1. Thermochemistry

2. Thermochemistry
Thermochemistry is concerned with the heat changes that occur during chemical reactions and changes in state. Thermal energy is heat energy
Energy is the capacity to do work or supply heat measured in Joules or calories 1000J = 1 Kilojoule (KJ)

3. Energy Transformations
17.1
Energy Transformations
Kinetic energy
Energy of motion or the energy an object possesses due to motion of its particles
Potential Energy stored energy
Atomic - energy stored in the nucleus of an atom
Chemical energy -energy stored in chemical bonds
TE = PE + KE

4. Energy Transformations
17.1
Energy Transformations
When fuel is burned in a car engine, chemical potential energy is released and is used to do work. 4

5. Energy Transformations
17.1
Energy Transformations
Energy Transformations
In what direction does heat flow
Misconceptions 5

6. Energy Transformations
17.1
Heat, represented by q, is thermal energy that transfers from one object to another because of a temperature difference between them.
Heat always flows from a warmer object to a cooler object.
WHERE IT’S HOT TO WHERE IT’S NOT
O C
37 C
Low
High 6

7. Exothermic and Endothermic Processes
17.1
Exothermic and Endothermic Processes
Exothermic and Endothermic Processes
What happens in endothermic and exothermic processes? Heat Flow
When fuel is burned in a car engine, chemical potential energy is released and is used to do work. 7

8. Exothermic and Endothermic Processes
17.1
Exothermic and Endothermic Processes
Exothermic (-)
A process that releases heat to its surroundings
Liquid → Solid
A + B → AB + nrg
Endothermic (+)
A process that absorbs heat from its surroundings.
Solid→Liquid
AB + nrg → A B 8

9. Exothermic and Endothermic Processes
17.1
Exothermic and Endothermic Processes
The law of conservation of energy states that in any chemical or physical process, energy is neither created nor destroyed. All energy can be accounted for as work, stored energy or heat 9

10. Temperature
Temperature is not a form of energy it is a measurement of the average KE of the particles in the substance
Two scales in Chemistry

11. Temperature Scales and Interconversions
Kelvin ( K ) - The “Absolute temperature scale” begins at absolute zero and only has positive values.
Celsius ( oC ) - The temperature scale used by science, formally called centigrade, most commonly used scale around the world; water freezes at 0oC, and boils at 100oC.
Fahrenheit ( oF ) - Commonly used scale in the U.S. for our weather reports; water freezes at 32oF and boils at 212oF.
Kelvin = oC
oC = Kelvin
oF = (9/5) oC + 32
oC = [oF ] 5/9 11

12. Kelvin Scale
Absolute Zero is the lowest point at which there is no motion or kinetic energy
C = K – 273
K = C

13. Kinetic Molecular Theory

14. crystallization

15. 13.3
A Model for Solids
The general properties of solids reflect the orderly arrangement of their particles and the fixed locations of their particles.

16. Crystal Structure and Unit Cells
13.3
Crystal Structure and Unit Cells
In a crystal, the particles are arranged in an orderly, repeating, three-dimensional pattern called a crystal lattice. 16

17. A Model for Solids Occurs at 0 Celsius or 273 Kelvin 13.3
The melting point (mp) is the temperature at which a solid changes into a liquid.
Occurs at 0 Celsius or 273 Kelvin 17

18. Theory of Physical Phase
Solids
Strong attractive forces
Liquids
intermediate attractive forces
Gases
-very weak forces therefore no definite shape or volume

19. Sublimation
Process in which a solid changes directly to a vapor without passing through the liquid phase
Solid CO2
Solid I2

20. 13.2
A Model for Liquids
Substances that can flow are referred to as fluids. Both liquids and gases are fluids. 20

21. 13.2
Evaporation
Evaporation occurs at the surface of the liquid and is the process a liquid changing to a gas
It occurs below the boiling point
It is a cooling process
Water releases a lot of heat as it cools. a) During freezing weather, farmers protect citrus crops by spraying them with water. The ice that forms has a protective effect as long as its temperature does not drop much below 0°C. b) Because it is mostly water, the filling of a hot apple pie is much more likely to burn your tongue than the crust. 21

22. 13.2
Evaporation
In an open container, molecules that evaporate can escape from the container. 22

