combine ΔH o f to calculate heat of rxn. ΔH o rxn = ∑ ΔH o f (prod.) - ∑ ΔH o f (react.)

BOND DISSOCIATION ENERGY We can use bond dissociation energies to approximate heat of rxn (recall prob 7.110) ΔH o = 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

COMBUSTION Type of reaction in which substance burns in oxygen.

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.

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 favor R  P Units of J/ o C-mol

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 ΔGΔGSpontaneous rxn < 0 R  P = 0equilibrium > 0 R  P