1. Thermochemistry
Energy

2. Glossary Terms
thermochemistry: the study of energy changes that occur during chemical reactions and changes in state
chemical potential energy: energy stored in chemical bonds
heat (q): energy that transfers from one object to another because of a temperature difference between the objects
system: a part of the universe on which you focus your attention
surroundings: everything in the universe outside the system
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3. Glossary Terms
law of conservation of energy: in any chemical or physical process, energy is neither created nor destroyed
endothermic process: a process that absorbs heat from the surroundings
exothermic process: a process that releases heat to its surroundings
heat capacity: the amount of heat needed to increase the temperature of an object exactly 1°C
specific heat: the amount of heat needed to increase the temperature of 1 g of a substance 1°C; also called specific heat capacity
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4. Energy Transformations
Energy is the capacity for doing work or supplying heat.
Unlike matter, energy has neither mass nor volume.
Energy is detected only because of its effects.
Thermochemistry is the study of energy changes that occur during chemical reactions and changes in state.
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5. 1st Law of Thermodynamics
Energy is always conserved. It can not be created or destroyed.

6. 2nd Law of Thermodynamics
Energy tends to disperse or spread out.

7. ENERGY (E)
Described as the capacity for doing work.

8. WORK The application of force through distance. W = F x d W = work
d = distance

9. FORMS OF ENERGY
Potential energy – energy of position ( stored energy).
Kinetic energy – energy of motion ( energy that is being used).

10. TYPES OF ENERGY Kinetic Potential Mechanic X X Electric X X
Nuclear X X
Radiant X
Chemical X

11. Energy Transformations
Every substance has a certain amount of energy stored inside it.
The energy stored in the chemical bonds of a substance is called chemical potential energy.
The kinds of atoms and the arrangement of the atoms in a substance determine the amount of energy stored in the substance.
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12. Energy Transformations
Every substance has a certain amount of energy stored inside it.
When you buy gasoline, you are actually buying the stored potential energy it contains.
The controlled explosions of the gasoline in a car’s engine transform the potential energy into
useful work, which can be
used to propel the car.
Heat is also produced,
making the car’s engine
extremely hot.
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13. Law of Conservation of Matter and Energy
Matter and energy are interchangeable, the total matter and energy in the universe is constant.
Matter can never be destroyed but converted from one form to another.
Example: Burn a log. What kinds of energy is being converted when a log burns?

14. Heat
The energy transferred between two objects as a consequence of the difference in temperature between the two objects.
Energy in transit.
Measure as a quantity of energy (calories)
Depends on mass.
Heat flows from a region of high heat intensity to a region of low heat intensity.

15. Thermal Equilibrium
When heat flow between two objects is equal, therefore both objects are of equal temperature.
No transfer of heat.
No change in phases.
Opposing tendencies are equal.
Thermal: pertains to heat and also temperature.

16. Temperature
The measure of an object ability to transfer heat to, or acquire heat from other object.
If two objects are in direct contact, energy flow from one to other.
Temperature is the property that determines the heat transfer.
Measure as the heat intensity (degrees) of matter.

17. Example of Temperature and Heat Interaction
If you put your hand in ice water and then in cool water How would your hand feel?
It would feel warm.
But if you put our hand in hot water first and then the cool water. How would your hand feel?
It would feel cold.
Explain
Temperature depends on the transfer of heat energy to the hand or away.
If a body of matter (object), has a greater temperature than its surrounding, energy flows from the object.

18. Example of Temperature
A burning match and a camp fire may both be at the same temperature, but the quantities of heat given off are different.

19. Measuring Temperature
Warm expand and cool contract.
Mercury thermometer is based on this concept.

20. Temperature Scales Temperature Scales
are Fahrenheit, Celsius, and Kelvin.
have reference points for the boiling and freezing points of water.
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21. Three Temperature Scales
Fahrenheit – normally used in U.S., freezing point of water 32° F and boiling point 212 ° F.
Celsius – most used by scientists, freezing point of water 0 ° C and 100° C.
Kelvin – absolute temperature 273 K.
The Celsius scale is determined by dividing the interval between 0° C and 100° C into 100 equal points.

22. Endothermic and Exothermic Processes
What happens to the energy of the universe during a chemical or physical process?
Chemical reactions and changes in physical state generally involve either the absorption or the release of heat.
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23. Endothermic and Exothermic Processes
You can define a system as the part of the universe on which you focus your attention.
Everything else in the universe makes up the surroundings.
Together, the system and its surroundings make up the universe.
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24. Endothermic and Exothermic Processes
Direction of Heat Flow
The direction of heat flow is given from the point of view of the system.
Heat is absorbed from the surroundings in an endothermic process.
Heat flowing into a system from its surroundings is defined as positive; q has a positive value.
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25. Endothermic and Exothermic Processes
Direction of Heat Flow
The direction of heat flow is given from the point of view of the system.
An exothermic process is one that releases heat to its surroundings.
Heat flowing out of a system into its surroundings is defined as negative; q has a negative value.
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26. Endothermic and Exothermic Processes
In an endothermic process, heat flows into the system from the surroundings.
In an exothermic process, heat flows from the system to the surroundings.
In both cases, energy is conserved.
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27. Specific Heat Capacity
Substances with low specific heat capacities can heat up and cool down easily.
Which substance will warm up the fastest?

