chapter 10 energy. chapter 10 table of contents copyright © cengage learning. all rights reserved 2...

Chapter 10

Energy

Chapter 10

Table of Contents

Copyright © Cengage Learning. All rights reserved 2

10.1 The Nature of Energy

10.2 Temperature and Heat

10.3 Exothermic and Endothermic Processes10.4 Thermodynamics

10.5 Measuring Energy Changes10.6 Thermochemistry (Enthalpy)

10.7 Hess’s Law

10.8 Quality Versus Quantity of Energy

10.9 Energy and Our World

10.10Energy as a Driving Force

Section 10.1

The Nature of Energy

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• 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.

Energy

Section 10.1


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• 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. Energy is the capacity to do work.

• Radiant energy - comes from the sun and is earth’s primary energy source

• Thermal energy - is the energy associated with the random motion of atoms and molecules

• Chemical energy - is the energy stored within the bonds of chemical substances

• Nuclear energy is the energy stored within the collection of neutrons and protons in the atom

Energy

Section 10.1


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• In the initial position, ball A has a higher potential energy than ball B.

Initial Position

Section 10.1


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• 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.

Final Position

Section 10.1


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• Heat involves the transfer of energy between two objects due to a temperature difference.

• Work – force acting over a distance.• Energy is a state function; work and heat are

not: State Function – property that does not

depend in any way on the system’s past or future (only depends on present state). Changes independently of its pathway

Energy

Section 10.1


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State functions are properties that are determined by the state of the system, regardless of how that condition was achieved.

Potential energy of hiker 1 and hiker 2 is the same even though they took different paths.

energy, pressure, volume, temperature

U = Ufinal - Uinitial

P = Pfinal - Pinitial

V = Vfinal - Vinitial

T = Tfinal - Tinitial

Section 10.2

Temperature and Heat

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• A measure of the random motions of the components of a substance.

Temperature

Temperature is a measure of the thermal energy.

Section 10.2

Temperature and Heat

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• 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.

Heat

Section 10.3

Exothermic and Endothermic Processes

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• System – part of the universe on which we wish to focus attention.

• Surroundings – include everything else in the universe.

Section 10.3


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open

mass & energy

Exchange:

closed

Energy exchange

isolated

No exchange

Section 10.3


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Energy Changes Accompanying the Burning of a Match

Section 10.3


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• Endothermic Process: Heat flow is into a system. Absorb energy from the surroundings.

• Exothermic Process: Energy flows out of the system.

• Energy gained by the surroundings must be equal to the energy lost by the system.

Section 10.3


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Concept Check

Is the freezing of water an endothermic or exothermic process? Explain.

Section 10.3


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Concept Check

Classify each process as exothermic or endothermic. Explain. The system is underlined in each example.

a) Your hand gets cold when you touch ice.

b) The ice gets warmer when you touch it.

c) Water boils in a kettle being heated on a stove.

d) Water vapor condenses on a cold pipe.

e) Ice cream melts.

Exo

Endo

Endo

Exo

Endo

Section 10.3


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Concept Check

For each of the following, define a system and its surroundings and give the direction of energy transfer.

a) Methane is burning in a Bunsen burner in a laboratory.

b) Water drops, sitting on your skin after swimming, evaporate.

Section 10.3


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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.

Section 10.3


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H2 + O2 have higher potential energy than H2O

energy is given offenergy is absorbed

Electrolysis of Water Burning of Hydrogen in Air

higher potential energy lower potential energy

Section 10.4

Thermodynamics

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• Study of energy• Law of conservation of energy is often

called the first law of thermodynamics. The energy of the universe is constant.

Section 10.5

Measuring Energy Changes

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• 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 = 4.184 joules

1Cal (food) = 1000 calories

Section 10.5


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Convert 60.1 cal to joules.

Example

4.184 J60.1 cal = 251 J

1 cal

Section 10.5


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1. The amount of substance being heated (number of grams).

2. The temperature change (number of degrees).

3. 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.

Energy (Heat) Required to Change the Temperature of a Substance Depends On:

Section 10.5


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Specific Heat Capacities of Some Common Substances

Section 10.5


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• Energy (heat) required, Q = s × m × ΔT

Q = energy (heat) required (J)

s = specific heat capacity (J/°C·g)

m = mass (g)

ΔT = change in temperature (°C)

To Calculate the Energy Required for a Reaction:

Section 10.5


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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.

Example

Section 10.5


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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?

Example

Q =

Q = 4.184 J/g C 6.25 g 39.0 C 21.0 C

Q = 471 J

s m T

Section 10.5


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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) 0.052 g

b) 4.6 g

c) 19 g

d) 590 g

Q =

Q =

4.184 J142 cal

1 cal = = 19 g

0.45 J/g C 92 C 23 C

s m T

ms T

m

Section 10.5


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Twice as much heat energy is required to raise the temperature of 200 g of water 10oC as compared to 100 g of water.

200 g water

20oC

A

100 g water

20oC

B

100 g water

30oC

200 g water

30oC

heat beakers

4184 J 8368 J

temperaturerises 10oC

Section 10.5


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Concept Check

A 100.0 g sample of water at 90°C is added to a 100.0 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

Section 10.5


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Concept Check

A 100.0 g sample of water at 90.°C is added to a 500.0 g sample of water at 10.°C.



b) 50°C


Calculate the final temperature of the water.

23°C

Section 10.5


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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)



b) 50°C


Calculate the final temperature of the water.

18°C

Section 10.6

Thermochemistry (Enthalpy)

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Calorimetry

• Enthalpy, H is measured using a calorimeter.

Section 10.6


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An energy transformation occurswhenever a chemical change occurs.

• If energy is absorbed during a chemical change, the products will have more chemical potential energy than the reactants.

• If energy is given off in a chemical change, the products will have less chemical potential energy than the reactants.

Section 10.6


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A sample of a metal with a mass of 212 g is heated to 125.0oC and then dropped into 375 g of water at 240.0oC. If the final temperature of the water is 34.2oC, what is the specific heat of the metal?

When the metal enters the water, it begins to cool, losing heat to the water. At the same time, the temperature of the water rises. This process continues until the temperature of the metal and the temperature of the water are equal, at which point (34.2oC) no net flow of heat occurs.

Section 10.6


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• Calculate the heat gained by the water.• Calculate the final temperature of the metal.• Calculate the specific heat of the metal.

Section 10.6


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Δt = 34.2oC – 24.0oC = 10.2oCtemperature rise of the water

Heat Gained by the Water

o(10.2 C) = (375 )g o

4.184 Jg C

heat gained by the water

= 41.60 x 10 J

Section 10.6


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Δt = 125.0oC – 34.2oC = 90.8oCtemperature drop of the metal

Once the metal is dropped into the water, its temperature will drop until it reaches the same temperature as the water (34.2oC).

Heat Lost by the Metal

heat lost by the metal

heat gained by the water

= = 41.60 x 10 J

Section 10.6


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heatspecific heat =

mass x Δ t

4

o

1.60 x 10 J(212g)(90.8 C)

o

0.831 Jg C)

specific heatof the metal

=

The heat lost or gained by the system is given by:

(mass) (specific heat) (Δt) = energy change

rearrange

Section 10.6


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Homework

• Reading assignment– Pages 288 through 302 and chapter review

• Homework Questions and Problems: pages 317 - 319– 3, 5, 9, 13, 17, 19, 25, 27, 29, 31, 33, 35.

• Due on

chapter 10 energy. chapter 10 table of contents copyright © cengage learning. all rights reserved 2...

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