Matter and Energy
Zumdahl Chapter 3
Grand Prismatic Spring in
Yellowstone National Park
Source: David Maisel/Stone/Getty Images
Why
does
soda
fizz
when
you
open
the
bottle?
Universe Classified



Matter is the part of the universe that has
mass and volume
Energy is the part of the universe that has
the ability to do work
Chemistry is the study of matter


The properties of different types of matter
The way matter behaves when influenced by
other matter and/or energy
3.1 Matter
AIMS
-learn about matter and its three states
What is matter??
Aristotle’s View….not science, but
logic.
Matter

What is Matter???

Anything that has mass and
occupies a volume

“Matter is the real substance of which
actual physical objects - the 'things' of this
world - are composed.”
Sir Roger
Penrose.
Can also be thought of as being composed
of protons, neutrons, and electrons (more
on that later)

All Matter exists in one of 4
Physical States




Gaseous-no definite
shape or volume
Liquid-definite
volume, but not shape
Solid-definite shape
and volume
Plasma-like gases, but
atoms are made up of
free electrons and ions
of the element.
Figure
3.1:
Liquid
water
takes the
shape of
its
container.
Some Characteristics of Gases, Liquids and Solids and the
Microscopic Explanation for the Behavior
gas
liquid
solid
assumes the shape
assumes the shape of the retains a fixed volume
and volume of its
part of the container
and shape
container
which it occupies
rigid - particles locked
particles can move
particles can move/slide
into place
past one another
past one another
compressible
lots of free space
between particles
not easily compressible
little free space between
particles
flows easily
particles can move
past one another
flows easily
particles can move/slide
past one another
not easily
compressible
little free space
between particles
does not flow easily
rigid - particles cannot
move/slide past one
another
Examples of substances and their
commonly found state

Solid

Ice cube, diamond, iron
bar, rock

Liquid

Gasoline, water,
alcohol, blood

Air, helium, oxygen,
natural gas

Gas
What determines the Physical State
of a particular substance?

Temperature

Melting point or
boiling point of
substance
What causes a substance to change
into another Physical State?


Addition of energy
-melting
-boiling
Subtraction of Energy
-freezing
-condensing
(condensation)
Densities of Physical States

Gases


Least dense
Occupy about
16oo greater
volume than
liquids
Liquids
Density similar
to solids, little
less
Solids
Most dense
of states
AeroGel—an ultra-low density solid
Aerogels have a very good thermal resistance
.
This photo shows a
2.5 kg
Brick being
supported by only 2
grams of aerogel
Mercury, on the other hand, is a very dense
liquid, denser than some solids
Figure 3.2: The three states of
water.
3.2 Physical and Chemical
Properties and Changes
AIMS
learn to distinguish between physical and
chemical properties
-learn to distinguish between physical and
chemical changes
Properties


Characteristics of the substance under
observation
Properties can be either
directly observable or
the manner something interacts
with other substances in the universe
Properties of Matter

Physical Properties are the characteristics of
matter that can be observed without changing its
composition


Characteristics that are directly observable
Chemical Properties are the characteristics that
determine how the composition of matter changes
as a result of contact with other matter or the
influence of energy

Characteristics that describe the behavior of matter
What is a Physical Property?

A physical property is a characteristic of a substance
that does not change as you determine that property.
Example: In order to determine that a domino is
black, all you do is look at it. The domino remains
unchanged during the test
Example: In order to determine if a rock contains
the mineral magnetite, you could place a paperclip
nearby and see if it sticks. The rock remains
unchanged during this test.
Examples of Physical Properties










Odor
Color
Physical state
Boiling/melting point
Density
Magnetic/not
Hardness
Radioactivity
Malleability
Ductility
Gallium
metal has
such a low
melting
point
(30°C) that
it melts
from the
heat of a
hand.
An iron
pyrite
crystal
(gold
color) on a
white
quartz
crystal.
Source:
Chip Clark
What is a Chemical Property?

A chemical property is a characteristic of a
substance which, when determining, will
change
an example would be …when determining the
flammability of a substance, we must burn it,
and thereby change it.
chemical properties refer to the ability of a
substance to form new substances
Some Chemical Properties


Flammability
-is the material flammable?
Reactivity


-does material react with another chemical, in what
way?
pH

What is the pH of the substance?
Classify Each of the following as Physical
or Chemical Properties

The boiling point of ethyl alcohol is 78°C.


Diamond is very hard.


Physical property – describes inherent characteristic of
alcohol – boiling point
Physical property – describes inherent characteristic of
diamond – hardness
Sugar ferments to form ethyl alcohol.

Chemical property – describes behavior of sugar –
forming a new substance (ethyl alcohol)
Changes in Matter

Physical Changes are changes to matter that do not
result in a change the fundamental components that
make that substance


State Changes – boiling, melting, condensing
Chemical Changes involve a change in the
fundamental components of the substance



Produce a new substance
Chemical reaction
Reactants  Products
What is a Physical change?

