1. Bond Enthalpies How does a chemical reaction have energy?

2. Bond Energy  Energy required to make/break a chemical bond  Endothermic reactions  Products have more energy than reactants  More energy to BREAK bonds  Exothermic reactions  Reactants have more energy than products  More energy to FORM bonds

3. Bond Enthalpy  Focuses on the energy/heat between products and reactants as it relates to chemical bonding  Amount of energy absorbed to break a chemical bond--- amount of energy released to form a bond.  Multiple chemical bonds take more energy to break and release more energy at formation  Amount of energy absorbed = amount of energy released to break chemical bond to form a chemical bond

4. Calculating ΔH rxn. by bond enthalpies (4 th method)  Least accurate method  ΔH = ΣBE (bonds broken) - ΣBE (bonds formed)

5. Example 1:  Using average bond enthalpy data, calcaulate ΔH for the following reaction.  CH 4 + 2O 2  CO 2 + 2H 2 O ΔH = ? BondAverage Bond Enthalpy C-H413 kJ/mol O=O495 kJ/mol C-O358 kJ/mol C=O799 kJ/mol O-H467 kJ/mol

6. Entropy

7. Spontaneous vs. Nonspontaneous 1)Spontaneous Process  Occurs WITHOUT help outside of the system, natural  Many are exothermic—favors energy release to create an energy reduction after a chemical reaction  Ex. Rusting iron with O 2 and H 2 O, cold coffee in a mug  Some are endothermic  Ex. Evaporation of water/boiling, NaCl dissolving in water

8. Spontaneous vs. Nonspontaneous 2) Nonspontaneous Process  REQUIRES help outside system to perform chemical reaction, gets aid from environment  Ex. Water cannot freeze at standard conditions (25°C, 1atm), cannot boil at 25°C **Chemical processes that are spontaneous have a nonspontaneous process in reverse **

9. Entropy (S)  Measure of a system’s disorder  Disorder is more favorable than order  ΔS = S (products) - S (reactants)  ΔS is (+) with increased disorder  State function  Only dependent on initial and final states of a reaction  Ex. Evaporation, dissolving, dirty house

10. Thermodynamic Laws 1 st Law of Thermodynamics  Energy cannot be created or destroyed 2 nd Law of Thermodynamics  The entropy of the universe is always increasing.  Naturally favors a disordered state

11. When does a system become MORE disordered from a chemical reaction? (ΔS > 0) 1)Melting 2)Vaporization 3)More particles present in the products than the reactants  4C 3 H 5 N 3 O 9 (l)  6N 2 (g) + 12CO 2 (g) + 10H 2 O (g) + O 2 (g) 4)Solution formation with liquids and solids 5)Addition of heat

12. When does a system become LESS disordered from a chemical reaction? (ΔS < 0) 1)Solution formation with liquids and gases

13. 3 rd Law of Thermodynamics The entropy (ΔS) of a perfect crystal is 0 at a temperature of absolute zero (0°K).  No particle motion at all in crystal structure  All motion stops

14. How do we determine if a chemical reaction is spontaneous? 1)Change in entropy (ΔS) 2)Gibbs Free Energy (ΔG)

15. Change in entropy (ΔS)  For a chemical reaction to be spontaneous (ΔS T > 0), there MUST be an increase in system’s entropy (Δs sys > 0) and the reaction MUST be exothermic (Δs surr > 0).  Exothermic reactions are favored, NOT endothermic reactions.  Exothermic (ΔH 0)  Endothermic (ΔH > 0, ΔS < 0)  ΔS T = Δs sys + Δs surr  If ΔS T > 0, then the chemical reaction is spontaneous

16. Example 1: Will entropy increase or decrease for the following? a)N 2 (g) + 3H 2 (g)  2NH 3 (g) b)2KClO 3 (s)  2KCl (s) + 3O 2 (g) c)CO (g) + H 2 O (g)   CO 2 (g) + H 2 (g) d)C 12 H 22 O 11 (s)  C 12 H 22 O 11

17. How do we calculate the entropy change (ΔS) in a chemical reaction?  Same method as using the enthalpies of formation to calculate ΔH and use the same table.  aA + bB  cC + dD ΔS° =[c (ΔS° C ) + d(ΔS° D )] - [a (ΔS° A ) + b (ΔS° B )]

18. Example 2: Calculate ΔS° for the following reaction at 25°C…. 4HCl (g) + O 2 (g)  2Cl 2 (g) + 2H 2 O (g)

19. Homework  pp. 382-383 #69, 71-73  pp. 742-743 #19, 27