1. CDO Chemistry 2015

2. Thermodynamics 1 st Law of Thermodynamics 1 st Law – energy cannot be created or destroyed it can just change forms Energy can be changed from Potential to Kinetic or vice versa

3. System Vs. Surroundings Two Parts to the Universe System – is the part of the universe being studied Surroundings – the rest of the universe that interacts with the system

4. Directionality of Heat Heat – (q) Flow from the surrounding to the system is positive; q > 0 Flow from the system into the surroundings is negative; q < 0

5. Endothermic Processes  Endothermic process (endo = in) is one that absorbs heat from the surrounding (it feels cold).

6. Exothermic Processes  Exothermic process (exo = out) is one that transfers heat to the surrounding (it feels hot)  All combustion and neutralization reactions

7. Heat and Temperature Temperature – is the measure of the average kinetic energy of the particles Heat energy does not necessarily change the temperature of 2 objects in the same way. Increase in temperature depends on Mass of the object The heat added Nature of the substance

8. Heat Capacity and Specific Heat The amount of energy required to raise the temperature of a quantity of a substance by 1 K (1  C) is its heat capacity. We define specific heat capacity (or simply specific heat - C) as the amount of energy required to raise the temperature of 1 g of a substance by 1 K. Specific Heat of water = 4.184 J/g o C

9. Finding Heat (q) q = m C  T q = quantity of heat in Joules m = mass in grams C = specific heat capacity = J/(g o C)  T = change in Temp

10. Example 1 How much heat is released when 10.0 g of Cu with a specific heat capacity of 0.385 J/g o C is cooled from 85.0 o C to 25.0 o C?

11. Example 2 If 500. J of heat is added to 100.0 g samples of water (specific heat = 4.184 J/g o C) and Silver (specific heat = 0.237 J/g o C) which substance will have the biggest change in Temperature?

12. Calorimetry We measure the transfer of heat (at a constant pressure) by a technique called calorimetry. In calorimetry... the heat released by the system is equal to the heat absorbed by its surroundings. the heat absorbed by the system is equal to the heat released by its surroundings. The total heat of the system and the surroundings remains constant.

13. Calorimetry We use an insulated device called a calorimeter to measure this heat transfer. A typical device is a “coffee cup calorimeter.” The calorimeter is often filled with water and the change in temperature of the water is measured

14. Calorimetry -q substance = +q water So……. -mCΔT = +mCΔT

15. Example If a 25.5 g sample of a metal at 95.0 o C is placed into a 100.0 g of water at 25.0 o C and the temperature of the water is raised to 32.6 o C, what is the specific heat of the metal?

16. Example An insulated cup contains 75.0g of water at 24.00 o C. A 26.00g sample of metal at 82.25 o C is added. The final temperature of the water and metal is 28.34 o C. What is the specific heat of the metal?

17. Enthalpy When a reaction is carried out a constant pressure, the heat (q) that is transferred in the reaction is given a special name “Enthalpy”, ΔH Enthalpy is measured in kilojoules per mole (KJ/mol)

18. ∆H Endothermic Reactions - Heat is transferred from the surrounding to the system, so the change in enthalpy is positive (ΔH > 0) Exothermic Reactions – Heat transferred from the system to the surrounding, so the change in enthalpy is negative (ΔH < 0)

19. Methods of Finding  H Measure it using a coffee cup calorimeter Calculate it using Hess’s Law Calculate it using average bond energies

20. Coffee Cup Calorimetry The ΔH for the reaction is equal in magnitude but opposite in sign to the q for the calorimeter q is heat lost or gained n is moles of substance

21. Important Assumptions Assume the volume of the water is equal to its mass in grams. If water is present use its mass in the q =mcΔT Always use the specific heat of water 4.184 J/g o C as the c value If two solutions are being reacted add volumes together and assume it is equal to mass in grams. Use the substance identified in the problem to determine n.

22. Example Calculate the enthalpy change of the combustion of ethanol (C 2 H 5 OH) from the following data. Assume all of the heat from the reaction is absorbed by the water.. Mass of water in calorimeter 200.00 g Temperature increase in water 13.00 o C Mass of ethanol burned0.45 g

23. Practice 1.56 grams of methanol (CH 3 OH) is burned in a calorimeter. This causes the 150.Og of water to increase its temperature from 20.0 o C to 38.6 o C. What is enthalpy of combustion for methanol?

24. Practice 100.0 mL of water was measured out and poured into a polystyrene cup with an initial temp of 18.3 o C. 5.20 g of NH 4 Cl was added to the cup and it dissolved. The minimum temp. was recorded to be 15.1 o C. Calculate the enthalpy of solution for NH 4 Cl per mole.

25. Practice 200.0 mL of water was measured out and poured into a polystyrene cup with an initial temp of 20.0 o C. 8.25 g of CaCl 2 was added to the cup and it dissolved. The maximum temp. was recorded to be 35.0 o C. Calculate the enthalpy of solution for CaCl 2.

26. Example When 1.00 L of 1.00M Ba(NO 3 ) 2 solution at 25.0 °C is mixed with 1.00 L of 1.00 M Na 2 SO 4 solution at 25 °C in a calorimeter, the white solid BaSO 4 forms and the temperature of the mixture increases to 28.1 °C. What is the enthalpy of reaction?

27. Practice When 1.00 L of 1.50 M HCl solution at 30.0 °C is mixed with 1.00 L of 1.50 M NaOH solution at 30.0 °C in a calorimeter, and the temperature of the mixture increases to 41.5 °C. What is the enthalpy of neutralization?

28. Hess’s Law If a chemical equation is the sum of multiples of other equations, the  H of this equation equals a similar sum of multiples of  H's for the other equations.

29. Enthalpy Relationships Given: H 2 (g) + 1/2 O 2 (g)  H 2 O(g) ∆H˚ = -242 kJ If multiplied: 2 H 2 (g) + O 2 (g)  2 H 2 O(g) ∆H˚ = -484 kJ

30. If Reversed the sign on the enthalpy value switches: H 2 O(g)  H 2 (g) + 1/2 O 2 (g) ∆H˚ = +242 kJ If a different Phase the enthalpy value changes: H 2 (g) + 1/2 O 2 (g)  H 2 O(liquid) ∆H˚ = -286 kJ

31. Example 8 CO 2 (g)  CO(g) + 1/2O 2 (g)  H = +283.0 KJ C(s) + O 2 (g)  CO 2 (g)  H = -393.5 KJ _________________________________________ C(s) + 1/2O 2 (g)  CO(g)  H = ?

32. Example 9 SO 2 (g)  S(s) + O 2 (g)  H = +296.8 KJ 2SO 3(g)  2S(s) + 3O 2 (g)  H = +791.4 KJ ________________________________ 2SO 2 (g) + O 2 (g)  2SO 3 (g)  H = ?

33. Average Bond Enthalpies This table lists the average bond enthalpies for many different types of bonds. Average bond enthalpies are positive, because bond breaking is an endothermic process. © 2009, Prentice-Hall, Inc.

34. Enthalpies of Reaction  H rxn =  (bond enthalpies of bonds broken) -  (bond enthalpies of bonds formed) Prior to using this formula the Lewis Structure for reactant and product must be drawn

35. Example Estimate the enthalpy change for the chlorination of ethylene: CH 2 CH 2 (g) + Cl 2 (g)  CH 2 ClCH 2 Cl

36. Example 2H 3 COH + 3O 2  2CO 2 + 4H 2 O