Thermo

1. Thermo And all that implies The Review

2. 1st Law of Thermodynamics • Energy is neither created nor destroyed • ΔE = q + w • q is heat • Heat transferred into the system q > 0 (+) • Heat transferred out of the system q < 0 (-) • w is work • Work done on the system w > 0 (+) • Word done by the system w < 0 (-)

3. Example • Consider the combustion process that occurs in the cylinder of an automobile: 2C8H18(l) + 25O2(g) 16CO2(g) + 18H2O(g) • because the reaction produces a greater amount of gas than is consumed, not to mention gives off excessive heat which causes the product gases to expand, the reaction pushes the piston upward against the force of gravity and the tension of the camshaft (or something like that). The point is that this process involves some work. • So what is the system � • The reactants and products • What are the surroundings � • The piston, cylinder, engine + the rest of the universe • What about q? • q < 0 • The reaction gives off heat which is dissipated in the engine (thus the need for radiator coolant). • What about w? • w < 0 • The system (octane, oxygen, CO2 and H2O) does work on the surroundings by pushing the piston upwards.

4. ΔH • Enthalpy- heat of reaction • Heat content at constant pressure • ∆H = Hproducts - Hreactants (this is given) • + enthalpy = endothermic • - enthalpy = exothermic

5. S(s) + O2 (g)  SO2 (g) ∆H = -296 kJ/mol • calculate the heat evolved when 275 g S is burned 275 g 1mol S 296 kJ = -2535.8 kJ 32.1 g S 1 mol = 2.54 X 10-3 kJ b) calculate the heat evolved when 150 g SO2 is produced 150 g SO2 1 mol SO2 1 mol S 296 kJ = 693 kJ 64.1 g SO2 1 mol SO2 1 mol S

6. Calorimetry Science of measuring heat Calorimeter constant pressure (coffee cup) or constant volume (bomb) Heat capacity C = heat absorbed (not given) increase in temperature • Specific heat capacity- energy required to raise the temperature of one gram of a substance by one degree Celsius q = m c  T (GIVEN) • Molar heat capacity-energy required to raise the temperature of one mole of a substance by one degree Celsius

7. Other H formulas • ΔH˚ = ΣΔH˚ products - ΣΔH˚ reactants (given) • ΔH = Σbond broken (reactants)- Σ bonds formed (products) (not given)

8. ENTROPY ‘S’The Second Law of Thermodynamics The universe is constantly increasing disorder. DSuniv = DSsystem + DSsurroundings ΔSsurroundingsbased on heat flow exothermic + DSsurr endothermic -DSsurr

9. ENTROPY (S) Disorder of a system (more disorder is favored) Nature tends toward chaos! Think about your room at the end of the week! Factors that can indicate entropy changes: • Phase changes • Temperature changes • Volume changes • Mixing substances • Change in number of particles • Change in moles of gas

10. Predicting the entropy of a system based on physical evidence: • The greater the disorder or randomness in a system, the larger the entropy. • The entropy of a substance always increases as it changes from solid to liquid to gas. • When a pure solid or liquid dissolves in a solvent, the entropy of the substance increases (carbonates are an exception! --they interact with water and actually bring MORE order to the system)

11. Predicting the entropy of a system based on physical evidence: • When a gas molecule escapes from a solvent, the entropy increases. • Entropy generally increases with increasing molecular complexity (crystal structure: KCl vs CaCl2) since there are more MOVING electrons! • Reactions increasing the number of moles of particles often increase entropy.

12. ENTROPYThe Third Law of Thermodynamics • The entropy of a perfect crystal at 0 K is zero. • not a lot of perfect crystals out there so, entropy values are RARELY ever zero—even elements

13. BIG MAMMA, verse 2 S°rxn = S°(products) - S°(reactants) S is + when disorder increases (favored) S is – when disorder decreases Units are usually J/K mol (not kJ ---tricky!)

14. S surr = - H T Give signs to ΔH following exo/endo guidelines! (If reaction is exo.; entropy of surroundings increases—makes sense!)

15. Whether a reaction will occur spontaneously may be determined by looking at the S of the universe. ΔS system + ΔS surroundings = ΔS universe • IF ΔS universe is +, then reaction is spontaneous • IF ΔS universe is -, then reaction is NONspontaneous

16. FREE ENERGY (ΔG) • Calculation of Gibb’s free energy is what ultimately decides whether a reaction is spontaneous or not. • NEGATIVE G’s are spontaneous.

17. G can be calculated one of several ways: Gºrxn = Go(products) - Go(reactants) (Given) ΔG°f = 0 (for elements in standard state) G = H - TS (given)

18. G = Go + RT ln (Q) (not given) Define terms: G = free energy not at standard conditions Go = free energy at standard conditions R = universal gas constant 8.3145 J/molK T = temp. in Kelvin ln = natural log Q = reaction quotient: (for gases this is the partial pressures of the products divided by the partial pressures of the reactants—all raised to the power of their coefficients) Q = [products] [reactants]

19. “RatLink”: (given) G = -RTlnK Terms: Basically the same as above --- however, here the system is at equilibrium, so G = 0 and K represents the equilibrium constant under standard conditions. K = [products] still raised to power of coefficients [reactants]

20. “nFe”: G = - nFE (given) Terms: Go = just like above—standard free energy n = number of moles of electrons transferred (look at ½ reactions) F = Faraday’s constant 96,485 Coulombs/mole electrons Eo= standard voltage ** one volt = joule/coulomb**

23. Heating Curves

24. Potential Energy Diagrams • Endo • Exo