1. Chapter 17 Thermochemistry 17.1 The Flow of Energy
17.2 Measuring and Expressing
Enthalpy Changes
17.3 Heat in Changes of State
17.4 Calculating Heats of Reaction
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2. Question of the Day high
Period 1 Day
Relative to other substances water has a _____ specific heat.
high
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3. Question of the Day
Pd 3 Day
1. Why does lava cool faster in water than in air? Do your best to answer.
2. The energy stored in the chemical bonds within a substance is called…
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4. Water requires more energy to raise its temp.
CHEMISTRY & YOU
Heat will flow from the lava to the surroundings until the lava and surroundings are at the same temperature. Air has a smaller specific heat than water. Why would lava then cool more quickly in water than in air?
Water requires more energy to raise its temp.
Therefore, lava in contact with water loses more heat energy than lava in contact with air, allowing it to cool more quickly.
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5. 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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6. 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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7. Energy Transformations
Energy changes occur as either heat transfer or work, or a combination of both.
Heat, represented by q, is energy that transfers from one object to another because of a temperature difference between the objects.
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8. Energy Transformations
Heat flows spontaneously from a warmer object to a cooler object.
If two objects remain in contact, heat will flow from the warmer object to the cooler object until the temperature of both objects is the same.
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9. D. chemical potential energy.
The energy released when a piece of wood is burned has been stored in the wood as
A. sunlight.
B. heat.
C. calories.
D. chemical potential energy.
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10. D. chemical potential energy.
The energy released when a piece of wood is burned has been stored in the wood as
A. sunlight.
B. heat.
C. calories.
D. chemical potential energy.
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11. Period 3 day
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12. Endothermic and Exothermic Processes
What happens to the energy of the universe during a chemical or physical process?
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13. 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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14. Energy
A system is a limited, well-defined part of the universe. It does not gain or lose matter.
A chemical system is a closed area in which a chemical reaction occurs. The reactants and products are stuck inside a container. 14
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15. The surroundings include everything outside the system.
Energy
Stopper & Flask =
Surroundings
The surroundings include everything outside the system.
Chemicals = System 15
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16. 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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17. Endothermic and Exothermic Processes
The law of conservation of energy states that in any chemical or physical process, energy is neither created nor destroyed.
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18. First Law of Thermodynamics
Energy Transfers
First Law of Thermodynamics
Energy lost by the system EQUALS energy gained by the surroundings
Energy gained by the system EQUALS energy lost by the surroundings 18
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19. Endothermic and Exothermic Processes
During any chemical or physical process, the energy of the universe remains unchanged.
If the energy of the system increases during that process, the energy of the surroundings must decrease by the same amount.
If the energy of the system decreases during that process, the energy of the surroundings must increase by the same amount.
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20. 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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21. 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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22. 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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23. Recognizing Endothermic and Exothermic Processes
Sample Problem 17.1
Recognizing Endothermic and Exothermic Processes
On a sunny winter day, the snow on a rooftop begins to melt. As the melted water drips from the roof, it refreezes into icicles. Describe the direction of heat flow as the water freezes. Is this process endothermic or exothermic?
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24. Solve Apply concepts to this situation.
Sample Problem 17.1
Solve Apply concepts to this situation. 2
In order for water to freeze, its temperature must decrease.
Heat flows out of the water and into the air.
Heat is released from the system to the surroundings.
The process is exothermic.
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25. Units for Measuring Heat Flow
Pd 3 Stopped on 2-21
Units for Measuring Heat Flow
Heat flow is measured in two common units:
the calorie
the joule
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26. Endothermic and Exothermic Processes
Units for Measuring Heat Flow
A calorie (cal) is defined as the quantity of heat needed to raise the temperature of 1 g of pure water 1°C.
The word calorie is written with a small c except when referring to the energy contained in food.
The dietary Calorie is written with a capital C.
One dietary Calorie is equal to one kilocalorie, or 1000 calories.
1 Calorie = 1 kilocalorie = 1000 calories
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27. Endothermic and Exothermic Processes
Units for Measuring Heat Flow
The joule (J) is the SI unit of energy.
One joule of heat raises the temperature of 1 g of pure water °C.
You can convert between calories and joules using the following relationships:
1 J = cal
4.184 J = 1 cal
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28. Athletes often use instant cold packs to soothe injuries
Athletes often use instant cold packs to soothe injuries. Many of these packs use the dissociation of ammonium nitrate in water to create a cold-feeling compress. Is this reaction endothermic or exothermic? Why?
The instant cold pack feels cold because it removes heat from its surroundings. Therefore, the dissociation of ammonium nitrate in water is endothermic. The system (the cold pack) gains heat as the surroundings lose heat.
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29. Heat Capacity and Specific Heat
Heat capacity = amount of heat needed to increase the temperature of an object exactly 1°C.
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30. Heat Capacity and Specific Heat
The heat capacity of an object depends on both its mass and its chemical composition.
The greater the mass of the object, the greater its heat capacity.
A massive steel cable requires more heat to raise its temperature by 1oC than a steel nail does.
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31. 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.
Water has a very high specific heat compared with the other substances.
Metals generally have low specific heats.
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
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32. Heat Capacity and Specific Heat
Specific Heat of Water
Just as it takes a lot of heat to raise the temperature of water, water also releases a lot of heat as it cools.
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33. Review section 17.1 and complete #s 5-11 on page 561.
Assignment
Review section 17.1 and complete #s 5-11 on page 561.
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34. Question of the Day
Pd 3 Day
1. In an exothermic reaction how does heat flow with respect to the system and surroundings (use the words system and surroundings in your answer).
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35. Heat Capacity and Specific Heat
Specific Heat of Water
Water in lakes and oceans absorbs heat from the air on hot days
and releases it back into
the air on cool days.
This property of water is responsible for moderate climates in coastal areas.
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36. Heat Capacity and Specific Heat
Calculating Specific Heat
To calculate the specific heat (C) of a substance, you divide the heat input by the mass of the substance times the temperature change.
C = q
m  ΔT =
mass (g)  change in temperature (oC)
heat (J or cal)
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37. 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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38. 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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39. q CCu = mCu  ΔT 849 J CCu = = 0.387 J/(g·°C) 95.4 g  23.0oC
Sample Problem 17.2
CCu = q
mCu  ΔT
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?
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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43. Key Concepts & Key Equation
Energy changes occur as either heat transfer or work, or a combination of both.
During any chemical or physical process, the energy of the universe remains unchanged.
The heat capacity of an object depends on both its mass and its chemical composition.
C = q
m  ΔT
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44. chemical potential energy: energy stored in chemical bonds
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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45. endothermic process: a process that absorbs heat from the surroundings
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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