Any change that does not change the chemical
makeup of a substance

Examples include:


Changes in physical state
Crushing, grinding, tearing
What is a Chemical change?

Any change that results in materials with a
different chemical nature.

Examples:
 Electrolysis
 Burning (combustion)
 Silver tarnishing
Figure 3.3:
Electrolysis,
the
decomposition
of water by an
electric
current, is a
chemical
process.
Electrolysis Link
Signs of a Chemical Change



Color Change
Formation of a gas (not from a phase change)
Energy absorbed or given off



Gets cold or hot
Gives off light/electric charge/other energy
Formation of a precipitate
Oxygen
combines
with the
chemicals in
wood to
produce
flames. Is a
physical or
chemical
change taking
Source:
Jim Pickerell/
Stone/Getty
Images
place?
Classify Each of the following as Physical
or Chemical Changes

Iron metal is melted.

Iron combines with oxygen to form rust.

Sugar ferments to form ethyl alcohol.
Classify Each of the following as Physical
or Chemical Changes

Iron is melted.


Iron combines with oxygen to form rust..


Physical change – describes a state change, but the
material is still iron
Chemical change – describes how iron and oxygen react
to make a new substance, rust
Sugar ferments to form ethyl alcohol.

Chemical change – describes how sugar forms a new
substance (ethyl alcohol)
Classification of Matter
Matter
Pure Substance
Constant Composition
Homogeneous

Homogeneous = uniform throughout, appears to be one thing



Mixture
Variable Composition
pure substances
solutions (homogeneous mixtures)
Heterogeneous = non-uniform, contains regions with
different properties than other regions
3.3 Elements and
Compounds
AIM: to understand the definitions of elements
and compounds
Elements




Cannot be separated into simpler substances
Fundamental form of matter
Examples include iron, oxygen, aluminum,
and hydrogen
There are 90 naturally occurring elements,
about 118 known elements
Compounds



Have same composition no matter where they
are found
ALWAYS made up of atoms of 2 or more
elements, chemically combined
Can be broken into the constituent elements
by chemical changes
Elements and Compounds


Substances which can not be broken down into
simpler substances by chemical reactions are
called elements
Most substances are chemical combinations of
elements. These are called compounds.




Compounds are made of elements
Compounds can be broken down into elements
Properties of the compound not related to the
properties of the elements that compose it
Same chemical composition at all times
A Model of an oxygen molecule….
Is this an element or a compound?
Element
Element
Compound
Would you
put elemental
sodium metal
and chlorine
gas on your
popcorn?????
Or would you put
a compound
composed of
sodium ion and
chloride ion
(NaCl) on it?
Mixtures vs. Pure Substances
Pure Substances vs. Mixtures



Pure Substances
 All samples have a unique set of physical and chemical
properties
 Constant Composition  all samples have the same
composition
 Homogeneous (meaning same throughout)
 Separate into components based on chemical properties
 Include Elements and Compounds
Mixtures
 Different samples may show different properties
 Variable composition
 Homogeneous or Heterogeneous (different throughout)
 Separate into components based on physical properties
All mixtures are made of pure substances
Creamtop
milk
bottles
Grain Probe, for sampling corn, soybeans, etc
Types of Mixtures


Solid mixtures can simply be a mix of various
substances, such as pancake mix, or a can of
peas and carrots
Solid mixtures also include alloys which are
physical mixtures of metals.

examples 14 K gold, steel, brass
Examples



Liquid mixture--Ocean water contains water,
dissolved salts, oxygen, and other gases
Solid mixture--14 K gold contains gold and
other metals (only physically combined)
Liquid mixture--Blood contains water,
dissolved ions, proteins, red blood cells
Element,
Compound,
or Mixture?
Types of Mixtures

Liquid mixtures fall into the following 3 categories...


Solution—particles are dissolved and don’t settle out.
Typically transparent, cannot be filtered


Colloid-particle size a little bigger, still don’t settle out.
Colloids exhibit Tyndall Effect, appear cloudy, murky, or
opaque


Salt water, Kool-Aid
Milk, cloudy water, fog
Suspension—particles are fairly large, and will settle out
of the mixture, separate upon standing still, can be filtered

Air/dust, river water/gravel/mud/sand,
Behaviors of Mixtures

Suspension - particles
are readily sedimented
by gravity or by
centrifugation.
Soil suspension sedimentation by
gravity
Suspension of
blood cells centrifugation
at 500 x g

No change after
centrifugation at
10,000 x g
True solutions
and colloidal
mixtures remain
dispersed even
following
centrifugation.
Hemoglobin
(true solution)
Milk (colloidal
mixture)
Separation of Mixtures

Separate mixtures based on different
physical properties of the components

Physical change
Different Physical Property
Technique
Boiling Point
Distillation
State of Matter
(solid/liquid/gas)
Adherence to a Surface
Filtration
Chromatography
Volatility
Evaporation
A honey extractor,
separates honey from
comb
Graded sieves
Filter paper
Used coffee filter
Separatory Funnel
Figure 3.4: When table salt is stirred into
water (left), a homogeneous mixture
called a solution forms (right).
Figure 3.5: Sand and water do not mix to
form a uniform mixture. After the mixture
is stirred, the sand settles back to the
bottom.
Figure 3.6: Distillation of a
solution consisting
of salt dissolved in water.
Identity Each of the following as a Pure
Substance, Homogeneous Mixture or
Heterogeneous Mixture

