Thermochemistry

1. Thermochemistry 16.1

2. Thermochemistry • study of the transfer of energy as heat • Virtually every chemical reaction is accompanied by a change in energy. • Chemical reactions usually absorb or release energy as heat. • Thermochemistry: is the study of the transfers of energy as heat that accompany chemical reactions and physical changes.

3. Thermochemistry • Temperature: average kinetic energy of the particles in a sample of matter • K=273.15 + °C • 1 Joule= 1N·m=1kg·m2/s2 • Joule:is the SI unit of heat as well as all other forms of energy.

4. Heat • energy transferred between samples of matter because of a difference in their temperatures • Heat can be thought of as the energy transferred between samples of matter because of a difference in their temperatures. • Energy transferred as heat always moves spontaneously from matter at a higher temperature to matter at a lower temperature.

5. The direction of energy transfer is determined by the temperature differences between the objects within a system. The energy is transferred as heat from the hotter brass bar to the cooler water. • This energy transfer will continue until the bar and the water reach the same temperature.

6. Specific Heat • Specific heat: amount of energy required to raise the temperature of one gram of a substance by 1 °C • cp=specific heat • q=energy lost or gained (in J) • m=mass of sample (in g) • ΔT=change in temperature (in °C or K) • Energy transferred as heat always moves spontaneously from matter at a higher temperature to matter at a lower temperature. • Heat can be thought of as the energy transferred between samples of matter because of a difference in their temperatures. q=mcDT

7. Page 533 You need ALL these for HW • Other cp • Steel = 0.452 • P = 0.80 • Concrete = 0.75 • Wood = 2.00 • Ethanol = 2.5 Write these Additional 5 down in your notes. You need them for HW, too!

8. Enthalpy ΔH = H(products)-H (reactants) • An enthalpy change is the amount of energy absorbed or lost by a system as heat during a process at constant pressure. • The energy absorbed or released as heat during a chemical reaction at constant pressure is represented by ∆H. • The H is the symbol for a quantity called enthalpy.

9. ΔH = H(products)-H (reactants) 2H2 (g) + O2 (g) 2H2O (g) DH=-483.6 kJ/mol 2H2O (g)  2H2 (g) + O2 (g)DH= +483.6 kJ/mol exothermic endothermic

10. If R>P then the reaction is exothermic • If R<P then the reaction is endothermic EA curves and Reaction Types

11. Classwork • Specific Heat worksheet • Problems # 1, 4, 6, 7, 10 • Finish Worksheet Homework

12. Specific Heat WS Answers • 3,100 J • 15,000 J • 4,520 J • 6,270,000 J • 2.5, ethanol • 0.8, P • 0.25kg • 14kg • 22C or 22K • -1C or -1K • steel

13. Activation Energy 17.1

14. Activation energy • EA= activation energy • The energy required to begin a reaction

15. If R>P then the reaction is exothermic • If R<P then the reaction is endothermic EA curves and Reaction Types

16. Catalyst • Lowers the EA of a reaction, while not being consumed • Will remain After the reaction

17. Classwork Homework • Activation Energy Worksheets 1 and 2

18. Reaction Rate 17.2

19. Collision Theory • Reaction rate is a function of the collisions between molecules • For a chemical reaction to occur, molecules must collide with the correct • Orientation • Energy

20. Rate-Influencing Factors

21. Nature of Reactants • Inert or stable atoms or molecules are less likely to react simply because of their nature

22. Surface Area • Increased surface area allows for a greater reaction rate since more material is able to come into contact with other reactants

23. Temperature • Increased temperature increases the likelihood of collisions • It also increases the energy of the collisions

24. Concentration • Increased concentration increases the likelihood of collisions

25. Concentration

26. Catalysts • Catalysts lower the energy required to get a reaction going • They are not used up in the process of the reaction

27. Collision Theory Worksheet • Reaction Rate Worksheet Classwork/Homework

28. Chemical Equilibrium Chapter 18

29. Reversible reactions • Chemical reactions in which products can reform into reactants • Endothermic in one direction, exothermic in the other direction • 2HgO2Hg + O2 • A + BC + D

30. Equilibrium • Occurs when the forward reaction and the reverse reaction occur at the same rate • Le Châtelier’s Principle: equilibrium is shifted in the direction that tends to release the stress on the system

31. Shift in equilibrium

32. Change in Pressure 3H2+ 1 N22 NH3 + 92kJ • (a) H2, N2, and NH3 are in equilibrium within a closed system. • (b) Addition of more N2 causes a stress on the initial equilibrium. • (c) The new equilibrium position for this system.

33. Change in Concentration • A + B  C + D • By adding A or B, equilibrium will shift to increasing C and D • By adding C or D, equilibrium will shift to increasing A and B

34. Change in Temperature Recall that: • Endothermic (-) • – energy is absorbed • - Temperature decreases • Exothermic (+) • - energy is released • - temperature increases Energy is a Reactant!! Energy is a Product!!

35. Change in Temperature • Different temperatures can cause an equilibrium system to shift and seek a new equilibrium position.

36. Reactions that go to completion • Form a Gas 2. Form a Precipitate 3. Form a Slightly Ionized Product

37. Form a gas • Often move toward completion because the gas is naturally removed as it diffuses away unless it is trapped • H2CO3  H2O + CO2

38. Form a precipitate • Precipitate is insoluble (or it wouldn’t form) and is therefore removed from system shifting equilibrium

39. Slightly ionized • H+ (or H3O+)and OH- are usually removed from system, shifting equilibrium

40. Equilibrium Constant • Keq = Products over reactants, when coefficients become exponents

41. Equilibrium Constant A  2B + C Write the equilibrium constant. Keq = [B]2[C] [A]

42. Classwork/Homework • Le Châtelier’s Principle Worksheet