Chapter 2
ENERGY &
MATTER
Wednesday, 10/1/14
Learning Target:
Know the 3 basic forms of energy and how
energy is calculated.
Learning Outcome:
I will complete energy conversion problems.
What is Energy?
 The capacity to do work or produce heat.
Law of Conservation of Energy
 Energy can neither be created nor
destroyed in any chemical or
physical process. It can be
converted from one form to
another.
https://www.youtube.com/watch?v=qybUFnY7Y8w
Greenhouse Effect
2-1 Energy
•
Energy is classified into three main forms
•
Radiant
•
Kinetic &
•
Potential
Radiant Energy
This is energy from the Sun
which is the result of nuclear
fusion
Kinetic Energy
This is the energy carried by objects in motion,
like a locomotive.
Kinetic Energy includes:
1. Mechanical energy carried by the
moving parts of a machine
2. Thermal Energy of the random
internal motion of particles in all substances
Kinetic Energy (KE)
 KE = ½ mv2
 KE = kinetic energy
Unit
J = Joule (kg.m2/s2)
 m = mass
kg
 v = velocity
m/sec
Calculate the KE of a 70.kg
person walking at 2.5m/s.
Potential Energy
This is the energy possessed by objects because
of the position or the arrangement of their
particles
In essence it is stored energy.
Types of Gravitational Potential Energy
1.
2.
3.
Gravitational
Electrical – different electrical charges
Chemical – Fuels and Food
Energy stored in food is often
given a unit that is related to
the calorie. The Calorie (Cal) is
1000cal or 1 kilocalorie.
The SI Unit of energy is the
Joule (J)
Joule in the long form is kg.m2/s2
4.184J = 1 cal
4.184 KJoule = 1Kcal or 1Cal
1 KJ = 1000 J
1 Cal = 1000 cal
Calorie (cal) [older unit]
 The amount of energy required to raise the
temperature of 1 gram of water by 1 degree
Celsius.
1cal =4.184 joules
chocolate bar=200 Cal
200Cal x 4.184 KJ/Cal=
Energy in one chocolate bar= 836.8KJ
Thursday, 10/2/14
Learning Target:
Know what temperature is the measure of and
how to convert between Kelvin, Celsius and Fahrenheit
temperature scales.
Learning Outcome:
I will complete temperature conversion
problems.
Thermometer
 The modern
thermometer used in
our class is filled with
colored alcohol.
 As the bulb is heated or
cooled the liquid with
expand or contract.
Thermal Energy (Heat)
 Sum total of all the KE of the
particles in a sample. This can only
be measured using indirect means
when a change of heat occurs.
Temperature
 Measure of the average KE of the
particles in a sample. Can be
measured directly.
The Celsius Temperature Scale
 The freezing point of
pure water at sea
level is 0º C.
 The boiling point of
pure water at sea
level is 100º C.
The Difference between Kelvin and Celsius
 The main difference is the
location of the zero point.
 The zero point for kelvin is
called absolute zero.
 Absolute zero is equal to
-273.15º C or 0K.
 Absolute zero is the point at
which the motion of particles
of matter has completely
stopped.
Kelvin Temperature Scale
 SI Unit for temperature is Kelvin (K).
 The degree unit is not used in Kelvin (K),
Converting Kelvin and Celsius
 ºC = K – 273.15
 K = ºC + 273.15
 For example, the
boiling point of water
is 100 ºC is 373.15 K.
K = 100 ºC + 273.15
K = 373.15 K
Convert 50. K to the Celsius scale
 ºC = K – 273.15
 ºC = 50. K – 273.15
 ºC = -223.15 ºC
Converting Fahrenheit to Celsius
 ºC = (ºF – 32) x 5/9
 Convert 67°F to °C
 ºC = (67º – 32) x 5/9 = 19.4 ºC
Converting Celsius to Fahrenheit
 ºF = 9/5 x (ºC) + 32
 Convert -14 ºC to ºF
 ºF = 9/5 x (-14º) + 32 = 6.8ºF
Wednesday, 10/3/14
Learning Target:
Understand the characteristic differences
between physical changes and chemical changes of
matter.
Learning Outcome:
I correctly identify physical changes and
chemical changes that are demonstrated.
Properties of Matter
 Extensive Properties-properties that are
dependent on the quantity of matter. (mass,
volume, shape)
 Intensive Properties-Not dependent on
the size of the sample, characteristic
properties of that substance. (melting point,
boiling point, density)
Indicators of Chemical Change
 1. Evolution of heat and/or light.
 2. Production of a gas (not from boiling)
 3. Production of a precipitate (ppt.) (solid
but not from freezing)
 4. Color change (be careful with this one,
indicators cause color change but that is not
chemical!)
Physical vs. Chemical
 Examples:
 rusting
iron
 dissolving
in water
Chemical
Chemical
 burning
a log
Chemical
 melting
ice
Physical
 grinding
spices
Physical
Thursday, 10/6/14
Learning Target:
Explain the differences between elements,
compounds and mixtures (heterogeneous and
homogeneous).
Learning Outcome:
I will pre-lab for the Separation of Mixtures Lab.
What is matter?
Anything that has mass and takes up space.
Matter is broken up into two ways:
1. Pure Substances
2. Mixtures
5 States of Matter [Actually 4 States]
 Bose-Einstein Condensate




