Chapter 6

1. Chapter 6 • Energy • Law of Conservation of Energy • Potential vs Kinetic • Heat – transfer of Energy because of temp. difference

2. State function • System vs Surroundings • Exothermic • Endothermic

3. Thermodynamics • 1st law – Energy of Universe is constant • Internal energy • E = q + w • w = -p V

4. Enthalpy • H = qp • H = Hprod – Hreact • Calorimetry • Specific heat • q= sm T • Specific heat of water is 4.18 J/g0C

5. What is the specific heat of 15 g of an unkown at 45 0C if when mixed with 55 g of water at 25 0C the final temperature is 32 0C? • - q hot = q cold • - smTunk. = sm TH2O • - s(15g)(32-45 0C) = (4.18J/g0C)(55g)(32- 25 0C) • s = - (4.18J/g0C)(55g)(7 0C)/((15g)(-130C)) • s = 8 J/g0C

6. Constant volume • E = qv • Bomb Calorimeter • q = C T C = heat capacity of the calorimeter

7. Hess’s Law • Htotal = H1 + H2 + …. • If reaction is reversed change the sign • If you multiply the reaction’s coefficients you also multiply the H • N2 + O2 2NO H1 = 180 kJ • 2NO + O2  2NO2H2 = -112 kJ • ----------------------------------------------- • N2 + 2O2  2NO2H = H1 + H2 = 68 kJ

8. Cgraph. + O2 CO2H = -394 kJ • Cdiam. + O2  CO2H = -396 kJ • Cgraph. + O2 CO2H = -394 kJ • CO2  O2 + Cdiam. H = +396 kJ • --------------------------------------------- • Cgraph.  Cdiam. H = 2 kJ

9. Standard Enthalpies of Formation • Hfo • Values for formation of 1 mole at standard conditions • Reactants are all elements • Elements have a value of zero • From Appendix 4 on pg. A21 • H rxn =  nprodHfoprod -  nreactHforeact

10. H rxn =  nprodHfoprod -  nreactHforeact • CH4(g) + 2O2(g) CO2(g) + 2H2O(l) H = ? • H = 1n(-394kJ/n)+2n(-286kJ/n) – 1n(-75kJ/n) – 0 • H = -891 kJ