1. hermochemistry

2. Thermochemistry: The study of energy during physical and chemical changes. Objectives:
Be able to define and correctly use energy-related terminology.
Identify and understand endothermic and exothermic processes.

3. Thermochemistry: The study of energy during physical and chemical changes.
Chemical Bonds:
Ability to do work (push things)
Store energy
Chemical potential energy

4. Heat and Thermal Energy:
Heat: represented by (q)
Thermal energy that flows from warmer to cooler areas.
Thermal energy:
total KE of particles in a substance(depends on #g and temp.)
Enthalpy: (∆H)
The total amount of energy stored in a system at constant pressure. (KE+PE)

5. Heat transfer: conduction— colliding particles
radiation— electromagnetic waves
convection— currents

6. Heat Transfer: Exothermic Process
System releases energy and the surroundings tend to get warmer.
Surroundings
Enthalpy Decreases
-∆H
ENERGY
System

7. Heat Transfer: Endothermic Process
System gains energy from its surroundings. (surroundings get colder)
Surroundings
ENERGY
System
Enthalpy Increases
+∆H

8. Objectives
Be able to define and correctly use the common units of thermal energy.
Be able to define and understand the concept of specific heat.
Be able to make calculations related to thermal energy and temperature changes.

9. Calorimetry Precise calculation for measuring heat transfer
Heat flow is measured in two common units, the calorie (cal) and joule (J).
1 J= cal It takes about 4000 J to heat kg of water by 1oC. (4 J for 1 g)
4.184 J = 1 cal

10. Specific Heat: heat capacity
the amount of energy required to
raise 1.00 g of a substance by 1.00oC.
Molecules such as water, have a high
specific heat (C).
Metals have relatively low C, change
temperature (T) quickly.

11. Energy and Temperature
Temperature is a property of matter whereas heat is transferred energy from one object to another.
Temperature changes involve KE.
Q = m•C•∆T ∆T= Tf-Ti
Q = enthalpy (#J)
+Q = enthalpy increase = endothermic
-Q = enthalpy decrease = exothermic

12. Problems Calorimetry
Example: What is the change in enthalpy when a cup of water (227 g) cools from boiling to room temperature (97oC to 22oC)?
Example: A wedding ring absorbs 16.4 J of energy when it is placed on a finger (the temperature rises from 21oC to 38oC). If the mass of the ring is 4.80 g, what is the “specific heat” of the metal?

13. Objectives
Understand the concept of latent heat and how it corresponds to potential energy.
Be able to make latent heat calculations.

14. Latent Heat: Changes in state involve changes in Potential Energy.
This stored energy is called latent heat or the quantity of heat absorbed or released by a substance undergoing change of state.
KE (temperature) is constant during a phase change.
Draw diagram.

15. Latent Heat Values:
Size of value depends on the strength of intermolecular bonds

16. Flathead Cherries
Cherries are sprayed with water to protect them from freezing, Why?

17. Latent Heat Calculations
Temperature remains constant, so we use:
Q = (m/M) · ∆H
How much energy is needed to boil g of ethanol (CH3CH2OH)?
How much water (at 0oC) is freezing if 2.5 kJ of energy is released?

18. Objectives
Be able to draw a heating curve or cooling curve for a substance.
Be able to correctly label the regions where ∆KE and ∆PE are occurring on a heating or cooling curve.

19. Heating and Cooling Curves
∆ PE = molecules
pulled apart when boiling
changes of state
Q = m/M · ∆H
∆KE = molecules
speed up
∆ PE = molecules
pulled apart when melting
∆ KE = molecules
speed up
Imagine heating an ice cube. What energy changes take place as it is continually heated?
heating or cooling
Q = m · ∆T · C
∆KE = molecules speed up

20. Objectives Understand the concept of a standard heat of formation.
Be able to calculate the heat of reaction using Hess’s Law and determine if a reaction is endothermic or exothermic.

21. Standard Heat of Formation
standard heat of formation (∆Hf0): change in enthalpy that accompanies the formation of one mole of a compound from its elements at 25oC and 101.3kPa.
∆Hf0 for any uncombined element in its normal state = 0 kJ/mol

22. Q = S(n·∆Hf0)products - S(n·∆Hf0)reactants
Hess’s Law
heat of reaction (Q): the change in enthalpy (energy lost or gained) in a chemical reaction
Use Hess’s Law to calculate the heat of reaction:
Q = S(n·∆Hf0)products - S(n·∆Hf0)reactants