1. Chapter 10
Energy

2. 10.3 Exothermic and Endothermic Processes 10.4 Thermodynamics
The Nature of Energy
Temperature and Heat
10.3 Exothermic and Endothermic Processes
Thermodynamics
10.5 Measuring Energy Changes
10.6 Thermochemistry (Enthalpy)
10.7 Hess’s Law
Quality Versus Quantity of Energy
10.9 Energy and Our World
10.10 Energy as a Driving Force
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3. Ability to do work or produce heat.
Energy
Ability to do work or produce heat.
That which is needed to oppose natural attractions.
Law of conservation of energy – energy can be converted from one form to another but can be neither created nor destroyed.
The total energy content of the universe is constant.
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4. Potential energy – energy due to position or composition.
Kinetic energy – energy due to motion of the object and depends on the mass of the object and its velocity.
Telekinesis Video Part 1
Telekinesis Video Part 2

5. Energy Transformations
Types of Energy
1. Kinetic energy
2. Potential energy
3. Chemical energy
4. Heat energy
5. Electric energy
6. Radiant energy

6. Initial Position
In the initial position, ball A has a higher potential energy than ball B.

7. Final Position
After A has rolled down the hill, the potential energy lost by A has been converted to random motions of the components of the hill (frictional heating) and to the increase in the potential energy of B.

8. Both of the above are at 70oF. Which has the most heat?
Temperature = average kinetic energy.
Heat= total energy.
Both of the above are at 70oF. Which has the most heat?

9. A measure of the random motions of the components of a substance.
Temperature
A measure of the random motions of the components of a substance.
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10. Heat
A flow of energy between two objects due to a temperature difference between the objects.
Heat is the way in which thermal energy is transferred from a hot object to a colder object.
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11. Energy Changes Accompanying the Burning of a Match
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12. Energy Changes and Chemical Reactions
4 C3H5N3O N H2O CO2 + O kJ Nitroglycerine Exothermic

13. Endothermic Process:
Heat flow is into a system.
Absorb energy from the
surroundings. q = positive(+)
Exothermic Process:
Energy flows out of the
system. q = negative (-)
Energy gained by the surroundings must be equal to the energy lost by the system.

14. Concept Check
Is the freezing of water an endothermic or exothermic process? Explain.
Exothermic process because you must remove energy in order to slow the molecules down to form a solid.
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15. Concept Check
Classify each process as exothermic or endothermic. Explain. The system is underlined in each example.
Your hand gets cold when you touch ice.
The ice gets warmer when you touch it.
Water boils in a kettle being heated on a stove.
Water vapor condenses on a cold pipe.
Ice cream melts.
Exo
Endo
Exothermic (heat energy leaves your hand and moves to the ice)
Endothermic (heat energy flows into the ice)
Endothermic (heat energy flows into the water to boil it)
Exothermic (heat energy leaves to condense the water from a gas to a liquid)
Endothermic (heat energy flows into the ice cream to melt it)
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16. Concept Check
For each of the following, define a system and its surroundings and give the direction of energy transfer.
Methane is burning in a Bunsen burner in a laboratory.
Water drops, sitting on your skin after swimming, evaporate.
System – methane and oxygen to produce carbon dioxide and water; Surroundings – everything else around it; Direction of energy transfer – energy transfers from the system to the surroundings (exothermic)
System – water drops; Surroundings – everything else around it; Direction of energy transfer – energy transfers from the surroundings (your body) to the system (water drops) (endothermic)
System=burner, surroundings= lab, energy to system
System=drops, surroundings= air, energy to surroundings
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17. Hydrogen gas and oxygen gas react violently to form water.
Concept Check
Hydrogen gas and oxygen gas react violently to form water.
Which is lower in energy: a mixture of hydrogen and oxygen gases, or water? Explain.
Water is lower in energy because a lot of energy was released in the process when hydrogen and oxygen gases reacted.
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18. Study of energy
Law of conservation of energy is often called the first law of thermodynamics.
The energy of the universe is constant.

