1. Energy

2. Thermochemistry
The study of energy changes that occur during chemical reactions and changes in state.

3. The Nature of Energy Energy - the ability to do work or produce heat
Energy is stored in the chemical bonds of the molecules.
Energy exists in two basic forms:
Potential energy
Kinetic energy

4. Potential Energy
Potential Energy - energy due to the composition or position of an object
ex. water behind a dam
chemical potential energy - energy stored in a substance because of its composition
ex. gasoline

5. Kinetic Energy and the law of conservation of energy
The energy of motion
ex. Water falling over a dam
In any chemical reaction or physical reaction, energy can be converted from one form to another, but it is neither created nor destroyed.
Law of conservation of energy-in any chemical reaction or physical reaction, energy can be converted from one form to another, but it is neither created nor destroyed. This is the first law of thermodynamics.
When water falls over a dam, the waters energy goes from potential to kinetic

6. Heat and Temperature
Heat (q)- a form of energy that flows from a warmer object to a cooler object
When an object loses heat the temperature of that object falls
When an object gains heat, the temperature of that object rises

7. Heat vs Temp. Cont’
Temperature - a measure of the average kinetic energy of the particles in a sample of matter
As heat increases, the temperature increase, and the kinetic energy increases.

8. Phase Changes The three phases are solid, liquid, and gas
For a phase change to occur, energy must be lost or gained (with constant pressure)
a substance can go from a solid to a gas, a solid to a liquid, a liquid to a solid, a liquid to a gas, a gas to a solid, or a gas to a liquid

9. Solid to liquid; liquid to solid
Melting - the temperature at which the forces holding a crystal lattice together are broken an it becomes a liquid
Freezing - the temperature at which a liquid is converted into a crystalline solid

10. Liquid to gas; gas to liquid
Vaporization - the process by which a liquid changes from a liquid to a gas
evaporation - vaporization only at the surface
When a dog pants, water is evaporated off of it’s tongue. – We sweat so the water is evaporated from our skin, which cools the body
boiling point - the temperature at which vapor pressure of a liquid is equal to the external or atmospheric pressure

11. Liquid to gas; gas to liquid
Vapor pressure - the pressure exerted by a vapor over a liquid
condensation - when a gas or vapor becomes a liquid

12. Solid to gas; gas to solid
Sublimation - the process by which a solid changes directly into a gas without becoming a liquid first.
Deposition - the process by which a substance changes from a gas or vapor to a solid without becoming a liquid

13. Phase Change Diagram

14. Measuring Heat
The flow of energy and the resulting change in temperature are determining factors in how heat is measured
The SI unit of heat and energy is the joule (J)

15. Specific Heat
The amount of heat required to raise the temperature of one gram of any substance by one degree Celsius
Different substances have different specific heats
The specific heat of water J/g•°C

16. calories
One calorie is the energy required to raise the temperature of one gram of water by one degree Celsius
Specific heat! So one calorie is J
Kilocalorie?
Calorie on food labels- one food Calorie equals 1000 chemistry calories= 1kcal
Carbohydrates, Fats, and Proteins provide calories
Energy is stored in the chemical bonds

17. Endothermic and Exothermic reactions
Endothermic- a chemical reaction or process in which a greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds form the products. (energy in) Energy moving into the system happens with melting, vaporization/ evaporation, and sublimation. (gaining heat, temperature increases)
Exothermic- a chemical reaction or process in which more energy is released than is required to break bonds in the initial reactants. (energy out) Energy is typically released as heat. Burning something would be an exothermic reaction. Energy moving out of the system takes place in condensation, freezing, and deposition. (losing heat, temperature gets lower)

18. System and Surroundings
System- the specific part of the universe containing the reaction or process being studied. Ex: If you are heating water in a pot on the stove the stove is the system since it contains the reaction you are studying.
Surroundings- everything in the universe except the system. Ex: The pot on the stove and everything around it is the surroundings (air, table, etc).

19. Enthalpy and Entropy
Enthalpy- the heat content of a system at constant pressure.
Entropy- a measure of the number of ways that the energy of a system can be distributed (free particles).
Enthalpy (H) heat of reaction
∆Hrxn= Hproducts – H reactants

20. Calculation for heat evolved and absorbed
q = m x c x ΔT
q = the heat absorbed or released
m = the mass of the sample in grams
c = the specific heat of a substance
ΔT = change in temperature (final temperature – initial temperature)

21. Question
If the temperature of 35.0 g of water increases from 25.0°C to 50.0°C, how much heat has been absorbed by the water?
Known: m = 35.0 g
c = J/g•°C
ΔT = 50.0°C °C = 25.0°C

22. Answer
q= m x c x ΔT
q = 35.0 g x J/g•°C x 25.0°C
q = 3660 J

23. A 4. 10 g nugget of pure gold absorbed 256 J of heat
A 4.10 g nugget of pure gold absorbed 256 J of heat. What was the final temperature of the gold if the initial temperature was 21.0°C? The specific heat of gold is J/(g·°C).

24. A 4. 10 g nugget of pure gold absorbed 256 J of heat
A 4.10 g nugget of pure gold absorbed 256 J of heat. What was the final temperature of the gold if the initial temperature was 21.0°C? The specific heat of gold is J/(g·°C).
q= m x c x ΔT
ΔT= q/ m x c
ΔT= final temperature – initial temperature
Final temp. = ΔT + initial temp.
ΔT = ____256J__________
4.10 g x 0.129J/(g·°C)
ΔT = °C
Final temp = °C + 21°C = °C

25. The temperature of a sample of water increases from 20. 1°C to 45
The temperature of a sample of water increases from 20.1°C to 45.3°C as it absorbs 5607 J of heat. What is the mass of the sample?

26. The temperature of a sample of water increases from 20. 1°C to 45
The temperature of a sample of water increases from 20.1°C to 45.3°C as it absorbs 5607 J of heat. What is the mass of the sample? (specific heat capacity of water is 4.184J /(g·°C))
q= m x c x ΔT
m= q/c x ΔT
M= ______5607J_______
4.184J/(g·°C) x 25.2°C
M= g