1. Transfers of energy as heat in chemical reactions and physical changes
17-1 Thermochemistry
Transfers of energy as heat in chemical reactions and physical changes

2. Remember…
The First Law of Thermodynamics: energy cannot be created or destroyed only converted from one form to another

3. Heat vs.Temperature Heat (Joule, J)
measure of energy change in a system.
Temperature (Celsius, °C or Kelvin, K)
measure of the kinetic energy (movement) of the particles in a system.
Gaining or losing heat energy in a substance can change its temperature.
Exothermic
System loses energy to surroundings
Endothermic
System gains energy from surroundings

4. Specific Heat
is a property of matter, and different species have different Specific Heat.
The heat energy required to raise one gram of a pure substance 1° C or 1 K
The symbol we use is cp

5. Specific Heat Metals have very low cp, Water has a very high cp,
which is why metals often feel cold to the touch.
Table on page 513
Water has a very high cp,
4.184 J/g·0C
Substances with lower cp will rise in temperature faster and require less energy to do so than do substances with high cp.

6. Specific Heat 1 calorie= 4.184 Joules Substance J/g/oC or J/g/K
cal/g/oC or cal/g/K
Water (0 oC to 100 oC)
4.186
1.000
Zinc
.387
0.093
Ice (-10 oC to 0 oC)
2.093
0.500
Steam (100 oC)
2.009
0.480
Brass
.380
0.092
Wood (typical)
1.674
0.400
Soil (typical)
1.046
0.250
Air (50 oC)
Aluminum
.900
0.215
Tin
.227
0.205
Glass (typical)
.837
0.200
Iron/Steel
.452
0.108
Copper
0.0924
Silver
.236
0.0564
Mercury
.138
0.0330
Gold
.130
0.0310
Lead
.128
0.0305
1 calorie= Joules

7. Specific Heat Equation
cp = q/(m ΔT) OR q = cp x m x ΔT

8. Heat of Reaction
Quantity of energy released or absorbed during a chemical reaction
Thermochemical equation – shows the quantity of heat
Example: 2H2 + O2  2H2O kJ
Energy released or absorbed?

9. Heat of Reaction Example: 2H2 + O2  2H2O + 483.6 kJ
Energy released (on product side)
EXOTHERMIC
What about when it’s on the reactant side?

10. Heat of Reaction Example: 2H2O + 483.6 kJ  2H2 + O2
Energy absorbed (on reactant side)
ENDOTHERMIC

11. ΔH = Hproducts - Hreactants
Enthalpy Change (ΔH)
The amount of energy absorbed or lost
ΔH = Hproducts - Hreactants
Thermochemical equations usually written this way
2H2 + O2  2H2O ΔH = kJ/mol
When ΔH is negative, the system loses energy and it is EXOTHERMIC

12. ΔH = Hproducts - Hreactants
Enthalpy Change (ΔH)
The amount of energy absorbed or lost
ΔH = Hproducts - Hreactants
2H2O  2H2 + O2 ΔH = kJ/mol
When ΔH is positive, the system gains energy and it is ENDOTHERMIC

13. Endothermic or Exothermic?
C6H12O6 + 6O2 6CO2 + 6H20 ΔHrxn = kJ/mol

14. Heat of Formation (ΔHf)
energy released or absorbed to form 1 mole of a compound from its elements
p. 902
If a large amount of energy is released when compound is formed…
Endothermic or exothermic?
Positive or Negative ΔHf?

15. Heat of Formation (ΔHf)
energy released or absorbed to form 1 mole of a compound from its elements
p. 902
If a large amount of energy is released when compound is formed…
Endothermic or exothermic?
Positive or Negative ΔHf?
HIGH NEGATIVE ΔHf = VERY STABLE!

16. Heat of Formation (ΔHf)
POSITIVE ΔHf = UNSTABLE!
Pure elements ΔHf = O
ΔHf of carbon dioxide = kJ/mol
More stable than C and O alone
HgC2N2O2 ΔHf = kJ/mol
Very unstable, used in explosives

17. Heat of Combustion (ΔHc)
Reactants?
Products?
Exothermic or Endothermic?
Positive or Negative?

18. Heat of Combustion (ΔHc)
Reactants? C and H with O2
Products? CO2 and H2O and heat and light
Exothermic or Endothermic?
Positive or Negative?
P. 896

19. Stability of these compounds?
Al2O3 (s) kJ/mol
CaCO3 (s) kJ/mol
NO (g) kJ/mol
O kJ/mol