19. The common energy units for heat are the calorie and the joule.
calorie – the amount of energy (heat) required to raise the temperature of one gram of water 1oC.
Joule – 1 calorie = joules

20. Example
Convert 60.1 cal to joules.

21. The amount of substance being heated (number of grams).
Energy (Heat) Required to Change the Temperature of a Substance Depends On:
The amount of substance being heated (number of grams).
The temperature change (number of degrees).
The identity of the substance.
Specific heat capacity is the energy required to change the temperature of a mass of one gram of a substance by one Celsius degree.
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22. Specific Heat Capacities of Some Common Substances
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23. To Calculate the Energy Required for a Reaction:
Energy (heat) required, Q = s × m × ΔT
Since q=ΔH (change of heat) ΔH = s x m x ΔT
Q = energy (heat) required (J)
s = specific heat capacity (J/°C·g)
m = mass (g)
ΔT = change in temperature (°C)
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24. Energy of a Snickers Bar
There are 311 Cal in a Snickers Bar. If you added this much heat to a gallon of water at 25oC what would be the final temperature?
311 Cal = 311,000 cal
1 gallon of water weighs g
Heat = SH x m x ΔT
311,000 cal = cal/goC x g x ?oC
Solving for the change of temperature gives us 78.5oC
Add this to the 25oC to get 104oC which is above the boiling point of water (100oC)

25. The mass of water and the temperature increase.
Example
Calculate the amount of heat energy (in joules) needed to raise the temperature of 6.25 g of water from 21.0°C to 39.0°C.
Where are we going?
We want to determine the amount of energy needed to increase the temperature of 6.25 g of water from 21.0°C to 39.0°C.
What do we know?
The mass of water and the temperature increase.
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26. What information do we need?
Example
Calculate the amount of heat energy (in joules) needed to raise the temperature of 6.25 g of water from 21.0°C to 39.0°C.
What information do we need?
We need the specific heat capacity of water.
4.184 J/g°C
How do we get there?
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27. Exercise
A sample of pure iron requires 142 cal of energy to raise its temperature from 23ºC to 92ºC. What is the mass of the sample? (The specific heat capacity of iron is 0.45 J/gºC.)
a) g
b) 4.6 g
c) 19 g
d) 590 g
The correct answer is c.
Q = s×m×ΔT
m = Q/ s×ΔT
m = {(142 cal)(4.184 J/1 cal)} / {(0.45 J/g°C)(92°C–23°C)}
m = 19 g

28. (90o(100g)+ 10o(500g))/600gT = 23o Concept Check
A g sample of water at 90.°C is added to a g sample of water at 10.°C.
The final temperature of the water is:
a) Between 50°C and 90°C
b) 50°C
c) Between 10°C and 50°C
Calculate the final temperature of the water.
The correct answer is c). The final temperature of the water is 23°C.
- (100.0 g)(4.18 J/°C·g)(Tf – 90.) = (500.0 g)(4.18 J/°C·g)(Tf – 10.)
Tf = 23°C
(90o(100g)+ 10o(500g))/600gT = 23o

29. The final temperature of the water is: a) Between 50°C and 90°C
Concept Check
You have a Styrofoam cup with 50.0 g of water at 10.C. You add a 50.0 g iron ball at 90.C to the water. (sH2O = 4.18 J/°C·g and sFe = 0.45 J/°C·g)
The final temperature of the water is:
a) Between 50°C and 90°C
b) 50°C
c) Between 10°C and 50°C
Calculate the final temperature of the water. 18°C
The correct answer is c). The final temperature of the water is 18°C.
- (50.0 g)(0.45 J/°C·g)(Tf – 90.) = (50.0 g)(4.18 J/°C·g)(Tf – 10.)
Tf = 18°C
50.0gFe(0.45SHFe)(90-XoC)=50.0gw(4.18SHw)(X-10oC)

30. Consider the reaction: S(s) + O2(g) → SO2(g)
Example
Consider the reaction:
S(s) + O2(g) → SO2(g)
ΔH = –296 kJ per mole of SO2 formed. Calculate the quantity of heat released when 2.10 g of sulfur is burned in oxygen at constant pressure.
Where are we going?
We want to determine ΔH for the reaction of 2.10 g of S with oxygen at constant pressure.
What do we know?
When 1 mol SO2 is formed, –296 kJ of energy is released.
We have 2.10 g S.
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31. Consider the reaction: S(s) + O2(g) → SO2(g)
Example
Consider the reaction:
S(s) + O2(g) → SO2(g)
ΔH = –296 kJ per mole of SO2 formed. Calculate the quantity of heat released when 2.10 g of sulfur is burned in oxygen at constant pressure.
How do we get there?
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32. Consider the combustion of propane:
Exercise
Consider the combustion of propane:
C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l)
ΔH = –2221 kJ
Assume that all of the heat comes from the combustion of propane. Calculate ΔH in which 5.00 g of propane is burned in excess oxygen at constant pressure.
–252 kJ
(5.00 g C3H8)(1 mol / g C3H8)(-2221 kJ / mol C3H8)
ΔH = -252 kJ

