 Download Download Presentation Unit 3: Thermochemistry

# Unit 3: Thermochemistry

Download Presentation ## Unit 3: Thermochemistry

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
##### Presentation Transcript

1. Unit 3: Thermochemistry Chemistry 3202

2. Unit Outline • Temperature and Kinetic Energy • Heat/Enthalpy Calculation • Temperature changes (q = mc∆T) • Phase changes (q = n∆H) • Heating and Cooling Curves • Calorimetry (q = C∆T & above formulas)

3. Unit Outline • Chemical Reactions • PE Diagrams • Thermochemical Equations • Hess’s Law • Bond Energy • STSE: What Fuels You?

4. Temperature and Kinetic Energy Thermochemistryis the study of energy changes in chemical and physical changes eg. dissolving a solid burning phase changes

5. Temperature - a measure of the average kinetic energy of particles in a substance - a change in temperature means particles are moving at different speeds - measured in either Celsius degrees or degrees Kelvin Kelvin = Celsius + 273.15

6. p. 628

7. 300 K 500 K # of particles Kinetic Energy

8. The Celsius scale is based on the freezing and boiling point of water The Kelvin scale is based on absolute zero- the temperature at which particles in a substance have zero kinetic energy.

9. Heat/Enthalpy Calculations system - the part of the universe being studied and observed surroundings - everything else in the universe open system - a system that can exchange matter and energy with the surroundings eg. an open beaker of water a candle burning closed system - allows energy transfer but is closed to the flow of matter.

10. isolated system – a system completely closed to the flow of matter and energy heat - refers to the transfer of kinetic energy from a system of higher temperature to a system of lower temperature. - the symbol for heat is q WorkSheet: Thermochemistry #1

11. Part A: Thought Lab (p. 631)

12. Part B: Thought Lab

13. Heat/Enthalpy Calculations specific heat capacity - the amount of energy , in Joules (J), needed to change the temperature of one gram (g) of a substance by one degree Celsius (°C). • The symbol for specific heat capacity is a lowercase c

14. A substance with a large value of c can absorb or release more energy than a substance with a small value of c. ie. For two substances, the substance with the larger c will undergo a smaller temperature change with the same amount of heat applied

15. q = mc∆T q = heat (J) m = mass (g) c = specific heat capacity ∆T = temperature change = T2 – T1 = Tf – Ti FORMULA

16. eg. How much heat is needed to raise the temperature of 500.0 g of water from 20.0 °C to 45.0 °C? Solve q = m c ∆T for c, m, ∆T, T2 & T1 • p. 634 #’s 1 – 4 • p. 636 #’s 5 – 8 WorkSheet: Thermochemistry #2

17. heat capacity- the quantity of energy , in Joules (J), needed to change the temperature of a substance by one degree Celsius (°C) • The symbol for heat capacity is uppercase C • The unit is J/ °C or kJ/ °C

18. C = mc q = C ∆T C = heat capacity c = specific heat capacity m = mass ∆T = T2 – T1 FORMULA Your Turn p.637 #’s 11-14 WorkSheet: Thermochemistry #3

19. Enthalpy Changes • The difference between the potential energy of the reactants and the products during a physical or chemical change is the Enthalpy change or ∆H. • AKA: Heat of Reaction

20. Products ∆H Reactants PE Endothermic Reaction Reaction Progress

21. Products ∆H Reactants PE Endothermic Reaction Enthalpy ∆H Reaction Progress

22. Products Reactants Enthalpy ∆H is + Endothermic

23. products ∆H is - reactants Enthalpy Exothermic

24. Enthalpy Changes in Reactions • All chemical reactions require bond breaking in reactants followed by bond making to form products • Bond breaking requires energy (endothermic) while bond formation releases energy (exothermic) see p. 639

25. Enthalpy Changes in Reactions endothermic reaction - the energy required to break bonds is greater than the energy released when bonds form. exothermic reaction - the energy required to break bonds is less than the energy released when bonds form.

