1. Thermochemistry Energy and Chemical Change

2. Energy Energy can change for and flow, but it is always conserved.

3. The Nature of Energy Energy – the ability to do work or produce heat – Potential energy – Kinetic energy

4. Chemical systems contain both kinetic and potential energy – Kinetic energy – Potential energy

5. Chemical potential energy Energy that is stored in a substance because of its composition

6. Heat Heat = energy Symbol = q Always flows from warmer object to cooler object

7. Measuring Heat calorie – the amount of energy needed to raise the temperature of one gram of pure water by one degree Celsius – Calorie vs. calorie

8. Joule – the SI unit of energy 1 calorie = 4.184 J A chemical reaction releases 213 J of energy. How many calories has this reaction released?

9. A big mac contains 550 Calories, express this energy in Joules.

10. Specific Heat Amount of heat needed to raise the temperature of 1 g of a substance by 1 o C – Symbol = c – Unit = J/g o C High specific heat = absorbs a lot of energy w/small changes in temperature – Water = 4.184 J/g o C – Concrete = 0.84 J/g o C

11. Calculating changes in heat q = (m)(c)( T)

12. If the temperature of 34.4 g of ethanol increases from 25.0 o C to 78.8 o C, how much heat has been absorbed by the ethanol? (specific heat of ethanol = 2.44 J/g o C)

13. The temperature of a 10.0 g piece of iron changed from 50.4 oC to 25.0 oC, how much energy was release by the iron? (specific heat of iron is 0.449 J/g o C)

14. Heat Calorimeter – insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process.

15. Chemical Energy and the Universe Thermochemistry – the study of heat changes that accompany chemical reactions and phase changes – System – specific part of the universe that contains the reaction or process you want to study – Surroundings – everything else in the universe other than the system

16. Law of conservation of energy 1 st law of thermodynamics – energy cannot be created or destroyed, only transferred

17. Energy being absorbed/released is indicated from the system’s point of view -q = energy released by system(exothermic) + q = energy absorbed by system(endothermic) q system = -q surroundings (m 1 )(c 1 )( Δ T 1 ) = - (m 2 )(c 2 )(Δ T 2 )

18. A 125 g sample of iron at 93.5 o C is dropped into an unknown mass of water at 25.0 o C. The final temperature of the mixture is 32.0 o C. The C of iron is 0.451 J/g o C, the C of water is 4.18 J/g o C. What is the mass of the water?

19. A 118 g piece of tin at 85 o C is dropped into 100 g of water at 35 o C. The final temperature of the mixture is 38 o C. C of water is 4.18 J/g o C. What amount of heat is absorbed by the water? What amount is released by the tin? What is the specific heat of tin?

20. Enthalpy and enthalpy change Enthalpy (H) – heat content of a system at a constant pressure Cannot measure enthalpy directly but can measure change in enthalpy (heat absorbed or released in a chemical reaction) Enthalpy of reaction ( H rxn )– change in enthalpy for a reaction

21. ΔH rxn = H products – H reactants Exothermic reaction Endothermic reaction

22. Thermochemical Equations Thermochemical equations express the amount of heat released or absorbed by chemical reactions

23. Writing Thermochemical Equations Thermochemical equation – balanced chemical equation that includes the states of all reactants and products and the energy change CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l) ΔH comb = -891 kJ

24. Enthalpy (heat) of combustion (ΔH comb ) – the enthalpy change for the complete burning of one mole of the substance Enthalpy (heat) of vaporization (ΔH vap ) – heat required to vaporize one mole of a liquid Enthalpy (heat) of fusion (ΔH fus ) – heat required to melt one mole of a solid

25. How much heat is released when 54.0 g of glucose is burned? ΔH comb = -2808 kJ/mol

26. H 2 O(l)  H 2 O(g) ΔH vap = 40.7 kJ H 2 O(s)  H 2 O(l) ΔH fus = 6.01 kJ

27. Phase Change Diagram C steam = 1.70 J/g o C Liquid  steam: 40.7 kJ/mol C ice = 2.10 J/g o C C Water = 4.18 J/g o C solid  liquid: 6.01 kJ/mol

28. How much heat must be absorbed to melt 150.0 g of water?

29. How much heat is released when 50.0 g of steam cools to 40 o C?

30. Challenge!!! A 39.0g sample of ice at -125 o C changes into steam at 125 o C. How much energy is absorbed during this process?

31. Calculating Enthalpy Change Hess’s Law – if you can add two or more thermochemical equations to produce a final equation for a reaction, then the sum of the enthalpy changes for the individual reactions is the enthalpy change for the final reaction

32. What is the energy change for the following reaction: 2S(s) + 3O 2 (g)  2SO 3 (g) a.S(s) + O 2 (g)  SO 2 (g) ΔH = -594 kJ b.2SO 3 (g)  2SO 2 (g) + O 2 (g) ΔH = 198 kJ

33. What is the energy change for the following? H 2 O 2 (l)  2H 2 O(l) + O 2 (g) a. 2H 2 (g) + O 2 (g)  2H 2 O(l) ΔH = -572kJ b. H 2 (g) + O 2 (g)  H 2 O 2 (l) ΔH = -188kJ

34. Standard heat of formation The change in enthalpy that accompanies the formation of one mole of the compound in its standard state ΔH f ΔH f of an element = 0

35. ΔH rxn = sum of ΔH f products – sum of ΔH f reactants

36. What is ΔH rxn for the following equation: CH 4 (g) +2O 2 (g)  CO 2 (g) + 2H 2 O(l)

37. What is the ΔH rxn for the following: 4NH 3 (g) + 7O 2 (g)  4NO 2 (g) + 6H 2 O(g)

38. Reaction Spontaneity Changes in enthalpy and entropy determine whether a process is spontaneous

39. Spontaneous processes Any physical or chemical change that once begun, occurs with no outside intervention – Iron rusting – Paper burning Often some energy from the surroundings must be supplied to get the process started

40. Entropy Entropy(S) – a measure of the number of possible was that the energy of a system can be distributed – Determined by the freedom of the systems particles to move and the number of ways they can be arranged – Disorder or randomness of a system

41. Second law of thermodynamics – spontaneous processes always proceed in such a way that the entropy of the universe increases

42. Predicting changes in entropy +ΔS = entropy increases - ΔS = entropy decreases

43. Changes resulting in +ΔS Changes in state that allow more molecule movement – (s)  (l) – (l)  (g) Number of particles increases in a reaction CaCO 3  CaO + CO 2 Solid or liquid dissolves in solvent Increase in temperature

44. Changes resulting in -ΔS Phase changes that decrease molecule movement – (g)  (l) – (l)  (s) Number of particles decreases in a reaction Dissolving of gas in a solvent Decrease in temperature

45. Predict the sign of ΔS for the following: – ClF(g) + F 2 (g)  ClF 3 (g) – NH 3 (g)  NH 3 (aq) – CH 3 OH(l)  CH 3 OH(aq) – C 10 H 8 (l)  C 10 H 8 (s)

46. Gibbs Free Energy ΔG = ΔH – TΔS -ΔG = spontaneous reaction + ΔG = nonspontaneous reaction

47. For a process, ΔH = 145 kJ and ΔS = 322 J/K. is the process spontaneous at 382K?

48. ΔHΔHΔSΔSΔGΔG Reaction Spontaneity NegativePositiveNegativeAlways spontaneous Negative Negative or positive Spontaneous at low temperatures Positive Negative or positive Spontaneous at high temperatures PositiveNegativePositiveNever spontaneous