33. Enthalpy, H is measured using a calorimeter.
Calorimetry
Enthalpy, H is measured using a calorimeter.

34. In going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.

35. N2(g) + 2O2(g) → 2NO2(g) ΔH1 = 68 kJ
This reaction also can be carried out in two distinct steps, with enthalpy changes designated by ΔH2 and ΔH3.
N2(g) + O2(g) → 2NO(g) ΔH2 = 180 kJ
2NO(g) + O2(g) → 2NO2(g) ΔH3 = – 112 kJ
N2(g) + 2O2(g) → 2NO2(g) ΔH2 + ΔH3 = 68 kJ
ΔH1 = ΔH2 + ΔH3 = 68 kJ
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36. Characteristics of Enthalpy Changes
If a reaction is reversed, the sign of ΔH is also reversed.
The magnitude of ΔH is directly proportional to the quantities of reactants and products in a reaction. If the coefficients in a balanced reaction are multiplied by an integer, the value of ΔH is multiplied by the same integer.
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37. Consider the following data:
Example
Consider the following data:
Calculate ΔH for the reaction
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38. Problem-Solving Strategy
Work backward from the required reaction, using the reactants and products to decide how to manipulate the other given reactions at your disposal.
Reverse any reactions as needed to give the required reactants and products.
Multiply reactions to give the correct numbers of reactants and products.
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39. Reverse the two reactions:
Example
Reverse the two reactions:
Desired reaction:
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40. Example
Multiply reactions to give the correct numbers of reactants and products:
4( ) 4( )
3( ) 3( )
Desired reaction:
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41. Final reactions: Desired reaction: ΔH = +1268 kJ Example
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42. -792/2 + 297 = -99 kJ Concept Check Calculate ΔH for the reaction:
SO2 + ½O2 → SO3 ΔH = ?
Given: (1) S + O2 → SO ΔH = –297 kJ (2) 2S + 3O2 → 2SO ΔH = –792 kJ
a) –693 kJ
b) 101 kJ
c) 693 kJ
d) –99 kJ
The correct answer is d.
We have to rearrange the equations so that they add up to the desired equation.
First, reverse equation (1) to give : SO2 → S + O ΔH = +297 kJ
Then add ½ equation (2) : S + 3/2O2 → SO ΔH = –396 kJ
The two new eqns add up: SO2 + ½O2 → SO3 ΔH = –99 kJ
Turn the first reaction around and divide the second reaction by 2. Do the same with the ΔH’s.
-792/ = -99 kJ

43. Carbon based molecules from decomposing plants and animals
Fossil Fuels
Carbon based molecules from decomposing plants and animals
Energy source for United States
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44. Thick liquids composed of mainly hydrocarbons.
Petroleum
Thick liquids composed of mainly hydrocarbons.
Hydrocarbon – compound composed of C and H.
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45. Gas composed of hydrocarbons.
Natural Gas
Gas composed of hydrocarbons.
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46. Coal
Formed from the remains of plants under high pressure and heat over time.
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47. Effects of Carbon Dioxide on Climate
Greenhouse Effect
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48. Effects of Carbon Dioxide on Climate
Atmospheric CO2
Controlled by water cycle
Could increase temperature by 10oC

49. Global Warming

50. Global Warming

51. Global Warming

52. Global Warming

53. New Energy Sources
Solar
Nuclear
Biomass
Wind
Synthetic fuels

54. Consider a gas trapped as shown
Natural processes occur in the direction that leads to an increase in the disorder of the universe.
Example
Consider a gas trapped as shown
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55. What happens when the valve is opened?
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56. Gives off energy when solid forms. Spreads out when dissolves.
Two Driving Forces
Gives off energy when solid forms. Spreads out when dissolves.
Energy spread
Matter spread
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57. In a given process concentrated energy is dispersed widely.
Energy Spread
In a given process concentrated energy is dispersed widely.
This happens in every exothermic process.
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58. Molecules of a substance spread out to occupy a larger volume.
Matter Spread
Molecules of a substance spread out to occupy a larger volume.
Processes are favored if they involve energy and matter spread.
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59. Entropy, S
Function which keeps track of the tendency for the components of the universe to become disordered.
Measure of disorder or randomness

60. Entropy, S
What happens to the disorder in the universe as energy and matter spread?
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61. Second Law of Thermodynamics
The entropy of the universe is always increasing.
The heat death of the universe will occur when all particles of matter ultimately have the same average kinetic energy and exist in a state of maximum disorder.
From this
To this