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Energy and Energy Changes
• Energy: ability to do work or produce heat
– Chemical, mechanical, thermal, electrical, radiant,
sound, nuclear
– Potential and kinetic
• Energy may affect matter.
– e.g. Raise its temperature, eventually causing a state
change, or cause a chemical change such as
decomposition
• All physical changes and chemical changes
involve energy changes.
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Energy and Energy Changes
• Potential Energy:
energy due to
composition or position
• Kinetic Energy: energy
due to motion
– - ½ mv2
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10 | 2
Energy and Energy Changes (cont.)
• Law of Conservation of Energy: energy can
be converted from one form to another, but
cannot be created or destroyed
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Work and Energy
• Work: force acting over a distance
– w=f•d
– Work done on a system will increase the
energy of the system, whereas work done by
the system will decrease the energy of the
system
• State function: a property that changes
independent of pathway
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Temperature and Heat
• Heat: a flow of energy due to a temperature
difference
• Temperature: a measure of the random
motions of the components of a substance
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Exothermic vs. Endothermic
• System: that part of the universe that we wish to
study
• Surroundings: everything else in the universe
• Exothermic process: a process that results in the
evolution of heat
- Example: when a match is struck, it is an
exothermic process because energy is produced as
heat.
• Endothermic process: absorbs heat
- Example: melting ice to form liquid water is an
endothermic process because the ice absorbs heat in
order to melt
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Exothermic Process
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Thermodynamics
• The Law of Conservation of Energy is also known
as The First Law of Thermodynamics. It can be
stated as “the energy of the universe is constant.”
• Internal Energy (E) = kinetic energy + potential
energy
• ΔE = q + w = change in internal energy
q = heat absorbed by the system
w = work done on the system
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Units of Energy
• One calorie = amount of energy needed to
raise the temperature of one gram of water by
1°C
– kcal = energy needed to raise the temperature of
1000 g of water 1°C
• joule
– 4.184 J = 1 cal
• In nutrition, calories are capitalized.
– 1 Cal = 1 kcal
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Example - Converting Calories to Joules
Convert 60.1 cal to
joules.
1 cal  4.184 joules
4.184 J
60.1cal 
 251J
1 cal
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Energy & Temperature of Matter
• The amount the temperature of an object increases
depends on the amount of heat added (q).
– If you double the added heat energy the temperature
will increase twice as much.
• The amount the temperature of an object increases
when heat is added depends on its mass
– If you double the mass it will take twice as much heat
energy to raise the temperature the same amount.
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Specific Heat Capacity
• Specific heat (s): the amount of energy
required to raise the temperature of one
gram of a substance by one degree Celsius
J
By definition , the specific heat of water is 4.184
g C
Amount of Heat = Specific Heat x Mass x Temperature Change
Q = s x m x T
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Specific Heat Capacity
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Example #1:
Calculate the amount of heat energy (in joules)
needed to raise the temperature of 7.40 g of
water from 29.0°C to 46.0°C.
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Example #1 (cont.)
Specific heat of water = 4.184
J
g°
C
Mass = 7.40 g
Temperature change = 46.0°C – 29.0°C = 17.0°C
Q = s • m • T
J
Heat  4.184
 7.40g 17.0C  526 J
g C
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Example #2
A 1.6 g sample of metal that appears to be
gold requires 5.8 J to raise the temperature
from 23°C to 41°C. Is the metal pure gold?
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Example #2
Q  s  m  T
Q
s
m  T
T  41C - 23C  18C
5.8 J
J
s
 0.20
1.6 g x 18C
g C
Table 10.1 lists the specific heat of gold as 0.13
Therefore the metal cannot be pure gold.
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J
g C
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Enthalpy
• Change in enthalpy (ΔHp = qp): the amount
of heat exchanged when heat exchange
occurs under conditions of constant pressure
• Enthalpy is a state function
• ΔH is independent of the path taken
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Hess’s Law
• Hess’s Law: in going from a set of reactants to
a set of products, the change in enthalpy is the
same whether the reaction takes place in one
step or in a series of steps.
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Hess’ Law (cont.)
• ΔHreaction= ∑Δhsteps
• If the direction of a reaction is reversed, the
sign of ΔH is reversed.
• ΔHforward = -Δhreverse
• Magnitude of ΔH α quantities of reactants
and products
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Hess’s Law (cont.)
• Overall reaction:
N2 + 2O2
2NO2 ΔH = 68 kJ
• This reaction can be carried out in 2 steps:
N2 + O2
2NO
ΔH = 180 kJ
2NO + O2
2NO2
ΔH = -112 kJ
-------------------------------------------------------N2 + 2O2
2NO2
ΔH = 68 Kj
Note: the sum of the two reactions gives the overall
reaction and the same is true for the sum of the
enthalpy change values.
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Calorimetry
• The amount of heat flow transferred during
a reaction is determined from temperature
measurements made in a calorimeter.
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Calorimetry (cont.)
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Energy Quality & Quantity
• While the total amount or quantity of energy
in the universe is constant (1st Law) the
quality of energy is degraded as it is used.
Burning of petroleum:
High grade concentrated energy
energy (heat)
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Low grade
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Fuels
• Petroleum
– A fossil fuel composed mainly of hydrocarbons
• Natural gas
– Consists largely of methane
– Also contains ethane, propane, and butane
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Fuels (cont.)
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Fuels (cont.)
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Global Warming
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Global Warming (cont.)
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Energy Use and Sources
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Energy as a Driving Force
• Most processes that occur spontaneously
involve an “energy spread.”
– Heat flows from high to low temperature and
“spreads”
…or a “matter spread”
– Salt dissolves or “spreads” in water
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Entropy
• Entropy (S) is a measure of disorder or
randomness.
– As a system becomes more disordered, ΔS >0
• Second Law of Thermodynamics: the
entropy of the universe is always increasing.
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