Gasoline


A stream with gravel on the bottom


a homogenous mixture
a heterogeneous mixture
Copper metal

A pure substance (all elements are pure substances)
Figure 3.10:
The
organization
of matter.
Organization of Matter
Quia - Mixtures, elements and
compounds “Rags to Riches” Game
Energy and Energy Changes


Capacity to do work
 chemical, mechanical, thermal,
electrical, radiant, sound, nuclear
Energy may affect matter


e.g. raise its temperature, eventually causing a
state change
All physical changes and chemical changes
involve energy changes
Heat

Heat: a flow of energy due to a temperature
difference
1.
Exothermic = A process that results in the
evolution of heat.

2.
Example: when a match is struck, it is an
exothermic process because energy is produced as
heat.
Endothermic = A process that absorbs energy.

Example: melting ice to form liquid water is an
endothermic process.
Units of Energy

One calorie is the amount of energy needed to raise
the temperature of one gram of water by 1°C


joule


kcal = energy needed to raise the temperature of 1000 g of
water 1°C
4.184 J = 1 cal
In nutrition, calories are capitalized

1 Cal = 1 kcal
Example - Converting Calories to Joules
Convert 60.1 cal to joules
Converting kcal to Joules

Convert 56.7 kcal to joules
Converting kJ to kcal

Convert 45.06 kJ to kcal
Figure 3.11: In ice, the water molecules
vibrate randomly about their positions in the
solid. Their motions are represented by
arrows.
Figure 3.12: Equal masses of hot water
and cold water separated by a thin metal
wall in an insulated box.
Figure 3.13: The H2O molecules in hot
water have much greater random motions
than the H2O molecules in cold water.
Figure 3.14: The water samples now have
the same temperature (50°C) and have the
same random motions.
Energy and the 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 depends on its mass

If you double the mass it will take twice as
much heat energy to raise the temperature the
same amount.
A burning
match
releases
energy.
Source:
ElektraVision/
PictureQuest
Specific Heat Capacity

Specific Heat (c) is the amount of energy
required to raise the temperature of one
gram of a substance by one Celsius degree
J
By definition , the specific heat of water is 4.184
g C
Amount of Heat = Mass x specific heat x Temperature Change
Q = m x c x T
Example – 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
JJ
Specific Heat of Water = 4.184
gg-CC
Mass = 7.40 g
Temperature Change = 46.0°C – 29.0°C = 17.0°C
Q = m x c x T
J
Heat  4.184
 7.40g  17.0C  526 J
g C
Example – 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.
Q  s  m  T
Is the metal pure gold?
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 3.2 lists the specific heat of gold as 0.13
Therefore the metal cannot be pure gold.
J
g C
Heating and Cooling Curves
Heating Curve





As heat added to solid, it first raises the temperature of the
solid to the melting point
Then added heat goes into melting the solid
 Temperature stays at the melting point
 Heat of Fusion
As more heat added it raises the temperature of the liquid
to the boiling point
Then added heat goes into boiling the liquid
 Temperature stays at the boiling point
 Heat of Vaporization
As more heat added it raises the temperature of the gas
Figure 13.2: The heating/cooling curve for
water heated or cooled at a constant rate.
Heating/Cooling Curve web site

http://www.kentchemistry.com/links/Matter/H
eatingCurve.htm
Energy Requirements for
State Changes

In order to change a liquid to a gas, must
supply the energy required to overcome the
all the intermolecular attractions


Not break bonds (intramolecular forces)
The energy required to boil 1 mole of a
liquid is called the Heat of Vaporization

Hvaporization = 40.6 kJ/mol for water at 100°C
Energy Requirements for
State Changes


In order to change a solid to a liquid must
supply the energy required to overcome the
some of the intermolecular attractions
The energy required to melt 1 mole of a
solid is called the Heat of Fusion

Hfusion = 6.02 kJ/mol for ice at 0°C
Structures of the States of Matter

In solids, the molecules have no translational
freedom, they are held in place by strong
intermolecular attractive forces


In liquids, the molecules have some translational
freedom, but not enough to escape their attraction
for neighboring molecules


may only vibrate
they can slide past one another, rotate as well as vibrate
In gases, the molecules have “complete” freedom
from each other
Why do Molecules Attract Each Other?



Intermolecular attractions are due to attractive forces
between opposite charges
+ Ion to - ion
+ End of polar molecule to - end of polar molecule



H-bonding especially strong
Larger the charge = Stronger attraction
Even non-polar molecules have attractions due to
opposite charges

London Dispersion Forces
Figure 13.4: Intramolecular (bonding) forces exist
between the atoms in a molecule and hold the
molecule together. Intermolecular forces exist
between molecules.