Download is energy

Energy and
Transformation
Aim 7– What are some of the different
types of energy we encounter?
Energy in the US
• Energy is the capacity
for doing work
• It exists in many forms,
and sources of energy in
the US come from
specific substances
and technologies
• The chart at the right
represents a breakdown
of energy sources in the
United States as of 2010
• What percentage of the
total resources are
finite (will run out)?
• Energy may exist in various forms
• Kinetic energy
– Energy associated with motion
– Moving objects have kinetic energy
– There are different subforms of
kinetic energy…
• Mechanical energy
– Moving machine parts
transfers kinetic energy
within the machine
– Example: a car’s
engine and drive train
• Electrical Energy
– Moving electrons
carry energy through
electrical devices
– Example: power
running through an
iPhone or light bulb
• Electromagnetic radiation
– Energy carried by photons
• the elementary particles of radio waves,
light, X-rays, gamma rays, etc.
– Longer the wavelength
– the lower the energy
• Potential energy
– Energy associated with position
– Often referred to as stored energy
– There are different subforms of potential
energy
• Gravitational
Potential Energy
– Energy stored in
an object due to
its position
– Example:
– A rollercoaster
• Chemical Energy
– Energy stored within a
chemical substance
– released when a
substance goes through
a chemical reaction
– Examples:
– energy stored in:
• a candle’s wax
• the chemical bonds of
dynamite
• the chemical
reactions of a battery
• Nuclear Energy
– Energy stored within
the nucleus of an
atom
• Examples:
• Nuclear weapons
– split atoms in
uncontrolled
reactions
• Fission power plants
– split atoms in
controlled reactions
Examples of Energy Transformations
State what energies are being converted:
• Coal is burned to run a generator
chemical energy to electrical energy
• A rock is dropped off a building and hits a car
gravitational potential energy to kinetic energy
• A cell phone is plugged in to a socket and charged
electrical energy to chemical energy
• A cheeseburger is digested and gives a runner energy
chemical energy to mechanical (muscles/bones)
• Light waves allow a plant to do photosynthesis
electromagnetic energy to chemical energy (sugar)
• Mr. Foley set off a nuclear bomb in Hauppauge
nuclear energy to heat and light energy
Aim 8 – What is heat flow?
• Thermal energy
– Thermal energy is a result of the
kinetic energy of the molecules’ motion
– The total kinetic energy of all the molecules
combined is called thermal energy
– The amount of thermal energy contained in the
particles depends on
• how fast they are moving and
• how many particles there are
– Thermal energy may be transferred from one
system to another
• Heat
– is a transfer of thermal energy
– flows from a body of higher temperature to a body
of lower temperature
• Units of Heat flow
• Joule (J)
– a base unit of
metric heat
measure
– 4.18 joules heat
will raise one
gram of water 1oC
• calorie (cal)
– another base unit of heat measure
– 1 calorie of heat will raise one gram of water 1oC
• Kilocalorie (kcal)
– Is equal to 1,000 calories
– 1 kilocalorie is equal to 1 nutritional Calorie(the
unit used to measure food intake)
• Thermal energy is NOT the same as
temperature
• Temperature is NOT a form of energy
• Temperature
– (def) a measure of the average kinetic energy
of the particles in an object and a system
– Measuring temperature
allows us to calculate:
• the heat in a system
or object;
• The heat being
transferred
between systems
• Temperature Scales
• Three scales are used
to measure temperature
– Fahrenheit (rarely
used in chem) and
Celsius are based
on the freezing and
boiling points of water
– Kelvin is based on
absolute zero, the
absence of all particle
motion
– therefore, to convert from Kelvin to Celsius we
use:
Kelvin = oC +273
Practice Problems
Convert each of the following using the
formula K = oC + 273
a) 273 K
=
_________ oC
b) 373 oC
=
_________ K
c) -100 K
=
_________ oC
d) 100 oC
=
_________ K
e)
5K
=
_________ oC
f) 38 oC
=
_________ K
• Energy Conversions
and the Law of Science
– Theories that have been proven over
and over again to be valid
– Examples:
– Law of Gravity – all objects exert a gravitational
force in proportion to the sized of the object
– Kepler’s Law of Planetary Motion – all planetary
bodies move in an elliptical orbit
• Chemistry has laws as well!
• The first three are:
– The Law of Conservation of Energy
– The Law of Conservation of Matter
– The Law of Conservation of Energy and Matter
• The Law of Conservation of Energy
– states that the total energy of an isolated
system cannot change
– it is said to be conserved over time
– Energy can be neither created nor destroyed,
but can change form
– Example:
– chemical energy
can be converted
to kinetic energy
in the explosion
of a stick of
dynamite
• The Law of Conservation of Matter
– States that in chemical reactions, the
total mass of the reactants must equal
the total mass of the reactants.
