1. It was a hot summer day. Mattie poured herself a glass of lemonade
It was a hot summer day. Mattie poured herself a glass of lemonade. The lemonade was warm, so she put some ice in the glass. After 10 minutes she noticed the ice was melting and the lemonade was cold. Mattie wondered what made the lemonade get cold.
The coldness from the ice moved into the lemonade
The heat from the lemonade moved into the ice.
The coldness and the heat moved back and forth until the lemonade cooled off.

2. Thermochemistry
Chapter 5

3. First Law of Thermodynamics
states that energy is conserved
Energy that is lost by a system must be gained by the surroundings or vice versa.

4. Enthalpy (H)
accounts for heat flow in chemical changes occurring at constant pressure

5. Enthalpy (H)
State function:
H=Hfinal-Hinitial

6. Enthalpy (H)
When DH > 0, the system has gained heat from surroundings (endothermic)
When DH < 0, the system has released heat to the surroundings (exothermic)

7. Enthalpy (H)

8. Enthalpies of Reaction
Thermochemical equations are balanced chemical equations that show the associated enthalpy change

9. Enthalpies of Reaction
DH = H(products) – H(reactants)

10. Guidelines
(1)H is an extensive property meaning it depends on the amount of reactant consumed in the process

11. Guidelines
(2) The DH for a reaction is equal in magnitude, but opposite in sign, to the reverse reaction

12. Guidelines
(2)

13. Guidelines
(3) The DH for a reaction depends on the state of the reactants and products and we assume they are at constant T

14. Practice 5.5
How much heat is released when 4.50 g of methane gas is burned in a constant pressure system?
Given:
CH4(g) + 2O2(g) CO2(g) +2H2O(l) DH = -890kJ

15. Enthalpies of Reaction
DH can be determined experimentally by measuring the heat flow of a reaction at constant pressure which is done by measuring Temperature changes

16. Enthalpies of Reaction
Calorimetry is the measurement of heat flow
A calorimeter will measure this heat flow

17. Heat Capacity
the amount of heat required to raise the temperature of an object by 1K or 1oC.

18. heat capacity of one mole of a substance
Molar Heat Capacity
heat capacity of one mole of a substance

19. Specific Heat Capacity
heat capacity of 1g of substance

20. Specific Heat Capacity

21. Enthalpies of Reaction
Heat = specific heat x grams of substance x DT
q=cm DT

22. Sample Exercise 5.5 Relating Heat, Temperature Change, and Heat Capacity
(a) How much heat is needed to warm 250 g of water (about 1 cup) from 22 °C (about room temperature) to near its boiling point, 98 °C? The specific heat of water is 4.18 J/g-K. (b) What is the molar heat capacity of water?

23. Constant- Pressure Calorimetry

24. Constant- Pressure Calorimetry
(1)Assume the calorimeter prevents loss or gain of heat from the solution to its surroundings

25. Constant- Pressure Calorimetry
(2a) For an exothermic rxn, heat is lost by rxn and gained by soln so the T of soln rises

26. Constant- Pressure Calorimetry
(2b) For an endothermic rxn, heat is gained by rxn and lost by soln so the T of soln goes down

27. Constant- Pressure Calorimetry
(3) qsoln = specific heat of solution x gram of solution x DT = -qrxn

28. Sample Exercise 5.6 Measuring ΔH Using a Coffee-Cup Calorimeter
When a student mixes 50. mL of 1.0 M HCl and 50 mL of 1.0 M NaOH in a coffee-cup calorimeter, the temperature of the resultant solution increases from 21.0 °C to 27.5 °C. Calculate the enthalpy change for the reaction in kJ/mol HCl, assuming that the calorimeter loses only a negligible quantity of heat, that the total volume of the solution is 100. mL, that its density is 1.0 g/mL, and that its specific heat is 4.18 J/g-K.

29. Bomb Calorimetry
Constant-Volume Calorimetry

30. Bomb Calorimetry
The heat released in a combustion rxn is absorbed by the calorimeter contents, causing a rise in the temperature of the water.

31. Bomb Calorimetry
qrxn = -Ccal x DT
Ccal is the heat capacity of the bomb calorimeter

32. Sample Exercise 5.7 Measuring qrxn Using a Bomb Calorimeter
Methylhydrazine is used as a liquid rocket fuel. The combustion of methylhydrazine with oxygen produces N2(g), CO2(g), and H2O(l):
2 CH6N2(l) + 5 O2(g) → 2 N2(g) + 2 CO2(g) + 6 H2O(l)
When 4.00 g of methylhydrazine is combusted in a bomb calorimeter, the temperature of the calorimeter increases from °C to °C. In a separate experiment the heat capacity of the calorimeter is measured to be kJ/°C. Calculate the heat of reaction for the combustion of a mole of CH6N2.

33. Hess’ Law
If a reaction is carried out in a series of steps, DHrxn will equal the sum of DH for the individual steps.
DHrxn = DH1 + DH2 + …

34. Hess’ Law
Hess’ law provides ways to calculate energy changes that are difficult to measure directly.

35. Sample Exercise 5.8 Using Hess’s Law to Calculate ΔH

36. Hess’ Law

37. Hess’ Law (5.9)

38. Hess’ Law (5.9)
Practice Exercise

39. Standard enthalpy of a reaction
is the enthalpy change when all reactants and products are in their standard states, DHorxn

40. Standard States
Standard state is when a substance is in its pure form at 1 atm and 298K.

41. Standard enthalpy of formation
DHof,, reported in kj/mol, is the change in enthalpy for the reaction that forms 1 mole of the compound from its elements, with all substances in their standard state.

42. Standard enthalpy of a formation
DHfo of the most stable form of an element is zero because there is no formation reaction needed for an element in its standard state.

43. Standard enthalpy of a formation

44. Standard enthalpy of a formation (5.10)
Do these represent standard enthalpy of formation reactions?

45. Enthalpy of Reaction
DHorxn = S n DHof (products) - S m DHof (reactants) , where n and m are the coefficients in the balanced chemical equation

46. 5.11 Enthalpy of Reaction
Calculate the standard enthalpy change for the combustion of 1 mol of benzene(l), to form CO2(g) and H2O(l).

47. 5.12 Enthalpy of Reaction
The standard enthalpy change for the decomposition reaction of CaCO3 is kJ. From the values for the standard enthalpies of formation of CaO(s) and CO2(g), calculate the standard enthalpy of formation of CaCO3(s).