23. 13.2
Evaporation
During evaporation, only those molecules with a certain minimum kinetic energy can overcome the attraction of nearby molecules and escape from the surface of the liquid
Water releases a lot of heat as it cools. a) During freezing weather, farmers protect citrus crops by spraying them with water. The ice that forms has a protective effect as long as its temperature does not drop much below 0°C. b) Because it is mostly water, the filling of a hot apple pie is much more likely to burn your tongue than the crust. 23

24. Evaporation Animation 15
Observe the phenomenon of evaporation from a molecular perspective. 24

25. 13.2
Evaporation
In a closed container, the molecules cannot escape. They collect as a vapor above the liquid. Some molecules condense back into a liquid. 25

26. Condensation
Process in which vapor or gas changes to a liquid

27. Phase Changes Heating & Cooling Curves
As heat is added the temperature increases (diagonal lines) ,therefore KE increases and PE remains constant
As heat is removed & temperature decreases, KE decreases and PE remains the same or constant

28. Phase Changes Heating & Cooling Curves
During a phase change you will notice that temperature and Kinetic Energy remain constant, additional heat causes an increase in the potential energy of the substance and this energy is stored

29. Phase Changes Heating & Cooling Curves https://www. youtube. com/watch
During a phase change, none of the heat is being used to increase the randomly moving particles. Heat is needed to change ice at 0oC to water at 0oC. The particles are rearranging themselves. Energy is stored as potential energy.
The mp and fp take place at the same temperature

30. Phase Change Examples Exothermic: Endothermic: Gas , Liquid , Solid
KE decreases on diagonal lines and PE decreases during phase changes
Endothermic:
Solid, Liquid, Gas
KE increases on diagonal lines and PE increases on (horizontal lines) during phase change

31. Heating & Cooling Curve

32. Latent Heat Heats of Fusion and Solidification
17.3
Latent Heat
Heats of Fusion and Solidification
How does the quantity of heat absorbed by a melting solid compare to the quantity of heat released when the liquid solidifies? 32

33. Heats of Fusion and Solidification
17.3
Heats of Fusion and Solidification
The quantity of heat absorbed by a melting solid is exactly the same as the quantity of heat released when the liquid solidifies; that is, Hfus = Hsolid.
Represented as an equation
H2O (s) + energy H20 (l) ( melting )
H2O (l) H2O(s) + energy (freezing) 33

34. 17.3 34

35. Heats of Fusion and Solidification
17.3
Heats of Fusion and Solidification
The heat of fusion (Hfus) is the quantity of heat needed to change a unit mass of a solid to a liquid at constant temperature.
It takes 334J/g to change ice to liquid water at 0oC at STP.
Hf = 334 J /g see table B
Melting 334 J/g is absorbed
Freezing 334 J/g is released
The heat of solidification (Hsolid) is the quantity of heat needed to change a unit mass of a liquid to a solid at constant temperature J/g is released 35

36. 17.3 36

37. Heats of Vaporization and Condensation
17.3
Heats of Vaporization and Condensation
Heats of Vaporization and Condensation
How does the quantity of heat absorbed by a vaporizing liquid compare to the quantity of heat released when the vapor condenses? 37

38. Heats of Vaporization and Condensation
17.3
Heats of Vaporization and Condensation
The quantity of heat absorbed by a vaporizing liquid is exactly the same as the quantity of heat released when the vapor condenses; that is,
∆Hvap = ∆Hcond. Hv = 2260 J/g (table B)
Heat of vaporization is the quantity of energy needed to vaporize a unit mass of liquid at a constant temperature
2260J/g is absorbed to vaporize
2260 J/g is released to condense

39. Represented as an equation
H20 (l) + energy H20 (g)
Boiling or vaporizing
H20 (g) H2O (l) + energy
Condensing
These examples represent Physical changes
H2O on both sides of equation

40. Heats of Vaporization and Condensation
Animation 21
Observe the phase changes as ice is converted to steam when heat is added.

41. 13.2
Vapor Pressure
Vapor pressure is a measure of the force exerted by a gas above a liquid. See Table H 41

42. Vapor Pressure Pressure exerted by vapor
The greater the vapor pressure of a substance the weaker the IMF (forces holding molecules in liquid phase)

43. Vapor Pressure & Boiling
Normal Boiling Point (nBP)
Temperature at which the vapor pressure is equal to the atmospheric (air) pressure
Normal BP is the Boiling point at standard pressure kpa

44. Would you need more energy to increase the temperature of a bathtub filled with 400,000 ml of water 0.5o C or to increase a glass of a water filled with 250ml by 80oC?