– very low volume, close to
absolute zero.
Solid-definite shape & volume,
maintains shape.
Liquid-definite volume but
indefinite shape, takes the
shape of its container but does
not fill.
Gas-indefinite shape & volume,
fills any container placed in.
Plasma-highly ionized form of
gas that exists at high temps.
(surface of the stars, fluorescent
lights)
Physical Characteristics
 Physical Properties-These are
observed or tested without changing the
substance.
 Physical change -These include
changes of state such as melting,
boiling, dissolving, grinding, filtering.
Chemical Characteristics
 Chemical Properties-How a substance reacts with
other substances. This is only observed in a chemical
reaction.
 Chemical Change-When a substance is converted
into a new substance. All properties and
characteristics will change!
 Format: Reactants
(start)

Products
(yields)
(ending)
Pure Substances
 Elements & Compounds
 These
always have the same
properties
 The
same composition
 They
can not be separated without
changing properties.
Element
 A substance that can not be broken down
into another substance by chemical means.
 The smallest part is an atom
 There are approximately 90 naturally
occurring elements.
Compound
 A substance that can be broken down
into another substance by chemical
means.
 The smallest part is a molecule or ion.
Mixtures
 Mixture-Physical combination of 2
or more substances.
 2 Classifications:
Heterogeneous-different composition
present
[examples: sand, granite, blood]

Homogeneous-same composition
present throughout
[examples: salt water, coffee, apple
juice]

Separation of Mixtures
 separate mixtures based on different physical
properties of the components
Different Physical Property
Technique
Boiling Point
Distillation
State of Matter (solid/liquid/gas)
Filtration
Dissolves in water
Evaporation
Filtration
Evaporation
Liquid vaporizes leaving less volatile liquid or
solid.
Distillation
Tuesday, 10/6/14
Learning Target:
Explain the differences heterogeneous and
homogeneous mixtures, and know techniques used to
separate them.
Learning Outcome:
I will complete the separation of mixtures lab.
Physical vs. Chemical
 Examples:
 melting
point
physical
 flammable
chemical
 density
physical
 magnetic
physical
 tarnishes
air
in
chemical
WARM UP
 A runner burns about 10. kcal per minute. If the
runner completes a race in one hour and fourteen
minutes, how many kJ did he burn? How many J did
he burn?
WARM UP
 A runner burns about 740 kcal in a cross country
race, how many kJ did she burn? How many J did
she burn?