1. Chapter 17 Energy and Chemical Change

2. Thermochemistry The study of heat changes in chemical reactions

3. Law of Conservation of Energy Defn – energy can be converted from one to another, but neither created nor destroyed Ex: potential to kinetic solar to chemical

4. Energy and Chemical Change Energy (defn) – ability to do work or produce heat Exists in two forms a) potential energy – stored energy b) kinetic energy – energy in motion

5. Chemical Potential Energy Defn – energy stored in a substance based on composition Ex: methane (CH 4 ) vs. propane (C 3 H 8 ) C HH H H C HH H H C H H C H H Propane has more energy b/c it has more bonds

6. Heat (q) Defn – energy that goes from warmer object to colder object Energy transfer moves from an area of high energy to an area of low energy

7. Units of Heat calorie (cal) – amount of heat required to raise the temp of 1 gram of water one degree Celsius –Calorie vs calorie 1 Calorie = 1000 calories = 1 kilocalorie (kcal) Joule – SI unit of heat and energy 1 calorie = 4.18 J

8. Comparing Specific Heat water (liquid) aluminum iron water (ice) 4.18 2.03 0.897 0.449 ethanol 2.44

9. Specific Heat Defn – amount of heat required to raise temp of one gram of any substance by one degree Celsius q = m c ΔT heat massspecific heat change In temp Unit: J g ̊ C

10. Specific Heat If heat is released from the reaction: -q If heat is absorbed into the reaction: +q Change in temperature = t final - t initial

11. Sample problem (a) The temperature of a 10.0 g sample of iron is changed from 50.4°C to 25.0°C with the release of 114 J of heat. What is the specific heat of iron? q = mcΔT 114 J = (10.0 g) c (25.4°C) c = 0.449 J/g·°C

12. Sample problem (b) If the temperature of 34.4 g ethanol increases from 25°C to 78.8°C, how much heat is absorbed by ethanol? (specific heat of ethanol = 2.44 J/g°C) q = mcΔT = (34.4 g) ( 2.44 J/g°C) (53.8°C) = 4515 J

13. Heat In Chemical Reactions Calorimeter – insulated device used to measure amount of heat absorbed or released during a chemical or physical process Thermochemistry – study of heat changes in chemical reactions

14. 3 parts we look at 1) system – specific part of universe that contains the reaction 2)surroundings – everything in universe other than the system 3)universe – system + surroundings

15. Enthalpy (H) Defn – heat content of a system Enthalpy (heat) of rxn (ΔH rxn ) a) defn – change in enthalpy for a reaction b) formula ΔH rxn = H products – H reactants A + B  C + D

16. Enthalpy (H) c)endothermic vs exothermic rxns if +ΔH rxn = endothermic rxn if –ΔH rxn = exothermic rxn

17. Reaction Energy Diagrams This is an exothermic reaction – the reactants have more energy than the products, so energy has been released

18. Reaction Energy Diagrams This is an endothermic reaction – the reactants have less energy than the products, so energy has been absorbed

19. Reaction Energy Diagrams A: energy held by the activated complex B: energy of the reactants C: energy of the products F: heat of reaction ( ΔH) I: activation energy

20. HH products reactants products reactants EXOTHERMICENDOTHERMIC ΔHΔH ΔHΔH ΔH < 0 ΔH > 0

21. Example reactions 4 Fe + 3 O 2  2 Fe 2 O 3 + 1625 kJ i) exo- or endo-? ii) what is ΔH rxn ? Exothermic (heat written on right side of equation) ΔH rxn = -1625 kJ Another way to write equation: 4 Fe + 3 O 2  2 Fe 2 O 3 ΔH rxn = -1625 kJ

22. 2 Fe 2 O 3 4 Fe + 3 O 2 ΔH = -1625 kJ H

23. Example reactions 27 kJ + NH 4 NO 3  NH 4 + + NO 3 1- i) exo- or endo-? ii) what is ΔH rxn ? Endothermic (heat written on left side of equation) ΔH rxn = +27 kJ Another way to write equation: NH 4 NO 3  NH 4 + + NO 3 1- ΔH rxn = +27 kJ

24. H NH 4 + + NO 3 - NH 4 NO 3 ΔH = +27 kJ

25. Standard Enthalpy of Formation (ΔH f ° ) Defn – change in enthalpy when one mole of a compound is formed from its elements in their standard states Standard State – normal physical state of substance at room conditions (25°C and 1 atm) ex: standard state of Hg is liquid N 2 is gas