45. Calorimetry Measures the Quantity of heat released or absorbed
17.2
Calorimetry
Measures the Quantity of heat released or absorbed
Calorimetry is the precise measurement of the heat flow into or out of a system for chemical and physical processes. Measured by a Calorimeter 45

46. Units for Measuring Heat Flow
17.1
Units for Measuring Heat Flow
Heat flow is measured in two common units, the calorie and the joule.
The energy in food is usually expressed in Calories.
Food calories
A calorie is defined as the quantity of heat needed to raise the temperature of 1 g of pure water by 1 Celsius

47. 17.2
Calorimetry
The insulated device used to measure the absorption or release of heat in chemical or physical processes is called a calorimeter. 47

48. Joule The SI unit of heat and energy
Increases 1 gram of water by C
4.18 J = 1 calorie
When one pound of coal burns, it releases 3.5 x 106 calories. Convert this amount to joules. Convert to kilojoules
1.5 x 107 Joules
1.5 x 104 KJ

49. 17.2
Calorimetry
Calorimetry experiments can be performed at a constant volume using a bomb calorimeter.
In a simple constant-pressure calorimeter, a thermometer records the temperature change as chemicals react in water. The substances reacting in solution constitute the system. The water constitutes the surroundings. 49

50. PROBLEMS Quantity of heat transferred can be measured by 3 factors
Mass (m) of water in the calorimeter
The change in the temperature (Tfinal – Tinitial)
Constant (specific heat of a substance)
q = mc t

51. Nutritionists use bomb calorimeters to measure the energy content of the foods you eat. To see some of their data (expressed in Calories per serving), you can look at a nutrition label. Calculating According to the nutrition label below, each serving contains 140 Calories. How many kilojoules is this? 51

52. Calculate the amount of heat energy that must be added to a 48g sample of liquid water in order to raise its temperature 30oC

53. Calculate the maximum mass of water that can be heated from room temperature (25oC) to its boiling point by the addition of Joules of heat energy.
65 g

54. Calculate the amount of heat energy that must be removed from a 55g sample of water at 351K in order to lower its temperature to 298K

55. Heat Capacity and Specific Heat
17.1
Heat Capacity and Specific Heat
The heat capacity of an object depends on both its mass and its chemical composition.
The amount of heat needed to increase the temperature of an object exactly 1°C is the heat capacity of that object.
Specific Heat

56. Heat Capacity and Specific Heat
17.1
Heat Capacity and Specific Heat
The specific heat capacity, or simply the specific heat, of a substance is the amount of heat it takes to raise the temperature of 1 g of the substance 1°C.

57. Heat Capacity and Specific Heat
17.1
Heat Capacity and Specific Heat
Water releases a lot of heat as it cools. During freezing weather, farmers protect citrus crops by spraying them with water.

58. Heat Capacity and Specific Heat
17.1
Heat Capacity and Specific Heat 58

59. Heat Capacity and Specific Heat
17.1
Heat Capacity and Specific Heat
Because it is mostly water, the filling of a hot apple pie is much more likely to burn your tongue than the crust.

64. Section Quiz.
1. The energy released when a piece of wood is burned has been stored in the wood as
sunlight.
heat.
calories.
chemical potential energy.

65. Section Quiz.
2. Which of the following statements about heat is false?
Heat is the same as temperature.
Heat always flows from warmer objects to cooler objects.
Adding heat can cause an increase in the temperature of an object.
Heat cannot be specifically detected by senses or instruments.

66. 17.1 Section Quiz.
3. Choose the correct words for the spaces: In an endothermic process, the system ________ heat when heat is ________ its surroundings, so the surroundings _____________.
gains, absorbed from, cool down.
loses, released to, heat up.
gains, absorbed from, heat up.
loses, released to, cool down.

67. 17.1 Section Quiz.
5. Assuming that two samples of different materials have equal mass, the one that becomes hotter from a given amount of heat is the one that
has the higher specific heat capacity.
has the higher molecular mass.
has the lower specific heat capacity.
has the higher density.

68. Theory of Physical Phase
Solids
Close
Strong attractive forces
Liquids
-the molecules that have higher KE escape and enter gas phase
Gases
-very weak forces therefore no definite shape or volume