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Energy
11/12 & 11/13
Warm Up

Refer to the diagram below:
◦ What is going to happen in this situation?
Forms of Energy

Energy- ability to do work and produce
heat

W=FX
◦ Work= Force x Distance
Measurement of Energy

Measured by amount of work it can do
(physics)
- ORamount of heat it can be changed into
(chemistry)

SI unit of measure = Joule (J)
Forms of Energy
A. Mechanical Energy
◦
Energy that can exert a force or produce motion
 Potential: energy of position, stored energy

chemical potential energy – energy stored in a substance b/c of its
composition
 Kinetic: energy of motion
B. Chemical Energy
◦
Energy associated with chemical change
C. Electrical Energy
◦
Current carries energy to do work
D. Electromagnetic radiation
E. Magnetic
F. Nuclear
G. Heat- energy in the process of flowing from warmer to
cooler object
H. Sound
Conservation of Energy

Energy is changed, or converted, from one
form to another

Law of Conservation of Energy
◦ Energy cannot be created or destroyed
 AKA: 1st Law of Thermodynamics
◦ The total amount of energy remains the same
during all energy changes.

Temperature of all things flow towards
equilibrium
◦ It equals out.
◦ 2nd Law of thermodynamics

Absolute Zero=No movement=No
Energy
◦ 3rd Law of thermodynamics
Describe the energy conversion in
this picture
Thermochemistry
Thermochemistry is the study of heat
changes that accompany chemical
reactions and phase changes.
 In thermochemistry, the system is the
specific part of the universe that contains
the reaction or process you wish to study.

Everything in the universe other than the
system is considered the surroundings.
 Therefore, the universe is defined as the
system plus the surroundings.

universe = system + surroundings
Temperature vs. Heat

Heat
◦ Energy that transfers from one object to
another because of a temperature difference
between them.

Temperature
◦ A measure of the average kinetic energy of
the particles in a sample of matter
Heat and temperature
Heat is energy. It can do work.
 Temperature is a man-made, arbitrary scale indicating
which direction heat is flowing…is heat going into the
system, temperature rising or is heat leaving the system,
temperature declining.
 Heat is measured with an instrument called a
calorimeter.
 Heat is NOT measured with a thermometer.
 Temperature is measured with a thermometer.
 Heat is measured in Joules.
 Temperature is measured in degrees.

How to measure heat and energy…
◦ calorie
 Quantity of heat that will increase the temperature of 1
gram of water by 1oC
◦ 1 calorie (cal) = 4.19 Joules (J)
◦ 1 Calorie (food) = 1000 cal (heat)

How would you calculate food Calories to
Energy?
◦ How many joules of energy will a bowl of cereal
containing 230 Cal produce?
What are the Units of Energy?

It all depends…

In the food we eat, the units are Calories

In science, we use the calorie or the Joule

To convert between calories and joules, use
the following conversion factor
◦ 1 calorie = 4.19 Joules (J)
Example Problems

Covert 465 Joules to calories.

Convert 110 calories to Joules.
So when we talk about Calories (or calories)
we are talking about energy today and the
amount of energy we are taking in to our body.
 We must USE that energy we have taken in OR
our bodies will convert it to the storage form
of energy…FAT.
 Fat is simply the body’s way of saying…“don’t
want to use that energy now? OK. I will
save it for you for later”.

Specific Heat

When heated, different substances change
temperature at different rates.

Specific Heat
◦ The amount of heat it takes to raise the
temperature of 1 g of substance by 1oC.
 Specific heat of water = 1 cal/g.oC
 Specific heat of Iron = 0.11 cal/g.oC

What are the implications for this?