26. Standard Enthalpy of Formation (ΔH f ° ) Examples H 2 (g) + S (s)  H 2 S (g) ΔH f ° = -21 kJ S (s) + O 2 (g)  SO 2 (g) ΔH f ° = -297 kJ

27. Hess’s Law Defn – overall enthalpy change of reaction is equal to the sum of the enthalpy changes of individual steps A  D ΔH = ? A + B  CΔH = x C  D + BΔH = y Overall: A  D ΔH = x + y

28. Example problem #1 Calculate the enthalpy of reaction, ΔH, for the reaction: 2 H 2 O 2  2 H 2 O + O 2 (a) H 2 + O 2  H 2 O 2 ΔH = -188 kJ (b) 2 H 2 + O 2  2 H 2 OΔH = -572 kJ

29. Example problem #1 2 H 2 O 2  2 H 2 + 2 O 2 ΔH = +376 kJ 2 H 2 + O 2  2 H 2 OΔH = -572 kJ 2 H 2 O 2 + 2 H 2 + O 2  2 H 2 O + 2 H 2 + 2 O 2 2 H 2 O 2  2 H 2 O + O 2 ΔH = -196 kJ 1

30. Example problem #2 Calculate the enthalpy of reaction, ΔH, for the reaction: 2 S + 3 O 2  2 SO 3 (a) S + O 2  SO 2 ΔH = -297 kJ (b) 2 SO 3  2 SO 2 + O 2 ΔH = 198 kJ

31. Example problem #2 2 S + 2 O 2  2 SO 2 ΔH = -594 kJ 2 SO 2 + O 2  2 SO 3 ΔH = -198 kJ 2 S + 3 O 2 + 2 SO 2  2 SO 2 + 2 SO 3 2 S + 3 O 2  2 SO 3 ΔH = -792 kJ

32. Heat of Reaction (ΔH rxn ) Defn – amount of heat lost or gained in a reaction Formula ΔH rxn = Σ ΔH f (products) – Σ ΔH f (reactants)

33. Example problem Find ΔH rxn for the following reaction. Is reaction endo- or exothermic? H 2 S + 4 F 2  2 HF + SF 6 ΔH rxn = = -1745 kJ [ (2)(-273) + (1)(-1220) ] - [ (1)(-21) + (0)(4) ] exothermic -273 kJ-1220 kJ-21 kJ0 kJ

34. Entropy (S) Defn – measure of disorder or randomness in a system Formula ΔS = S products – S reactants Reaction Tendency –Nature favors a disordered state –The more entropy/disorder, the greater ΔS Unit: J/K

35. Entropy (S) Predicting ΔS - go to more disorder  + ΔS - go to less disorder  - ΔS Ex problem: predict ΔS system for these changes a)H 2 O (s)  H 2 O (l) + ΔS; Solid to liquid is more disorder

36. Entropy (S) b) 2 SO 3 (g)  2 SO 2 (g) + O 2 (g) Keep in mind: the reverse reactions have opposite signs There are more product particles (3) than reactant particles (2) + ΔS

37. Entropy example problem Calculate entropy change for this reaction: 2 PbO 2 (s)  2 PbO (s) + O 2 (g) ΔS rxn = = +209.2 J/K [(2)(68.7) + 205] - [(2)(66.6)] 68.7 205 66.6

38. Entropy, the Universe, and Free Energy The universe entropy ΔS universe > 0 Two natural universe processes a)things tend to go towards lower energy (-ΔH, exothermic) b)things tend to go towards higher disorder (+ΔS)

39. Spontaneous Process Defn – physical or chemical change that occurs with no outside intervention

40. Free Energy (G) Defn – energy available to do work Gibbs Free Energy Equation ΔG = ΔH – TΔS T is in Kelvin What does ΔG tell you? - ΔG  spontaneous rxn + ΔG  not spontaneous rxn

41. Free Energy Problem calculate the change in free energy, ΔG, for the reaction at 25°C. Is the reaction spontaneous or nonspontaneous? N 2 (g) + 3 H 2 (g)  2 NH 3 (g) ΔH rxn = -91800 J, ΔS rxn = -197 J/K

42. Free Energy Problem ΔG rxn = ΔH rxn – TΔS rxn = -91800 J – (298 K)(-197 J/K) = -33100 J Spontaneous

43. How does ΔH and ΔS affect spontaneity? +ΔS+ΔS -ΔS-ΔS -ΔH-ΔH+ΔH+ΔH Always spontaneous never spontaneous Spontaneity depends on temp Spontaneity depends on temp