1. THERMOCHEMISTRY

2. ENERGY 1. Categories: Potential Stored energy, energy of position
Gravitational
Chemical
Nuclear
Kinetic
Energy of motion
Mechanical
Electrical
Heat
Sound
Radiant:
Electromagnetic radiation

3. ENERGY Law of Conservation of Energy
In any chemical or physical process, energy is neither created nor destroyed.
BUT energy CAN be transformed from one form to another.
Chemical → Light + Heat

4. HEAT vs TEMPERATURE Heat vs. Temperature – NOT the same a. Temperature
1. A measure of the average kinetic energy of
the particles (atoms/molecules)
2. Determines the DIRECTION of heat
flow
b. Heat - flows from the hotter object to
the colder object.

5. Heat Vs. Temperature Heat Temperature Heat can travel Temp Can NOT
Measured in Joules Measured in Degrees
Measured with calorimeter Measured with Thermometer
Heat is E it can do work Temp is a man made scale

6. HEAT
1. Factors affecting heat: Mass Specific heat Temperature

7. HEAT the amount of energy needed to raise the temperature of one
2. Measured by means of difference
Q = m * c * Δt
m = mass
c = specific heat
Δt = temperature change
UNITS of energy:
calorie (cal)
the amount of energy needed to raise the temperature of one
gram of water 1ºC.
kilocalorie (kcal) =
1000 cal or 1 Calorie
Joule (J)
1 cal = J
1 kcal = kJ
Kilowatt-hour kwh = 860 kcal

8. But wait….
1. food Calorie = 1,000 calories or 1 kilocalorie
Food Calories listed are actually kilocalories!
Note: That Big Mac (576 Calories) can actually raise _____________ grams of water 1°C.
576,000

9. HEAT 3. Specific heat (capacity)
Energy needed to raise 1 gram of a substance 1oC
Unit: cal or _J_
goC goC

10. HEAT FLOW ENDOTHERMIC Energy is absorbed ∆ H is +
a. Heat flows into the system
b. Ex. CHEMICAL CHANGES
1. PHOTOSYNTHESIS
6 CO2(g) H2O(g) → C6H12O6(s) O2(g)
2. ELECTROLYSIS
2 H2O(l) → 2 H2(g) + O2(g)
c. Ex. PHYSICAL CHANGES
1. BOILING:
H2O(l) → H2O(g)

11. HEAT FLOW
2. MELTING:
H2O(s) → H2O(l)

12. HEAT FLOW
EXOTHERMIC Energy is released ∆ H is – a. Heat flows out of the system b. Ex. CHEMICAL CHANGES 1. BURNING C(s) + O2(g) → CO2(g) 2. FORMATION OF A COMPOUND FROM ITS ELEMENTS 2 H2(g) + O2(g) → 2 H2O(l)

13. HEAT FLOW PHYSICAL CHANGES 1. FREEZING 2. CONDENSATION H2O(l) → H2O(s)
H2O(g) → H2O(l)
IF A GIVEN PROCESS IS EXOTHERMIC, THE REVERSE PROCESS IS
ENDOTHERMIC.
Most Synthesis reactions are EXOTHERMIC.
Most Decomposition reactions are ENDOTHERMIC.

14. H (ENTHALPY)
H (Enthalpy) – HEAT CONTENT of a system at constant pressure ∆ H (change in heat content) = H products – H reactants

15. Exothermic Heat Diagram

16. Endothermic Heat Diagram

17. HEATS OF REACTION Heat of formation: ∆Hf
The change in heat content when 1 mole of a compound is formed from its elements
a. ∆H is usually _____
Exothermic
b. Ex. The heat of formation of CaO is kcal/mol. Write the thermochemical equation.
1. Write the balanced equation.
2. Convert to show 1 mole product.
3. Show the physical states and give the value for ∆H.