Where would you rather sit on a 115
degree day, a metal bench or a kiddie pool
full of water?
Calorimetry
Measurement of the amount of heat
released during a reaction
 Heat measured using a calorimeter (refer
to diagram)

◦ Calorimeter = device to measure the transfer
of heat to water
Simple Calorimeter
Test Tube
Thermometer
Stirring Rod
Water
Reaction
Sealed
Container
How to calculate heat…
Equation Q=mc∆T
 To calculate the calories of heat
transferred during a chemical reaction,
multiply:

◦ Q=heat absorbed or released
◦ Mass of substance in calorimeter (g)  m
◦ Change in temperature (oC)  T
 Tfinal-Tinitial
◦ Specific heat  This value is given to you
 1calorie/g.oC Or you can use 4.19 joules/g.oC
When working numerical problems we
will quickly become confused if we don’t
adopt a universal convention for when we
use a positive sign or a negative sign.
 Sign Convention for heat, Q
If Heat is transferred into the system Q
>0
+ absorbing heat, ENDOthermic
If Heat is transferred out of the system Q<0
- releasing heat, EXOthermic

Examples (cal)

2000 grams of water has its temperature
raised by 3.0 oC. How much heat was
produced? (1 cal/g.oC)
◦ 6000 cal

How many calories must be added to 5000 g
of water to change its temperature from
20oC to 30 oC?
◦ 50,000 cal

If 500 g of water at 25oC loses 2500 calories,
what will be the final temperature?
◦ 20oC
More Examples (joules)

The temperature of a sample of iron with
a mass of 10g changed from 50.4oC to
25oC with the release of 114J of energy.
What is the specific heat of iron?
◦ 0.449 J/goC

The temperature of a sample of water
increases from 20oC to 46.6oC as it
absorbs 5650 J of heat. What is the mass
of the sample? (4.19J/goC)
◦ 50.7 g

If 335 g water at 65.5 oC loses 9750 J of
heat, what is the final temperature of the
water?
◦ 58.6oC
Temperature Conversions

Celsius Scale

Kelvin Scale
◦ Freezing point of water = 0 oC
◦ Boiling point of water = 100 oCelsius
◦ Interval between them is divided into 100 parts
◦ Absolute zero=NO movement
◦
◦
◦
◦

 Lowest temperature theoretically possible
Absolute zero = 0 K = -273 oC
Freezing point of water = 273 K
Boiling point of water = 373 K
Size of degree is same as Celsius
Fahrenheit Scale
◦ Freezing point of water = 32 F
◦ Boiling point of water = 212 F

Converting
◦
◦
◦
◦
Kelvin = Celsius + 273 (K = oC + 273)
Celsius = Kelvin – 273 (oC = K -273)
Fahrenheit = (oC × 9/5) + 32
Celsius = (F-32)(5/9)
Examples

Convert from K  oC
◦
◦
◦
◦

110K
476K
295K
1114K
Convert from oC K
◦ 11 oC
◦ 112 oC
◦ -15 oC

Covert from oCF
◦
◦
◦
◦

-111 oC
0 oC
45 oC
323 oC
Convert from F oC
◦ 45 F
◦0F
◦ 115 F
Kinetic Theory
Matter has small particles in continuous
motion
 The faster a particle moves, the greater the
kinetic energy
 Temperature

◦ Measure of the average kinetic energy of particles
in the sample.
 At absolute zero, the average kinetic energy is zero
 Higher temperatures have a greater average kinetic
energy
 Samples at the same temperature have the same average
kinetic energy
What kind of energy is this?

Figure 3.12: Equal masses of hot water
and cold water separated by a thin metal
wall in an insulated box.

The H2O molecules in hot water have
much greater random motions than the
H2O molecules in cold water.

