1. THERMOCHEMISTRY ENERGY CHANGES ASSOCIATED WITH CHEMICAL REACTION

2. ENERGY Capacity to do work or supply heat Kinetic Energy: KE = 1/2 mv 2 = energy due to motion (v ≠ 0), Joule is the unit Potential Energy: PE = stored energy due to position, energy in a chemical bond (recall endo and exo Expt 1), Joule Energy is conserved SI unit: Joule = kg (m/s) 2 ; 1 calorie = 4.184 J

3. HEAT Heat is the energy transfer between system (chem rxn of reactants and products = focus of study) and surroundings (everything else) due to temperature difference, Joule q > 0 if heat absorbed by chem rxn; endothermic. Fig 6.3 q < 0 if heat given off by chem rxn; exothermic. Fig 6.2 Heat is a path function

4. WORK Work is the energy transferred between system and surroundings, Joule w = F · d = force that moves object a distance d Consider work associated with gas expansion or contraction: w = -P ΔV where P = external pressure If w < 0, system does work on surroundings and system loses energy; e.g. gas expands If w > 0, surroundings does work on system and system gains energy; eg. gas is compressed Work is a path function Note that 1.00 (L-atm) = 101.3 J

5. Figure 6.4 The Piston, Moving a Distance Against a Pressure P, Does Work On the Surroundings

6. FIRST LAW OF THERMODYNAMICS The energy of the universe is constant; in a physical or chemical 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

7. 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.

8. ENTHALPY If a chem rxn occurs at constant pressure (ΔP = 0) and only PV work occurs, then the heat associated with this rxn is called enthalpy, Joule H = enthalpy = state function, tabulated in Appendix 4 H = E + PV; ΔH = ΔE + PΔV = q P ΔH = H final - H initial = H P - H R

9. ENTHALPY (2) ΔH < 0 energy lost by system, exothermic ΔH > 0 energy gained by system, endothermic Enthalpy depends on amount of substance (I.e. #mol, #g); extensive property. Chemical rxns are accompanied by enthalpy changes (ΔH can be > 0 and < 0) that are measurable and unique.

10. Figure 6.2 Exothermic Process

11. Figure 6.3 Endothermic Process

12. Problems 24, 28, 30, 34, 36

13. 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

14. CALORIMETRY Experimental method of determining heat (q) absorbed or released during a chem. rxn. Expts are either done at constant P (q P = ΔH) or constant V (q 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 = heat capacity of the calorimeter; J/ o C

15. CALORIMETRY (2) Here are two expressions of heat capacity s = specific heat (capacity) = amount of energy needed to raise the temp. of 1 g of material 1 o C; (units = J/ o C-g) Table 6.1 C m = Molar Heat Capacity = amt of energy needed to raise temp. of 1 mol of sample 1 o C; (units = J/mol- o C) q = s m ΔT or q = C m n ΔT

16. Table 6.1 The Specific Heat Capacities of Some Common Substances

17. Figure 6.5 A Coffee-Cup Calorimeter Made of Two Styrofoam Cups

18. Figure 6.6 A Bomb Calorimeter.

19. Problems 42, 46, 48, 54

20. THERMODYNAMIC STANDARD STATE The standard or reference state of a pure compound is its state at T = 25 o C and –P = 1.00 atm for a gas or –1.00 M concentration for a solution. For an element, the std state is 1 atm and 25 o C. ΔH o = standard enthalpy of rxn or heat of rxn when products and reactants are in their standard states.

21. PHYSICAL CHANGES There are ΔH values associated with phase or 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.

22. 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. (p 246)

23. The Principle of Hess’s Law

24. Stoichiometry and Thermochemical Equations Fe 2 O 3 (s) + 3CO(g)  2Fe(s) + 3CO 2 (g) ΔH = -23 kJ 2Fe(s) + 3CO 2 (g)  Fe 2 O 3 (s) + 3CO(g) ΔH = +23 kJ 2Fe 2 O 3 (s) + 6CO(g)  4Fe(s) + 6CO 2 (g) ΔH = (2) -23 kJ = -46 kJ

25. Stoichiometry and Thermochemical Equations (2) Fe 2 O 3 (s) + 3CO(g)  2Fe(s) + 3CO 2 (g) ΔH = -23 kJ per one mol Fe 2 O 3 (s) reacting Calculate the heat given off if 500 g of Fe 2 O 3 (s) reacts with excess CO. g Fe 2 O 3 (s)  mol Fe 2 O 3 (s)  heat given off

26. STANDARD ENTHALPY 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. 4; note elements have ΔH o f = 0. Combine ΔH o f to calculate heat of rxn. ΔH o rxn = ∑n P ΔH o f (prod.) - ∑n R ΔH o f (react.)

27. Table 6.2 Standard Enthalpies of Formation for Several Compounds at 25°C

28. Problems 58, 60, 66, 72

29. ENERGY SOURCES Variety of and emerging sources of energy and preparation of fuels Impact on the environment Combustion = type of reaction in which substance burns in oxygen.