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What is energy?
 The ability to do work
 The ability to transfer heat
Two types: Potential and Kinetic
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Potential Energy
 Stored Energy
 Energy due to
 Position
 Chemical Bonds
 Nuclear
 Position: boulder at the top of the hill
 Chemical: tank of gas, hamburger
 Nuclear: atomic bomb, nuclear reactor
Kinetic Energy
 Active Energy
 Energy of Motion
 Electromagnetic waves, (ex. Light, Xrays)
 Heat
 Motion
 Electrical current,
 A moving truck has the ability to flatten you - do work
on you!
Kinetic and Potential Energy
Electromagnetic Spectrum
First Law of Thermodynamics
 Energy can
neither be
created nor
destroyed, but
may change from
one form to
another.
Page 2 bottom
Entropy –
nd
2
Law
 Entropy is the amount of disorder in a system
 Entropy always increases over time (in the absence of an
input of outside energy)
 Example: cleaning up your room
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Exothermic vs Endothermic
Chemical OR Physical changes can
be exothermic or endothermic.
Endothermic
Exothermic
Definition
Stored energy
Energy of motion
Energy is
Absorbed
Released
Temperature
+∆H
Chemical -∆T
Physical +∆T
-∆H
Chemical +∆T
Physical -∆T
2H2O + energy  2H2 + O2
Energy is on the left
2H2 + O2  2H2O + energy
Energy is on the right
Type of Energy Conversion
Example
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Regents Question: 06/02 #64-66
A hot pack contains chemicals that can be activated to produce heat. A cold
pack contains chemicals that feel cold when activated.
1.
Based on energy flow, state the type of chemical change that occurs in a hot
pack.
Exothermic
2.
A cold pack is placed on an injured leg. Indicate the direction of the flow of
energy between the leg and the cold pack.
From the leg to the cold pack (Hot to Cold)
3.
What is the Law of Conservation of Energy? Describe how the Law of
Conservation of Energy applies to the chemical reaction that occurs in the hot
pack.
Energy cannot be created nor destroyed. It can only be changed from one
form to another. The heat produced in the hot pack was stored in the
chemical bonds.
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Measuring Energy
I.
Energy
A. There are two units which are commonly used:
B. Calories (c):
1. amount of energy it takes to raise one gram of water
one degree Celsius
C. Joules (J):
1. 4.18 Joules = 1 calorie
2. Metric system is most commonly used in chemistry
Criteria
Heat
Temperature
Similar/Different
Energy
Kinetic
Motion of Molecules
Both are about
Motion
Quantitative Aspect
How fast molecules
are moving.
Measured by
temperature
Kelvin and Celsius
A form of Kinetic
Energy that involves
movement of
molecules
The measurement of
how fast a molecule
is moving
Definition
Examples
Celsius based on
properties of water
Kelvin based on
Celsius
Temperature is a
measurement of
Kinetic Energy
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III. Converting between Celsius and Kelvin
A. Reference Table
B. Why is it not out already?
1.
Temperature
Towards bottom
K = ◦C + 273
Heat and Temperature
 Temperature measures the average speed of the atoms
 Heat is the amount of kinetic energy of the atoms
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To convert
between Kelvin
and Celsius use
◦
K= C+273
J Deutsch 2003
21
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Phases of Matter- Page 9
Ice
Ice
Ice
Regular
Irregular
Irregular
Regular
Regular
Irregular
Minimal
Moderate
Fast
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Graph page 11 onto page 13.
Be sure to have an appropriate
scale. Circle the points and
connect them.
Heating Curve
180
F
160
140
120
D
E
100
80
60
B
C
40
20
A
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
A change in phase is a change in Potential
Energy, not Kinetic Energy
Boiling
Point
Potential energy changes, so
temperature doesn’t
Melting
Point
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Energy and phase changes
 AB
 solid warms up (KE/PE constant)
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Energy and phase changes
 AB
 solid warms up (KE/PE constant)
 BC
 solid melts (KE constant/PE)
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Energy and phase changes
 AB
 solid warms up (KE/PE constant)
 BC
 solid melts (KE constant/PE)
 CD
 liquid warms up (KE/PE constant)
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Energy and phase changes
 AB
 solid warms up (KE/PE constant)
 BC
 solid melts (KE constant/PE)
 CD
 liquid warms up (KE/PE constant)
 DE
 liquid boils (KE constant/PE)
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Energy and phase changes
AB
 solid warms up (KE/PE constant)
BC
 solid melts (KE constant/PE)
CD
 liquid warms up (KE/PE constant)
DE
 liquid boils (KE constant/PE)
EF
 gas warms (KE/PE constant)
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J Deutsch 2003
Regents Question: 06/02 #28
As ice melts at standard pressure, its temperature remains at 0°C until
it has completely melted. Its potential energy
(1) decreases
(2) increases
(3) remains the same
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J Deutsch 2003
Regents Question: 08/02 #54
A sample of water is heated from a liquid at 40°C to a gas at 110°C. The
graph of the heating curve is shown in your answer booklet.
a On the heating curve diagram provided in your answer booklet, label
each of the following regions:
Liquid, only
Gas, only
Phase change
Phase change
Gas Only
Liquid Only
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J Deutsch 2003
Regents Question: cont’d
b For section QR of the graph, state what is happening to the water
molecules as heat is added.
They move faster, their
temperature increases.
c For section RS of the graph, state what is happening to the water
Their intermolecular bonds
molecules as heat is added.
are breaking, their potential
energy is increasing.
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J Deutsch 2003
Regents Question: 01/02 #47
What is the melting point of this substance?
(1) 30°C
(3) 90°C
(2) 55°C
(4) 120°C
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Graph page 14 onto page 16.
Be sure to have an appropriate
scale. Circle the points and
connect them.
A
B
C
D
E
F
Energy and phase changes
AB
 Gas cools down (KE/PE constant)
BC
 Gas condenses (KE constant/PE )
CD
 liquid cools down (KE  /PE constant)
DE
 liquid freezes (KE constant/PE )
EF
 Solid cools down (KE  /PE constant)
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Pages 17-18
Page 19
How do we calculate amount
of heat,(Q), if we are not given
a graphic?
3 equations for Q
Q = mCT
Q = mHf
Q = mHv
 Have to figure out which one to use for
a given problem.
 Depends which section of heating
curve.
 Look for hints in the problem.
Calculating Heat Transferred
Simple system: Pure substance in a
single phase. To calculate heat gained
or lost, use:
Q = mCT
•Q = amount of heat transferred
•m = mass of substance
•C = specific heat capacity of the substance (Table
B).
•T = temperature change = Tfinal – Tinitial
Q = mCT
 Temperature
changed
 Temperature
increased
 Temperature
decreased
 Initial / Start
temperature
 Final temperature
 Ending temperature
 From ____ to ____
 Water
Amount of energy required to convert 1 gram of a
pure substance from the solid to the liquid phase at
the melting point.
Heat of Fusion
Q = mHf
Use this equation to calculate energy
changes for phase changes between ice
& liquid water at 0C.
Q = mHf
 Ice
 Freezing
 Melting
 At 0C (for H2O)
 At constant temperature
Amount of energy required to convert 1 gram of a
pure substance from the liquid to the gas phase at
the boiling point.
Heat of Vaporization
Q = mHv
Use this equation to calculate energy
changes for phase changes between
steam & liquid water at 100C.
Q = mHv
 Steam
 Boiling
 Condensation
 At 100C (for H2O)
 At constant temperature