18. HEATS OF REACTION
c. Compounds with very negative heats of formation – VERY STABLE ex kcal/mol d. Compounds with positive or low negative heats of formation – GENERALLY UNSTABLE ex kcal/mol or kcal/mol

19. HEATS OF REACTION HEAT OF COMBUSTION
c. Compounds with high positive heats of formation
VERY UNSTABLE
ex. HgC2N2O2 ∆H = + 64 kcal
(mercury fulminate)
HEAT OF COMBUSTION
Change in heat content when 1 mole of a substance burns
a. ∆H is ________
EXOTHERMIC
b. General format for a hydrocarbon combustion:
CxHy + O2 → CO2 + H2O

20. HEATS OF REACTION
c. Construct a heat content diagram for the combustion of CO(g). See table 4.2 for ∆H.

21. USES OF ∆H
1. If ∆H for the forward reaction is positive, then ∆ H for the reverse reaction is negative. 2. ∆H is directly proportional to the amount of reactants or products in a process. Ex. C(s) + O2(g) → CO2(g) ∆H = kcal a. Calculate the heat when 2.80 mol CO2 forms.

22. USES OF ∆H Ex. C(s) + O2(g) → CO2(g) ∆H = -94.1 kcal
b. Calculate the heat when 1.00 g of Carbon burns.

23. USES OF ∆H
Ex. C(s) + O2(g) → CO2(g) ∆H = kcal c. What mass of Carbon is needed to generate 10.0 kilocalories of heat?

24. CHANGES OF STATE HEAT OF FUSION ∆Hfus
Change in heat content when 1 mole of a substance changes from a solid to a liquid (melts)
a. The heat of fusion of water is 80.0 cal/g. Express as
kcal/mol.
b. Write a thermochemical equation for the melting of ice.

25. CHANGES OF STATE HEAT OF VAPORIZATION: ∆Hvap
Change in heat content when 1 mole of a substance changes from a liquid to a gas
a. The heat of vaporization is 540. cal/g. Express in kcal/mol.

26. CALORIMETRY CALORIMETRY – MEASURING ∆H
When using a calorimeter, Q = ∆H, and the equation becomes:
∆H = m c ∆T
m = mass in g or kg
c = specific heat (cal/goC or kcal/kgoC)
∆T = temperature change
cH2O = 1 cal/goC or 1 kcal/kgoC)
KEY: PAY ATTENTION TO THE UNITS!

27. CALORIMETRY
1. A reaction takes place in a calorimeter during which 40.0 g of water is heated from 24.0oC to 50.0oC. Find the heat of reaction (∆H).

28. CALORIMETRY
2. The temperature of a piece of copper with a mass of 95.4 g increases from 25.0oC to 48.0oC when the metal absorbs calories of heat. What is the specific heat of the copper?

29. HESS’ LAW
If Equation (C) is the sum of equations (A) and (B), then: ∆H for (C) = ∆H for (A) + ∆H for (B) Ex. 1: Given: C(s) + ½ O2(g) → CO(g) ∆H = kcal CO(g) + ½ O2(g) → CO2(g) ∆H = kcal Find: ∆H for: C(s) + O2(g) → CO2(g)

30. HESS’ LAW
Ex Calculate ∆H for: NO(g) + ½ O2(g)  NO2(g)

31. HESS’ LAW Ex. 3 Given: 1) Sn(s) + ½ O2(g)  SnO(s) ∆H = -68 kcal
2) SnO2(s)  SnO(s) + ½ O2(g) ∆H = 70 kcal
  Calculate the heat of formation of SnO2.

32. HESS’ LAW
Ex. 4: Find ∆H for: CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g) See Table 4.3 and H2O(l)  H2O(g) ∆H = 9.72 kcal

33. HESS’ LAW Ex. 5 Calculate ∆H for: 2 C(s) + H2(g)  C2H2(g)
Given: C2H2(g) + 5/2 O2(g)  2 CO2(g) + H2O(l) ∆H = kcal
C(s) + O2 (g)  CO2(g) ∆H = kcal
H2(g) + ½ O2(g)  H2O(l) ∆H = kcal 
Calculate ∆H for: 2 C(s) + H2(g)  C2H2(g)