THERMOCHEMISTRY

1. THERMOCHEMISTRY ENERGY CHANGES ASSOCIATED WITH CHEMICAL REACTION

2. ENERGY • Capacity to do work or supply heat • Kinetic Energy: KE = 1/2 mv2 = 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 = Efinal - Einitial • If ΔV = 0, then ΔE = qV • ΔE < 0, energy lost by system • ΔE > 0, energy gained by system

7. 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.

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 = qP • ΔH = Hfinal - Hinitial = HP - HR

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 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

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

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 oC; (units = J/oC-g) Table 6.1 • Cm = Molar Heat Capacity = amt of energy needed to raise temp. of 1 mol of sample 1 oC; (units = J/mol-oC) • q = s m ΔT or q = Cm 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 = 25oC 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 25oC. • ΔHo = 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/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.

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 • Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g) ΔH = -23 kJ • 2Fe(s) + 3CO2(g)  Fe2O3(s) + 3CO(g) ΔH = +23 kJ • 2Fe2O3(s) + 6CO(g)  4Fe(s) + 6CO2(g) ΔH = (2) -23 kJ = -46 kJ

25. Stoichiometry and Thermochemical Equations (2) • Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g) ΔH = -23 kJ per one mol Fe2O3(s) reacting • Calculate the heat given off if 500 g of Fe2O3(s) reacts with excess CO. g Fe2O3(s)  mol Fe2O3(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 • ΔHof (1 atm and 25 oC) values are tabulated in App. 4; note elements have ΔHof = 0. • Combine ΔHof to calculate heat of rxn. • ΔHorxn = ∑nPΔHof (prod.) - ∑nRΔHof (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.