– Example 2 X + Y  2 X2Y
• In this reaction, 2 particles of X react with 1
particle of Y to make 2 particles of X2Y
• Nothing is lost when the reaction occurs
• it is simply rearranged
– Matter cannot be neither created or destroyed,
but may be changed from one form to another
The Law of Conservation of
Matter and Energy
• states that neither matter or
energy can neither be created
nor destroyed, but may be
converted into the other.
• This was explained in part by
a very short but extremely
famous formula, created
by a gentleman by the name of Albert Einstein
• E = mc2, where “E” is energy, “m” is the mass
equivalent, and “c2 “ is the speed of light squared
J Deutsch 2003
19
Phase Changes
Aim 9 – What
happens when
temperatures
increase?
Phase Changes and Energy
• Changing energy causes changes in
phase as attractive forces are broken or
formed
• These energy changes and temperature
changes can be represented on a graph
• This graph is
called a phase
change diagram
Phase Changes and Energy
• Particles with temps above 0 Kelvin are
moving
• Particles have
– kinetic energy as they move in place (solids)
or away from each other (liquids and gases)
– potential energy due to attractive forces that
hold solids and liquids together
• As temperature increases
– The motion of the particles increases
– Therefore kinetic energy increases
– But the attractive forces are not being broken
– So no phase change occurs
Phase Changes and Energy
• Eventually, enough energy is added to
begin breaking the attractive forces
• Temperature stops increasing
– Therefore kinetic energy stops going up
• Eventually, the attractive forces holding the
particles to each other in the solid and liquid
phases break
– Potential energy increases
– But particles don’t move any faster
– So the average kinetic energy doesn’t
increase
• At the Melting Point
– the temperature a solid reaches the
point where attractive forces break,
– melting it and forming a liquid state
• After the substance melts completely
– The temperature continues to go up till it
reaches the
• Boiling Point
– the temperature a liquid reaches where all
the remaining attractive forces are broken
– The liquid changes over to the gas state
Phase Diagrams - Heating Curves
• As heat is added to a system
• What types of changes occur?
– Particles move faster
– Average kinetic energy increases
– Temperature increases
– Forces of attraction are broken as
• solids change to liquids
• liquids change to gases
• Because energy is added to break these forces of
attraction
– Stored energy increases
– So potential energy increases
Heating Curve of Water
Phase Changes
Aim 10 – What
happens when
temperatures cool?
Phase Diagrams – Cooling Curve
• As heat is REMOVED from a system
• What types of changes occur?
– Particles start to slow down
– Average kinetic energy decreases
– Temperature goes down
• Forces of attraction are reformed as
– gases change to liquids
– liquids change to solids
• Forming attractions lowers energy
– Stored energy decreases
– So potential energy decreases
Phase Diagram - Cooling Curve
Energy changes during
Phase changes
• Enthalpy – refers to heat being either lost
or gained from any system
– Exothermic changes
• energy is lost from a system
• Example – a burning match
– Endothermic changes
• energy is gained from a system
• Example – a cold pack
•
Entropy
– Refers to the amount of disorder
particles have in a sample
– The more disordered, the higher the entropy
• Solids – most organized, lowest entropy
• Gases – least organized, greatest
entropy
– As substances go thru phase changes, their
entropy will either increase or decrease
• Question – you melt ice… what happens to
the ice’s entropy?
Phases of Matter Changes:
Phase change
Change
occurring
(example)
Fusion
(melting)
Ice(s) 
water (l)
Endothermic
Increase – less
organized
Solidification
(freezing)
Water (l) 
Ice (s)
Exothermic
Decrease – more
organized
Vaporization
(boiling)
Water (l) 
Steam (g)
Endothermic
Increase – less
organized
Condensation
Steam (g) 
Water (l)
Exothermic
Decrease – more
organized
Sublimation
Dry Ice (s) 
Vapor (g)
Endothermic
Increase – less
organized
Deposition
Iodine (g) 
Iodine (s)
Exothermic
Decrease – more
organized
J Deutsch 2003
Exothermic or
endothermic?
Increase
or decrease in
enthalpy?
32
• Label the diagram with the following terms:
boiling
melting
PE h
KE h
Solid only
Liquid only
Gas only
Aim 11 – How
much heat is
absorbed or
released during
phase changes?
Calculating the Energy Changes
during Phase Changes
• When a phase change occurs, heat is either
– gained (heating occurs)
– or lost (cooling occurs)
• In order to calculate the heat change in the
sample or system, you need:
– the mass of the sample undergoing the phase
change
– and the change in heat constant for that
particular substance during a particular phase
change
• A constant is a value that is the same for
every situation
• There are two constants we use for
phase changes
– Heat of fusion – the amount of energy
gained or lost when a substance either melts or
solidifies (freezes)
– Heat of vaporization – the amount of energy
gained or lost when a substance either vaporizes
(boils or evaporates) or condenses
• For example, for water
– Heat of fusion (Hfusion ) = 334 joules/gram
– Heat of vaporization (Hvap ) = 2,260 joules/gram
– Joule – a unit of heat energy, similar to calories
• Knowing the mass of the sample,
• and knowing what type of phase
change is occurring,
• we can calculate the amount of heat (q)
gained or lost by the system using the
formulas on Table T in the CRT
• Example: a 100.0 gram sample of ice is
heated at 0oC and completely melted. How
much heat (q) is used?
q = mass of the sample x Hfusion
q = 100.0 grams x 334 J/g
q = 33,400 joules of heat energy
• Example 2: a student is distilling a
100.0 gram sample of water at 100oC.
