1. Applications of Heat and Energy
Thermochemistry Branch of chemistry dealing with the relationship between chemical action and heat.
Applications of Heat and Energy

2. Energy All energy can be classified as either potential or kinetic.
Potential energy is any type of energy that is stored.
Examples: batteries (stored chemical
converts to electrical)
Top of a roller coaster hill (gravitational)
Candy Bar (stored energy in its chemical bonds)
Kinetic energy – any energy from the movement of matter.

3. Which has more heat - Lake Erie in December of a drop of boiling oil?
Frozen Lake Erie
Drop of boiling oil

4. Which has a higher temperature- Lake Erie in December of a drop of boiling oil?
Frozen Lake Erie
Drop of boiling oil

5. Heat vs. Temperature

6. Heat vs. Temperature
Heat is the amount of energy that is transferred from one substance to another due to the difference in temperature.
Temperature is the average kinetic energy of an object.

7. Heat Transfer
Heat always flows from an object with more heat (hotter) to an object with less heat.
Note: Something that is cold just lacks heat. There is no unit for “cold”.
When you feel cold, it is because you are losing heat (not gaining “cold”)

8. True or False: When you place ice cream in the freezer, heat is transferred from the ice cream to the freezer.
True
False

9. True or False: When we open the window, heat is transferred from your body to the air outside.

10. Units of Heat Heat is measure in Joules (J).
(The joule is the SI unit for all types of energy.)
Example: when you heat a cup of tea, you use about 75,000J (or 75 kilojoules) of heat.
Joule is pronounced jewel

11. Other units of heat Heat is also measured in calories (cal).
1000 calories equals a kilocalorie (Cal).
A calorie is the amount of heat needed to raise the temperature of 1 gram of water by 1 degree Celsius.
Also, 1 calorie = Joules
And … 1 Cal = kJ

12. What is the boiling point of water?
273K
32oF
212oC
100oC

13. What is the freezing point of water?
32K
100oF
273K
32oC

14. Units of Temperature

15. Units of Temperature
TF = 1.8TC + 32
TK = TC + 273

16. Specific Heat
Different substances absorb (and lose) heat at different rates.
Specific heat (SH) is the amount of heat (q) needed to raise the temperature of 1 gram of a substance by 1 degree Celsius.

17. SPECIFIC HEAT q = s x m x DT q = thermal energy (J)
SPECIFIC HEAT: The quantity of heat required to raise the temperature of one gram of a substance by one degree Celsius (or one Kelvin).
q = s x m x DT
q = thermal energy (J)
sh = specific heat (J/goC)
m = mass (g)
DT = change in temperature (Tfinal – Tinitial) (oC)

18. SPECIFIC HEAT
Determine the energy (in kJ) required to raise the temperature of g of water from 20.0 oC to 85.0 oC?
m = g DT = Tf -Ti = oC = 65.0 oC
q = m x s x DT s (H2O) = J/ g - oC
q = (100.0 g) x (4.184 J/g-oC) x (65.0oC)
q = J (1 kJ / 1000J) = 27.2 kJ
Determine the specific heat of an unknown metal that required 2.56 kcal of heat to raise the temperature of g from 15.0 oC to oC?
S = cal /g -oC

19. LAW OF CONSERVATION OF ENERGY
The law of conservation of energy (the first law of thermodynamics), when related to heat transfer between two objects, can be stated as:
The heat lost by the hot object = the heat gained by the cold object
-qhot = qcold
-mh x sh x DTh = mc x sc x DTc
where DT = Tfinal - Tinitial

20. Measuring Heat Changes
A calorimeter is an instrument
used to measure heat changes.
By placing an object in an insulated
container, the heat loss can be
measured by the temperature gain
of the water.
Heat lost by object = Heat gained by water
-qobject = qwater

