1. Energy in Chemical & Physical Changes

2. Thermochemistry
Study of changes that accompany chemical reactions and phase changes
The Universe is considered to be made of 2 parts:
1. System: part that contains the reaction or process
2. Surroundings: everything else

3. ENERGY defined as the ability to do work or transfer heat energy.
2 types of energy
Potential Energy (PE): Energy at rest due to the position of an object; chemical potential energy is the energy stored in a substance’s bonds.
2. Kinetic energy (KE): Energy of the motion of particles in a substance and is directly proportional to temperature. As temperature increases, KE also increases.

4. Law of Conservation
Law of Conservation of Energy states that energy is neither created nor destroyed, just changed in form
C8H18 + O2  H2O + CO2 + Energy
Stored PE converts to 25% work and 75% heat
(ENERGY)

5. Exothermic Reactions HOT PACK
An exothermic reaction is when the system releases energy; heat flows out of a reaction and the surroundings get warmer. They have a NEGATIVE H.
H products < H reactants 
4Fe + 3 O2  2 Fe2O kJ OR
4Fe + 3 O2  2 Fe2O H = kJ

6. Endothermic Reactions
COLD PACK
An endothermic reaction is when the system absorbs energy; heat flows into a reaction and the surroundings get cooler. They have a POSITIVE H
H products > Hreactants
 27kJ + NH4NO3(s) NH4(aq)+1+NO3(aq) OR
NH4NO3(s) NH4(aq)+1 + NO3(aq)-1H = + 27 kJ

7. Reaction Co-ordinates

8. What is the difference between Temperature & Heat?
Instrument: thermometer
Units: Celsius, Fahrenheit, Kelvin
Definition:
A measure of the average kinetic energy of the molecules in a substance
A measure of the motions of the molecules
A measure of how hot or cold something is

9. What is the difference between Temperature & Heat?
Instrument: calorimeter
Units: calories, joules
Definition:
The total amount of energy in a substance.
A form of energy that is transferred between objects because one is warner than the other.
Heat transfer is always from hot to cold
Depends on 3 things:
1. amount of substance (mass)
2. Temperature change
3. type of material (specific heat)

10. Units of Heat Energy
A calorie is defined as the amount of heat needed to raise the temperature of 1 g of water by 1 C
1 cal= J
Food “Calories” are kilocalories.
1kcal = 1000 calories.

11. Temperature ≠ Heat
Greater Thermal Energy

12. Specific Heat
Amount of heat required to raise the temperature of 1 g of a substance by 1 C
Different substances have different specific heats.
Water has a specific heat of J/gC.
Iron(Fe) has a specific heat of .449 J/gC.
Gold (Au) has a specific heat of .129 J/gC.
The higher the specific heat the more energy it takes to change its temperature.

13. Calculating Heat mT c= specific heat q = heat in joules or galories
m= mass
T = change in temperature = Tf – Ti
c= q_
mT

14. Example
A 155 g sample of an unknown substance was heated from 25.0 C to 40.0 C. The substance absorbed 5696 J of energy. What is the specific heat?

15. Example
How much heat is needed to change the temperature of 12.0 g of silver with a specific heat of cal/g°C from 25.0°C to 83.0 °C?

16. Measuring Heat in a Calorimeter
A coffee cup calorimeter measures heat at constant pressure; works on the premise that the amount of heat released in a reaction(-q) or physical change is equal to the amount of heat absorbed by the water(+q) q = +q 
Rearrange the specific heat equation:
q = m x c x T

17. Example
A piece of unknown metal with mass g is heated to an initial temperature of °C and dropped into g of water (with an initial temperature of °C) in a calorimeter. The final temperature of the system is 30.05°C. What is the specific heat of the metal? Specific heat of water = J/g°C

18. Example
A gram sample of vanadium was heated to °C (its initial temperature). It was then dumped into a calorimeter. The initial temperature of the calorimeter’s water was °C. After the metal was allowed to release all its heat to the calorimeter’s water, °C was the final temperature. What mass of distilled water was in the calorimeter?
Specific heat of vanadium = J/gC Specific heat of water = J/g°C

19. Measuring Heat during Phases Changes

20. Heat of Fusion/Solidification
Heat of fusion (Hfus ) is the heat energy required to melt one gram of a solid at its melting point
For water, Hfus = 334 J/g
q = Hfus x mass
Heat of solidification (Hsolid ) is the heat energy lost when one gram of a liquid freezes to a solid at its freezing point
For water, Hsolid = -334 J/g
q = Hsolid x mass

21. Heat of Vaporization/Condensation
Heat of vaporization (Hvap) is the heat to vaporize one gram of a liquid at its normal boiling point
For water, Hvap= J/g
q = Hvap x mass
Heat of condensation (Hcond ) is the heat energy released when one gram of a liquid forms from its vapor
For water, Hcond = J/g
q = Hcond x mass

22. Example
How much heat is needed to melt 500.0g of ice at 0 C?

23. Example
How much heat is evolved when 1255 g of water condenses to a liquid at 100°C?