28. Specific Heats of Some Common Substances
Interpret Data
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.
Specific Heats of Some Common Substances
Substance
Specific heat
J/(g·°C)
cal/(g·°C)
Liquid water
4.18
1.00
Ethanol
2.4
0.58
Ice
2.1
0.50
Steam
1.9
0.45
Chloroform
0.96
0.23
Aluminum
0.90
0.21
Iron
0.46
0.11
Silver
0.24
0.057
Water has a very high specific heat compared with the other substances.
Metals generally have low specific heats.
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29. Measuring Heat
Calorimetry – the measurement of the thermal properties of matter.
Thermal properties of matter are:
Heat of solution
Heat of combustion
Heat of formation
Heat of reaction

30. Units of Heat
Calorie (cal) – the unit of heat energy commonly used in calorimetric.
The quantity of heat required to raise the temperature of 1 gram of water through 1 Celsius degree.
Very small unit.
Kilocalorie (kcal) – most often used.
The quantity of heat required to raise the temperature of 1 kilogram of water through 1 Celsius degree.
1 kcal = 103 cal

31. Calorie is defined as auxiliary heat unit in terms of a dimensional unit of energy, E.
Kinetic energy (EK) – is determine by the mass of the moving body and its velocity (speed).
EK = ½ mv2
m = mass (kg)
v = speed (length/time – unit are meter/sec)
Mass x (length/time)
ml2/t2
EK = Kgm2/s2 is called a joule (J)
1 cal = 4.18 J
1kcal = 4.18 kJ

32. Heat Energy Conversion
Identical units common to both numerator and denominator cancel.
cal = J x 1cal/4.18J
J = 4.18J/1cal
Example:
Given = 23.8 cal
J = 23.8 cal x 4.18J/1cal
J = 99.7J
Given = 99.7J
cal = 99.7J x 1cal/4.18J
cal = 23.8cal

33. Problems
45.6 cal to J
54.8 J to cal
33.1 kcal to kJ
85.5 kJ to kcal

34. Energy Equations Energy (q) = m x ΔT x (Cp) M = mass (g)
ΔT = final temperature – initial temperature (°C)
Cp = specific heat capacity (J/(g·°C) or cal/(g·°C)).
ΔT = q/ m x Cp
Cp = q/ m x ΔT
m = q/C x ΔT
q = m x heat of fusion or vaporization (phase change)

35. Heat Capacity and Specific Heat
Calculating Specific Heat
C = q
m  ΔT =
mass (g)  change in temperature (°C)
heat (J or cal)
q is heat, expressed in terms of joules or calories.
m is mass.
ΔT is the change in temperature.
ΔT = Tf – Ti
The units of specific heat are either J/(g·°C) or cal/(g·°C).
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36. Calculating the Specific Heat of a Substance
Sample Problem 17.2
Calculating the Specific Heat of a Substance
The temperature of a 95.4-g piece of copper increases from 25.0°C to 48.0°C when the copper absorbs 849 J of heat. What is the specific heat of copper?
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37. Analyze List the knowns and the unknown.
Sample Problem 17.2
Analyze List the knowns and the unknown. 1
Use the known values and the definition of specific heat.
KNOWNS
mCu = 95.4 g
ΔT = (48.0°C – 23.0°C) = 23.0°C
q = 849 J
UNKNOWN
C = ? J/(g·°C)
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38. Calculate Solve for the unknown.
Sample Problem 17.2
Calculate Solve for the unknown. 2
Start with the equation for specific heat.
CCu = q
mCu  ΔT
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39. Calculate Solve for the unknown.
Sample Problem 17.2
Calculate Solve for the unknown. 2
Substitute the known quantities into the equation to calculate the unknown value CCu.
CCu = = J/(g·°C)
849 J
95.4 g  23.0oC
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40. Evaluate Does the result make sense?
Sample Problem 17.2
Evaluate Does the result make sense? 3
Remember that liquid water has a specific heat of 4.18 J/(g·°C).
Metals have specific heats lower than water.
Thus, the calculated value of J/(g·°C) seems reasonable.
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41. The specific heat of ethanol is 2. 4 J/(g·°C)
The specific heat of ethanol is 2.4 J/(g·°C). A sample of ethanol absorbs 676 J of heat, and the temperature rises from 22°C to 64°C. What is the mass of ethanol in the sample?
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42. The specific heat of ethanol is 2. 4 J/(g·°C)
The specific heat of ethanol is 2.4 J/(g·°C). A sample of ethanol absorbs 676 J of heat, and the temperature rises from 22°C to 64°C. What is the mass of ethanol in the sample?
C = q
m  ΔT
m = q
C  ΔT
m = = 6.7 g ethanol
676 J
2.4 J/(g·°C)  (64°C – 22°C)
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