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Zumdahl • Zumdahl • DeCoste
World of
CHEMISTRY
Chapter 10
Energy
Goals of Chapter 10
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General properties of energy
Temperature & Heat
Direction of energy flow as heat
How energy flow affects internal energy
How to measure heat
Heat (enthalpy) of chemical reactions
Hess’s Law
Changes in quality of energy as it’s used
World’s energy resources
Energy as driving force for natural processes
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10-3
The importance of energy
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Huge abundance of fossil fuels
Society with huge appetite for energy
We have become dependent on oil
Has led to tension between countries
Supplies are dwindling, prices are rising
Need to find alternatives to oil
Use relationship between chemistry & energy
to find these alternatives
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10-4
Energy: the ability to do work or produce heat
• Potential Energy
• Energy due to position or composition
• Examples: water behind a dam, gasoline
• Kinetic Energy
• Energy of motion – depends on mass &
velocity (KE = ½mv2)
• Examples: car driving, thrown baseball
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10-5
Law of Conservation of Energy
• Energy can be converted from one form to
another but cannot be created or destroyed
• Energy of the universe is constant
• Can convert from one form to another
• Example: roller coasters
• State function: property that is independent of
pathway
• Example: Ball bearing in roller coaster; PE is
same at top of first hill regardless of path to bottom
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10-6
Temperature & Heat
• Temperature: measure of the random
motions of the components of a
substance (warm water molecules move
faster than cold water molecules)
• Heat: flow of energy due to a
temperature difference (heat will flow
from warmer water to cooler water)
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10-7
Figure 10.2: Equal masses of hot and cold water.
Each side
contains
1 kg of
water
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10-8
Figure 10.3: H2O molecules in hot and cold water.
Water
molecules
move
more
rapidly in
hot water
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10-9
Figure 10.4: H2O molecules in same temperature
water.
Heat is
transferred
from the hot
water to the
cold water
until both are
the same
temperature
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10-10
Final temperature is the average of the two
original temperatures
Change in temp (hot)
ΔT = 90°C – 50°C = 40°C
Change in temp (cold)
ΔT = 50°C – 10°C = 40°C
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10-11
Exothermic vs. Endothermic
• System: part of universe on which we
wish to focus
• Surroundings: everything else in
universe except system
• Exothermic: evolution of heat, energy
flows out of system
• Endothermic: absorbs energy from
surroundings, heat flows into system
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10-12
Which is it - Endothermic or Exothermic?
1. Your hand gets cold when you touch ice.
(with respect to your hand)
2. The ice melts when you touch it
3. Ice cream melts
4. Propane burning in a propane torch
5. After swimming water drops evaporate from
your skin
6. Two chemicals mixing in a beaker give off
heat
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10-13
What happens in an exothermic chemical reaction?
• Energy is conserved
• Reactants have potential energy
• Energy gained by surroundings must equal
energy lost by the system
• In any exothermic reaction, some of the
potential energy stored in the chemical
bonds is converted to thermal energy
via heat (random kinetic energy)
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10-14
Figure 10.5: The energy changes
accompanying the burning of a match.
Reactants have > PE than products, difference is heat
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10-15
Thermodynamics
• Thermodynamics – the study of energy
• First Law of Thermodynamics = Law of
Conservation of Energy (energy of the
universe is constant)
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10-16
Internal Energy (E)
• Sum of kinetic and potential energies of all
“particles” in a system
• Can be changed by flow of work, heat, or both
• ΔE = q + w; change in energy = heat + work
• Sign reflects systems point of view
• Endothermic – energy flows into system = +q (energy is
increasing)
• Exothermic – energy flows out of system = -q (energy is
decreasing)
• Same rules apply to work (w)
• +w = work flows into system (surroundings do work)
• -w = work flows out of system (system does work)
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10-17
Measuring Energy Changes
• Calorie: amount of energy (heat)
required to raise the temperature of one
gram of water by one degree Celsius
• Food “calorie” = kilocalorie = 1000 cal
• Joule: SI unit of energy
• 1 calorie = 4.184 J
• It takes 4.184 J of energy to raise the
temperature of one gram of water by 1°C
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10-18
How much energy is required to raise the
temperature of 7.40 g of H2O from 29.0°C to
46.0°C?
E = specific heat x mass x temp change
E = (4.184 J/g°C)x(7.40 g)x(46°C – 29°C)
E = 526 J
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10-19
Thermochemistry (Enthalpy)
• H = Enthalpy: energy function that indicates
how much energy is produced or absorbed in
a reaction
• ΔHp = energy that flows as heat
• ΔH: the change in enthalpy
• p: indicates process has occurred under constant
pressure
• The enthalpy change is the same as the heat of
reaction
• See Example 10.5 & Self-Check 10.5 (pg. 302)
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10-20
Figure 10.6: A coffee-cup calorimeter.
Calorimeter: device used to
determine the heat associated
with a chemical reaction
Run reaction – observe
temperature change
Heat capacity of calorimeter
enables us to calculate the
heat energy released/absorbed
by reaction
Determine ΔH for reaction &
calculate ΔH for other
reactions
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10-21
Hess’s Law
• Enthalpy is a state function – it is
independent of the pathway
• When 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 a series of steps
• See example on page 304
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10-22
Two characteristics of ΔH for a reaction:
• 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 also multiplied by that integer
• See examples on page 304 & 305
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10-23
Figure 10.7: Energy sources used in the
United States.
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Figure 10.8: The earth’s atmosphere.
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Figure 10.9: The atmospheric CO2 concentration over
the past 1000 years.
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10-26
Energy as a Driving Force
• Two driving forces
• Energy spread: concentrated energy is
dispersed widely
• Matter spread: molecules of a substance
are spread out and occupy a larger volume
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10-27
Entropy
• Designated by letter S
• Measure of disorder or randomness
• More disorder (or entropy) means:
• More energy spread
• More matter spread
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10-28
Figure 10.10: Comparing the entropies of ice and
steam.
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10-29
Second Law of Thermodynamics
• The entropy of the universe is always
increasing
• All processes lead to a net increase in
the disorder of the universe
• We are plunging slowly toward total
randomness – the heat death of the
universe
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