21. PRACTICE PROBLEM #7
1. Iron metal has a specific heat of J/goC. How much heat is transferred to a 5.00 g piece of iron, initially at 20.0 oC, when it is placed in a beaker of boiling water at 1 atm?
2. How many calories of energy are given off to lower the temperature of g of iron from oC to 35.0 oC?
3. If 3.47 kJ were absorbed by 75.0 g H2O at 20.0 oC, what would be the final temperature of the water?
4. A 100. g sample of water at 25.3 oC was placed in a calorimeter g of lead shots (at 100 oC) was added to the calorimeter and the final temperature of the mixture was 34.4 oC. What is the specific heat of lead?
5. A 17.9 g sample of unknown metal was heated to oC. It was then added to g of water in an insulted cup. The water temperature rose from oC to 23.98oC. What is the specific heat of the metal in J/goC?
180. J
1.23 x 103 cal
31.1 oC
1.28 J/g oC
0.792 J/goC

22. LAW OF CONSERVATION OF ENERGY
Assuming no heat is lost, what mass of cold water at 0.00oC is needed to cool g of water at 97.6oC to 12.0 oC?
-mh x sh x DTh = mc x sc x DTc
- (100.0g) (1 cal/goC) ( oC) = m (1 cal/goC) ( oC)
8560 cal = m (12.0 cal/g)
m = cal / (12.0 cal/g)
m = 713 g
Calculate the specific heat of an unknown metal if a g piece at 100.0oC is dropped into mL of water at 17.8 oC. The final temperature of the mixture was 39.4oC.
s (metal) = cal/g oC

23. GROUP STUDY PROBLEM #7
_____1. A g metal bar requires kJ to change its temperature from 22.0oC to 100.0oC. What is the specific heat of the metal in J/goC?
_____2. How many joules are required to lower the temperature of g of iron from 75.0 oC to 25.0 oC?
_____ 3. If 40.0 kJ were absorbed by g H2O at 10.0 oC, what would be the final temperature of the water?
_____ 4. A 250 g of water at oC is mixed with mL of water at 5.0 oC. Calculate the final temperature of the mixture.
_____5. A 400 g piece of gold at 500.0oC is dropped into 15.0 L of water at 22.0oC. The specific heat of gold is J/goC or cal/goC. Calculate the final temperature of the mixture assuming no heat is lost to the surroundings.

24. A solid becoming a liquid is called:
Freezing
Melting
Evaporation
Sublimation
Condensation
Deposition

25. A gas becoming a liquid is called:
Freezing
Melting
Evaporation
Sublimation
Condensation
Deposition

26. A solid becoming a gas is called:
Freezing
Melting
Evaporation
Sublimation
Condensation
Deposition

27. A liquid becoming a gas is called:
Freezing
Melting
Evaporation
Sublimation
Condensation
Deposition

28. Attractive Forces In gases, these attractive forces are minimal.
In solids and liquids, the forces are strong enough to keep the materials from scattering everywhere.
These attractive forces also determine the melting point and boiling point of different compounds. (ex. NaCl melts at 801o C)

29. Changes in State
Materials experience a change in state when enough heat energy is applied to break apart (or form) the attractions between molecules.
When intermolecular bonds are broken, heat is absorbed from the surroundings; when the bonds are formed, heat is taken from the system and released to the surroundings.

30. Changes in State
Different states of matter (solids, liquids, and gases) have very different properties due to attractive forces that exist between atoms.
To change from a solid to a liquid, for example, these attractive forces in solids must be broken so that the liquid molecules have more freedom to move. In gases, the molecules have even more mobility.

31. When a solid melts, heat is removed from the surrounding environment to break those intermolecular forces.
True
False

32. When a gas condenses, heat is removed from the surrounding environment?
True
False

35. What happens at B? Solid starts to melt Liquid starts to freeze
Gas starts to condense
Liquid starts to evaporate

36. What happens at D? Solid starts to melt Liquid starts to freeze
Gas starts to condense
Liquid starts to boil

37. What happens at E moving from right to left in the graph?
Solid starts to melt
Liquid starts to freeze
Gas starts to condense
Liquid starts to evaporate

