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, Joule Potential Energy: PE = stored energy due to position, energy in a chemical bond, Joule Energy is conserved SI unit: Joule = kg (m/s) 2 ; 1 calorie = 4.184 Joule

3. HEAT 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 Path function

4. WORK Energy transfer 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 Path function Note 1 (L-atm) = 101.3 J

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

6. FIRST LAW OF THERMODYNAMICS The energy of the universe 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

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 the rxn occurs at ΔP = 0 and only PV work occurs, heat associated with rxn = enthalpy, Joule H = state function, tabulated in Appendix 4 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

9. Figure 6.2 Exothermic Process

10. Figure 6.3 Endothermic Process

11. ENTHALPY (2) 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.

12. 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

13. 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

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

15. CALORIMETRY (2) 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. Figure 6.5 A Coffee- Cup Calorimete r Made of Two Styrofoam Cups

17. Figure 6.6 A Bomb Calorimeter.

18. THERMODYNAMIC STANDARD STATE The standard or reference state of a pure compound is its state at T = 25 o C, P = 1 atm for a gas and concentration = 1 M 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.

19. 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.

20. 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)

21. The Principle of Hess’s Law

22. 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.)

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

24. 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.