How much heat must be added to
completely boil away the water?
q = mass of the sample x Hvaporization
q = 100.0 grams x 2260 J/g
q = 226,000 joules of heat energy
• Note – this is a huge number – how many
kilojoules is this sample?
226,000 J x 1 kJ = 226 kJ
1,000 J
1. How much heat energy is required to freeze
5.00 grams of ice?
q = m x Hfusion
q = 5.00 grams x -334 J/g
q = 1,670 J
2. How much heat energy is required to boil 20.00 g of
water?
q = m x Hvaporization
q = 20.00 x 2260 J/g
q = 45,200 J
3. How much ice will melt if 668,000 joules of energy
are added to a block of ice?
q = m x Hfusion
668,000 J = m x 334 J/g
m = 2,000 g
Aim 12 – how do
we calculate the
heat change in
my cup of tea?
• When a phase changes occur
– Heat is either gained or lost and the temperature
does not change
– Particles are not moving faster or slower
– Instead, attractive forces are being broken or formed
– Potential energy changes, but not average kinetic energy
• When we are heating a solid, a liquid, or a gas
and NOT causing a phase change
– Heat is either gained or lost and the temperature IS
increasing or decreasing
– The average kinetic energy is increasing or decreasing
– Potential energy is staying the same (no attractive forces
formed or broken)
– Heat is transferred to different materials at different rates
• Specific heat capacity (C)
– the types of materials determine the rate at which heat will
be absorbed.
– Each material has a specific heat capacity constant (C)
• Examples of specific heats
– water
= 4.18 J/g.oC
– aluminum = 0.91 J/g.oC
– iron
= 0.45 J/g.oC
– copper
= 0.39 J/g.oC
– J/g.oC is the number of joules of heat energy
needed per gram for a 1oC temperature change
• This is why when you touch the desk top,
and the metal chair leg, they feel different
temperatures!
• Demo – Ice Melting Blocks
• When temperature changes, the heat
gained or lost (q) can be calculated if
three things are known:
 The mass (m) of the sample
 The temperature change (T) of the
sample
(or, Tfinal – Tinitial)
 The specific heat capacity (C) of the
sample
• From Table T in the CRTs
q = m x T x C
1. A 100.0 gram sample of water has its
temperature changed from 10.0oC to
30.0oC. If the specific heat of water
(C) is 4.18 J/g.oC, what is the heat
energy (q)?
q = m x T x C
q = 100.0 g x (30.0 – 10.0)oC x 4.18 J/g.oC
q = 100.0 g x 20.0oC x 4.18 J/g.oC
q = 8,360 J
Example 2 – A 20.0 gram sample of water cools
from 40oC to 20oC. If the specific heat of water (C)
is 4.18 J/g.oC, what is the heat energy (q) LOST by
the water? How can you tell it was lost?
q = - 1,672 J
Example 3 – A 10.0 gram sample of aluminum has its
temperature changed from 25oC to 30oC. If the specific heat
of water (C) is aluminum is 0.91 J/g.oC, what is the heat
energy (q) gained by the metal?
q = 45.5 J
Example – Mr. Foley’s cup of Earl Grey tea holds 200.0 g of
tea (specific heat = 4.18 J/g.oC). First made, it has a
temperature of 90oC. By lunch time, it is 30oC. What is the
heat energy (q) lost by the tea?
q = - 50,160 J
• Calorimetry
– is used to determine
the heat content or
the heat transfer in a
system
– the amount of
temperature change
that occurs in the water
in the cup allows us to
measure the heat
– note – it is a
CALCULATED value
– At right – a simple
calorimeter
• Calorimeter
– is a heat measuring device.
– is generally used to measure the
amount of heat energy
– It then uses that to calculate the specific heat
of a substance or other heat related
information
Example 4 – A macadamia nut
burns in a calorimeter set up. A
50.0 gram sample of water has its
temperature changed from 20oC
to 30oC. If the specific heat of
water (C) is 4.18 J/g.oC, what is
the heat energy (q) gained by the
water? How much heat was
released by the macadamia nut?
q = m x T x C
q = 50.0 g x (30.0 – 20.0)oC x 4.18 J/g.oC
q = 50.0 g x 20.0oC x 4.18 J/g.oC
q = 4,180 J gained by the water, the
same lost by the nut