38. What happens at C moving from right to left in the graph?
Solid starts to melt
Liquid starts to freeze
Gas starts to condense
Liquid starts to evaporate

39. Heating Curves

40. Melting Point / Freezing Point
The melting point (same temperature as freezing point) is the temperature at which a solid turns to a liquid.
Latent Heat of Fusion – the amount of heat needed to freeze one gram of a substance (or the heat released when one gram of a substance melts.)
Note: units in J/g

41. Heat = mass x Heat of Fusion
Latent Heat of Fusion
Q = m x Hfus
Heat = mass x Heat of Fusion
****For ice to water: Hfus = 334 J/g
(every substance has a difference heat of fusion)

42. Boiling Point / Condensation Point
The boiling point (same temperature as condensation point) is the temperature at which a liquid turns to a gas.
Latent Heat of Vaporization – the amount of heat needed to vaporize one gram of a substance (or the heat release when one gram of a substance condenses.)
Note: units in J/g

43. Latent Heat of Vaporization
Q = m x Hvap
Heat = mass x Heat of Vaporization
****For water to steam: Hvap = 2260 J/g
(every substance has a difference heat of vaporization)

45. Heat Graph calculations
When the graph is flat, use latent heat equations because of change of state.
For melting: Q = m x Hfus
For evaporating: Q = m x Hvap

46. Specific Heat
Different substances absorb (and lose) heat at different rates.
Specific heat (SH) is the amount of heat (q) needed to raise the temperature of 1 gram of a substance by 1 degree Celsius.

47. Heat Graph calculations
When the graph is sloped, use specific heat equations because of change of temperature
Q = SH x m x Temp. Change or

48. Measuring Heat Changes
A calorimeter is an instrument
used to measure heat changes.
By placing an object in an insulated
container, the heat loss can be
measured by the temperature gain
of the water.
Heat lost by object = Heat gained by water
-qobject = qwater

49. Why do chemical reactions occur between some substances and not in others?

50. Chemical reactions occur so that the atoms in the elements involved attain a more stable state of being.

51. Collision Theory
Collision theory – molecules must collide with the proper orientation and sufficient energy to react.

53. Activation Energy
The activation energy is the amount of energy required to break the bonds between the atoms of the reactants.

55. Condition Necessary for Reactions to Occur
Collision: Reactants must collide.
2) Orientation: The reactants must align properly to react.
3) Energy: The activation energy must be attained to react.

56. Energy in Chemical Reactions
Many chemical reactions also produce energy changes.
Definitions:
System – the reactants and products in the reaction
Surroundings – everything else around the reaction (eg air in the room, reaction flask)

57. Heat of Reaction
Heat of Reaction (ΔH) – the amount of heat lost or gained in a reaction
Heat of Reaction: ΔH = Hproducts – Hreactants

58. Exothermic Reactions
Exothermic Reactions – energy is produced by a reaction; energy flows from the system to the surroundings
ΔH is negative because the reaction loses heat.

59. Exothermic Reaction Graph

60. Endothermic Reactions
Endothermic Reactions – energy is gained by a reaction; energy flows from the surroundings into the system
ΔH is positive because the reaction gains heat.

61. Endothermic Reaction Graph

62. Classify the reaction: 2H2O + 572kJ -> 2H2+ O2
Exothermic
Endothermic

63. ΔH of the reaction = -560kJ
Exothermic
Endothermic

64. After the reaction, your hand gets burnt from the heat
After the reaction, your hand gets burnt from the heat. The reaction must be:
Exothermic
Endothermic

65. Heat Values in Chemical Reactions
Heat of Reaction is a stoichiometric value and is proportional to the coefficients of the reactants and products.
2H2O + 572kJ -> 2H2+ O2
Therefore, for every 2 moles of water that react, 572kJ of energy are required.

66. Presence of a Catalyst – a substance that increases the rate without being permanently changed
- lowers activation energy

67. Also used:
inhibitors – “tie up” a reaction so that it does not occur (opposite of a catalyst)
- preservatives
- anti-rust agents