26. Enthalpy Changes in Reactions ∆H can represent the enthalpy change for a number of processes • Chemical reactions ∆Hrxn – enthalpy of reaction ∆Hcomb – enthalpy of combustion (see p. 643)

27. Formation of compounds from elements ∆Hof– standard enthalpy of formation The standard molar enthalpy of formation is the energy released or absorbed when one mole of a compound is formed directly from the elements in their standard states. ( see p. 642)

28. Phase Changes (p.647) ∆Hvap – enthalpy of vaporization ∆Hfus – enthalpy of melting ∆Hcond – enthalpy of condensation ∆Hfre – enthalpy of freezing • Solution Formation ∆Hsoln – enthalpy of solution

29. There are three ways to represent any enthalpy change: 1. thermochemical equation - the energy term written into the equation. 2. enthalpy term is written as a separate expression beside the equation. 3. enthalpy diagram.

30. eg. the formation of water from the elements produces 285.8 kJ of energy. 1. H2(g) + ½ O2(g) → H2O(l) + 285.8 kJ 2. H2(g) + ½ O2(g) → H2O(l) ∆Hf = -285.8 kJ/mol thermochemical equation

31. H2(g) + ½ O2(g) ∆Hf = -285.8 kJ/mol H2O(l) enthalpy diagram 3. Enthalpy (H) examples: pp. 641-643 questions p. 643 #’s 15-18 WorkSheet: Thermochemistry #4

32. FORMULA: q = n∆H q = heat (kJ) n = # of moles ∆H = molar enthalpy (kJ/mol) Calculating Enthalpy Changes

33. eg. How much heat is released when 50.0 g of CH4forms from C and H ? p. 642 q = nΔH = (3.115 mol)(-74.6 kJ/mol) = -232 kJ

34. eg. How much heat is released when 50.00 g of CH4 undergoes complete combustion? q = nΔH = (3.115 mol)(-965.1 kJ/mol) = -3006 kJ

35. eg. How much energy is needed to change 20.0 g of H2O(l) at 100 °C to steam at 100 °C ? Mwater = 18.02 g/mol ΔHvap = +40.7 kJ/mol q = nΔH = (1.110 mol)(+40.7 kJ/mol) = +45.2 kJ

36. eg. The molar enthalpy of solution for ammonium nitrate is +25.7 kJ/mol. How much energy is absorbed when 40.0 g of ammonium nitrate dissolves? q = nΔH = (0.4996 mol)(+25.7 kJ/mol) = +12.8 kJ

37. What mass of ethane, C2H6, must be burned to produce 405 kJ of heat? ΔH = -1250.9 kJ q = - 405 kJ q = nΔH n = 0.3238 mol m = n x M = (0.3238 mol)(30.08 g/mol) = 9.74 g

38. Complete: p. 645; #’s 19 – 23 pp. 648 – 649; #’s 24 – 29 19. (a) -8.468 kJ (b) -7.165 kJ 20. -1.37 x103 kJ 21. (a) -2.896 x 103 kJ 21. (b) -6.81 x104 kJ 21. (c) -1.186 x 106 kJ 22. -0.230 kJ 23. 3.14 x103 g

39. 24. 2.74 kJ 25.(a) 33.4 kJ (b) 33.4 kJ 26.(a) absorbed (b) 0.096 kJ 27.(a) NaCl(s) + 3.9 kJ/mol → NaCl(aq) (b) 1.69 kJ (c) cool; heat absorbed from water 28. 819.2 g 29. 3.10 x 104 kJ

40. p. 638 #’ 4 – 8 pp. 649, 650 #’s 3 – 8 p. 657, 658 #’s 9 - 18 WorkSheet: Thermochemistry #5

41. Heating and Cooling Curves Demo: Cooling of p-dichlorobenzene

42. KE PE KE Cooling curve for p-dichlorobenzene 80 Temp. (°C ) liquid 50 freezing solid 20 Time

43. KE KE PE Heating curve for p-dichlorobenzene 80 Temp. (°C ) 50 20 Time

44. What did we learn from this demo?? • During a phase change temperature remains constant and PE changes • Changes in temperature during heating or cooling means the KE of particles is changing

45. p. 651