1. Energy Changes and Rates of Reaction
Unit 3

2. Thermochemistry...
...is the study of the energy changes that accompany physical or chemical changes in matter.
Changes that occur in matter can be classified as:
Physical (hydrogen boils at -252℃); H2 (l) + heat → H2 (g)
Chemical (hydrogen burns as fuel ); 2H2 (g) + O2 (g) → 2H2O (l) + heat
Nuclear (hydrogen undergoes nuclear fusion in the Sun, producing helium).
These are all accompanied by a change in energy.

3. Thermal Energy (heat) – energy available from a substance as a result of the motion of its molecules
Temperature (T) – the average kinetic energy of the molecules in a sample, measured in °C or K
Heat (Q) – amount of energy transferred between substances, measured in Joules (J)
*Energy flows between substances because of their difference in temperature.

5. Exothermic Reactions release thermal energy
heat (Q) flows from the system to the surroundings, usually causing an increase in the temperature of the surroundings.
Q has a negative value (Q < 0); i.e. losing heat from the system.

6. Endothermic Reactions
absorb thermal energy
heat (Q) flows from the surroundings into the system, usually causing a decrease in the temperature of the surroundings.
Q has a positive value (Q > 0); i.e. adding heat to the system.

7. The First Law of Thermodynamics
The total energy of the universe remains constant (i.e. energy cannot be created or destroyed, only converted from one form to another).
Euniverse = constant OR ∆Euniverse = 0
Euniverse = Esystem + Esurroundings OR ∆Euniverse = ∆Esystem + ∆ Esurroundings = 0
∆E system = – ∆E surroundings
*For any system that gains energy, the energy must come from the surroundings; for any system that loses energy, that energy must enter the surroundings.

8. Calorimetry... Specific Heat Capacity...
Measuring Energy Absorbed or Released
Calorimetry...
… is an experimental technique used to measure energy changes in chemical systems.
*different substances vary in their ability to absorb amounts of heat.
Specific Heat Capacity...
… is the amount of energy required to raise the temperature of one gram of a substance by 1°C or one K; depends on type of substance and state of substance.
(Table 5.1, p. 280 or p.743)- values recorded at SATP (25C and 100kPa)

9. Q= mcΔT The amount of heat (Q) transferred depends on:
Heat (J)
Mass (g)
Specific Heat Capacity (J/g•°C)
Change in temperature (°C)
The amount of heat (Q) transferred depends on:
mass of sample (m) measured in grams
temperature change (ΔT) measured in °C or K
specific heat capacity (c) measured in J/g•°C or J/g•K

10. Example 1
Many water heaters use the combustion of natural gas (assume methane) to heat the water in the tank. When L of water at 10.0°C is heated to 65.0°C, how much heat flows into the water?
m = 150L x 1kg/L = 150 kg = g c = 4.19 J/g°C ΔT = (65.0 – 10.0) = 55.0°C
Q = mcΔT = ( ) x (4.19) x (55.0) = J or x 107 J or x 104 kJ

11. Example 2
If 25.0 g of aluminum cools from 310°C to 37°C, how many joules of heat energy are lost by the sample?
m = 25.0 g c = J/g°C ΔT = 37 – 310 = –273℃
Q =mcΔT = (25.0) x (0.897) x (–273℃) = – J = –6.12 kJ
* Negative sign indicates loss of energy

12. Example 3 Calculate the molar heat capacity (C) of water (J/mol•°C)
…. vs. specific heat capacity (J/ g•℃)
*RECALL:m = n x M…..
C = c x M
C = (4.19 J / g•℃) (18.02 g /mol)
= 75.5 J / mol•℃

13. The total energy of a system is equal to the sum of the potential and kinetic energy of all species in the system.
Kinetic Energy
moving electrons in atoms
vibration, rotation and translation of atoms and molecules
Potential Energy
nuclear potential energy of protons and neutrons
bond energy (stored energy)
intra and intermolecular forces
It is not possible to measure all of these energies for a system!

14. Instead, we study the energy absorbed or released to the surroundings during a change in the system – the change in enthalpy, “ΔH” (i.e. difference in energy between reactants and products).
This example of exothermic change shows that the change in potential energy of the system (ΔH) equals the change in kinetic energy of the surroundings (Q).

15. Enthalpy “H” The heat content of the system.
H= E + PV (total energy of the system + pressure • volume)
It is not possible to measure the total enthalpy of a system, but it is possible to measure the change in the enthalpy of a system.
∆H= ∆E + ∆(PV ); where ∆(PV ) is work done by the reaction on the surroundings.

16. Measuring Enthalpy Changes
In an open beaker in a laboratory, we assume that there is no change in pressure or volume. Therefore, ∆(PV ) = 0.
∆H = ∆E = Q; the heat exchanged between the system and the surroundings is equal to the enthalpy change of the system.

17. The Second Law of Thermodynamics
When 2 objects are in thermal contact, heat is always transferred from the object at a higher temperature to the object at a lower temperature, until the 2 objects are at the same temperature (= “thermal equilibrium”).
*Textbook- p.283

18. Professor Dave Explains….Thermochem!

19. Homework
Read 5.1- p
Do p. 281 #2, 4-6; p. 291 #1, 5, 6, 10