1. How are they related?
Chemistry and Energy

2. Energy Encountered Daily

3. What is Energy? Defined as the ability to do work or create heat.
Many types of energy
Thermal
Light
Gravitational
Kinetic
Potential

4. Light Energy Review How is light energy produced?
Electrons release light energy when they fall from a high energy level to a lower energy.
We’re now going to talk about energy released or used in a chemical reaction. Heat energy

5. Thermochemistry
The study of heat used or released in a chemical reaction.
Let’s investigate heat as it compares to temperature using the Heat vs. Temperature Handout

6. Specific Heat Calculations
q = mCΔT
q = heat (J or cal or Cal)
4.184 cal = 1 Joule
1000 cal = 1 Cal (dietary calorie)
m = mass (g)
C = specific heat (J/g oC or cal/g oC)
ΔT = change in temperature (o C or K)
= Tf - Ti

7. Specific Heat Specific heat of water = 1 cal /g o C or = 4.184 J / goC
Specific heat of most metals = < 1 J / goC
Do metals heat slowly or quickly compared to water?
Do metals stay warm longer or shorter than water?

8. Practice Problem
How much energy is required to heat g of water from 2.0 oC to 24.0oC?
q = mCΔT
m= g
C = J/goC
ΔT= (24.0 – 2.0)oC = 22.0oC
q = (120.0g)(4.184 J/goC)(22.0oC) =

9. Practice Problem
How much heat (in kJ) is given off when g of lead cools from 200.0oC to 10.0 oC? (Specific heat of lead = J/g oC)
q = mCΔT
m = 85.0 g
C = J/g oC
ΔT = (10.0 – 200.0)oC = oC
q = (85.0 g)(0.129 J/g oC)( oC) = -

10. How Do Chemical Reactions Create Heat energy?
Consider the combustion of gasoline (octane)
2 C8H O2  16 CO2 +18 H2O
Potential Energy: Stored energy
Potential energy is stored in the bonds of the reactant s and the products
When bonds are broken, the energy is available
When produce bonds form, some energy is used in these bonds
The excess energy is released as heat

11. Kinetic Energy
Directly related to temperature

12. Is Heat Used or Released?
Endothermic reactions used heat from the surroundings
Sweating
Refrigeration
Exothermic heat releases heat to the surroundings
Hot hands
Combustion
Exercise

13. Endothermic Reactions
Decrease in kinetic energy  decrease in temperature  heat will transfer from the environment to the system resulting in a cooler environment
Absorbs heat from its surrounding.
The system gains heat
Positive value for q
H = q = 0
Hproducts  Hreactants

14. Exothermic Reactions
Increase in kinetic energy  increase in temperature of system heat released to the environment resulting in a hotter environment
Releases heat to its surroundings
The system loses heat
Negative value for q
H = q = 0
Hproducts  Hreactants

15. Enthalpy Heat content for systems at constant pressure Symbol is H
Terms heat and enthalpy are used interchangeably for this course
H = q = m C T
Heat moves from ________ to ___________.

16. Law of Conservation of Energy
Energy is not lost or gained in a chemical reaction
 In a chemical reaction potential energy is transferred to kinetic energy

17. Thermochemical Equations
An equation that includes the heat change
Example: write the thermochemical equation for this reaction
CaO(s) + H2O(l) Ca(OH)2(s) H = kJ
CaO(s) + H2O(l) Ca(OH)2(s) kJ

18. Stoichiometry and Thermochemistry
Tin metal can be extracted from its oxide according to the following reaction:
SnO2(s) + 4NO2(g) + 2H2O(l) kJ  Sn(s) + 4HNO3(aq)
How much energy will be required to extract 59.5 grams of tin?

19. How to solve Use your stoichiometry
Treat heat as a reactant or product
SnO2(s) + 4NO2(g) + 2H2O(l) kJ  Sn(s) + 4HNO3(aq)
59.5 g Sn 1 mol Sn kJ
g Sn 1 mol Sn

20. If an Object feels hot, it means it is undergoing a change with a H that is:
a. positive
b. negative
c. whether the object feels hot or not is unrelated to its H
d. I don’t know

21. If the object feels hot, it means it is undergoing:
a. an exothermic reaction
b. an endothermic reaction
c. whether it feels hot or not is unrelated to whether it is undergoing an exothermic or an endothermic change

22. Heat During a Change of State
How does ice melt?
Heat During a Change of State

23. Molar Heat of Fusion
Heat absorbed by one mole of a substance during melting
Constant temperature
Hfus
H2O(s)  H2O(l) H = 6.01 kJ/mol

24. Molar Heat of Solidification
Heat lost when 1 mole of a liquid solidifies
Temperature is constant
Hsolid
Hfus = -Hsolid
H2O(l)  H2O(s) H = kJ/mol

25. Molar Heat of Vaporization
Heat needed to vaporize 1 mole of a liquid
Hvap
H2O(l)  H2O(g) Hvap = 40.7 kJ/mol

26. Molar Heat of Condensation
Heat released when 1 mole of vapor condenses
Hcond
H2O(g)  H2O(s) Hcond = kJ/mol
Hvap = -Hcond

27. Phase Change Diagram for Water

28. Phase Change Diagram

29. The House that Heats Itself

30. CaLORIMETRY

31. Calorimetry
Method used to determine the heat involved in a physical or chemical change.
Relies on the law of conservation of energy

32. Calorimeter

33. Simple Calorimeter

34. Calorimetry Math Heat gained by the water = q
Heat lost by the system = -q
mCT = q
T = Tf –Ti , m = mass, C = specific heat
q gained by water = q lost by system
q water = - q system
 mCT = -mCT
(mass H2O)(spec. heat H2O)(T H2O) = - (mass sys)(spec. heat sys)(T sys)

35. Standard Heat of Reaction
Heat change for the equation as it is written
H = Hf(products) - Hf(reactants)
Standard Heats of Formation (Hf)
Change in enthalpy when 1 mole of the compound is formed from its elements in their standard states at 25oC and kPa

36. Hess’s Law
 A way to calculate the heat of a reaction that may be too slow or too fast to collect data from.
Add together several reactions that will result in the desired reaction. Add the ΔH for these reactions in the same way.
Htotal = Hproducts - Hreactants