The water samples now have the same
temperature (50°C) and have the same
random motions
Heat Transfer
Transfer of heat into or out of a
sample
◦ Heat transferred into a sample can be used
to increased the average kinetic energy of
the particles
 This causes an increase in temperature
 When a sample cools, the particles lose
kinetic energy. Heat is given off.
Heat Transfer
◦ Heat can also enter or leave a sample
without causing a change in
temperature.
 During the process of ice melting, the
heat absorbed is used to rearrange the
particles, not to increase the kinetic
energy of the particles
 No change in temperature occurs
Phases Changes that require energy
What happens to molecules in a solid as
it melts?
 Melting

◦ The amount of energy (heat of fusion)
required to melt one mole of a solid depends
on the strength of the forces keeping the
particles together (Intermolecular force).
When liquid water is heated, some
molecules escape from the liquid and
enter the gas phase.
 If a substance is usually a liquid at room
temperature (as water is), the gas phase is
called a vapor.

Vaporization
Vaporization is the process by which a
liquid changes into a gas or vapor.
 As temperature increases, water
molecules gain kinetic energy

◦ At Boiling point, molecules throughout the
liquid have the energy to enter the gas or
vapor phase.
◦ The amount of energy required to do this is
the heat of vaporization
Sublimation

The process by which a solid changes
directly into a gas without first becoming
a liquid is called Sublimation.
◦ Solid air fresheners and dry ice are examples
of solids that sublime.
Condensation
Some phase changes release energy into
their surroundings.
 For example, when a vapor loses energy,
it may change into a liquid.
 Condensation is the process by which a
gas or vapor becomes a liquid. It is the
reverse of vaporization.

Water vapor undergoes condensation when
its molecules lose energy, their velocity
decreases.
 The freezing point is the temperature at
which a liquid becomes a crystalline solid.
 When a substance changes from a gas or
vapor directly into a solid without first
becoming a liquid, the process is called
deposition.

◦ Deposition is the reverse of sublimation. Frost is
an example of water deposition.
Phase changes
Temperature
Heating Curve
Energy or Time
Practice
Substance Freezing point (oC)
Boiling Point
(oC)
Water
0.0
100.0
Gallium
23.0
89.0
Iron
723.0
2780.0
 At room temperature (27 oC), Iron is a solid,
mixture, liquid or gas?
 At 800 oC, Iron is a solid, mixture, liquid or gas?
 During the process of heating water from 27 to
85 oC :
◦ Did the potential energy change? Kinetic energy?
◦ Is it an endothermic or exothermic reaction?
Entropy
Entropy (S) is a measure of the disorder or
randomness of the particles that make up a
system.
 Spontaneous processes always result in an
increase in the entropy of the universe.
 Several factors affect the change in entropy of a
system.

◦
◦
◦
◦
Changes of state
Dissolving of a gas in a solvent
Change in the number of gaseous particles
Dissolving of a solid or liquid to form a
solution
◦ Change in temperature
Overview of Conduction,
Convection & Radiation

http://www.wisconline.com/objects/index_tj.asp?objid=SC
E304
Methods of Heat Transfer

Conduction:
◦ Transfer of heat between substances that are in direct
contact with each other
 Occurs mainly in solid
 Better conductor  More rapid heat transfer
 Examples of good and poor conductors?

Convection:
◦ Up and down movement (circulation) of gases and
liquids caused by heat transfer
 Does not occur in solid (molecules not free to
move around)
 Examples of convection?

Radiation:
◦ Electromagnetic waves traveling through space
 Does not require a medium to transfer heat
◦ Waves transfer heat to the object
 Examples of radiation heat transfer?
Practice




Boiling water over a campfire
Melting a tub of ice cream on the kitchen counter
Electric Stove versus Gas Stove
◦ Which stove will boil water faster? Why?
Why is the second floor usually warmer than the
first floor? Why?
Unit 4 Test

Nuclear Reactions
◦
◦
◦
◦

Alpha, Beta, Gamma, Electron Capture, Positron
Balancing Nuclear Equations
Half Life
Radioactive Decay
Energy
◦ Definitions
◦ Calculations
 Energy
 Temperature Conversions

Phase Changes
◦ Diagram
◦ Understanding how KE & PE relate to the
diagram
◦ Definitions

Conduction, Convection, & Radiation
◦ Know examples