2.1 Properties of Matter
• Define and prove the Law of Conservation of Matter
• Describe a substance according to its physical and
chemical properties.
• Distinguish between extensive and intensive
properties.
• Describe the six states of matter.
• Identify physical changes to matter.
• Identify the chemical properties of a substance.
• Describe chemical changes and differentiate them
from physical changes.
• Use various visual clues to identify whether a
chemical reaction is taking place.
2.2 Classification of Matter
• Define a mixture and understand why
mixtures are different than pure substances.
• Classify mixtures as homogeneous or
heterogeneous
• Classify liquids as a solution, suspension,
colloid, or alloy
• Describe several ways to separate mixtures.
• Distinguish between elements and
compounds.
Classifying Matter
• Scientists classify matter according to make-up.
Matter
Elements
Compounds
Mixtures
Elements and Compounds
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

An element is the simplest form of matter
and has a unique set of properties.
A compound contains two or more
elements chemically combined in fixed
proportions.
Compounds can be broken down,
elements cannot.
Elements
• Simplest pure substance – fixed composition.
• Pure substance is made of only one material. Same
throughout. Referred to as homogeneous matter.
• Are made of alike atoms, which are the basic
building blocks of matter.
• Represented by chemical symbols - H, O, Al, Cu,
Au.
• 118 elements organized in the periodic table in
squares called an element key.
Compounds
• Two or more elements chemically combined.
• Represented by chemical formulas.
• Examples: H2O, CO2, NaCl, C12H22O11
• Molecules are the smallest natural unit of a compound
or diatomic element. A molecule is made up of two or
more atoms chemically combined.
Pure Substances: Elements are Composed
of identical atoms
Chlorine gas
Sodium metal
2 or More Elements Combine to
Produce a Compound
Chlorine gas
+
=
Sodium metal
+
Sodium Chloride
=
Mixtures
• Two or more substances mixed together but not
chemically combined.
• Has unlike parts and a composition that varies
from sample to sample.
• Properties:
– Sub. in a mixture keep their separate identities.
– Sub. can be present in any amount.
– Sub. can be separated out by physical means.
Mixtures
Mixtures
Homogeneous
Heterogeneous
Mixtures: Variable combination of
two or more substances:

Homogeneous
uniform throughout

Heterogeneous
not uniform throughout
Phase



Used to describe any part of a sample with
uniform composition and properties
Heterogeneous: 2 or more phases
Homogeneous: 1 phase
Solutions
• A homogeneous mixture.
• One substance dissolves in another. Solutions remain
constantly & uniformly mixed.
•Solute and Solvent
•Particles are smaller than 1 nm in size
•Particles do not settle
•Particles pass through filter paper
Examples: lemonade, soda, ocean water, antifreeze,
metal alloys
Alloy


A homogeneous mixture combining two or
more metals to maximize the benefits of
each.
Ex. Gold Jewelry: 14 kt Gold (Gold for
shine, Silver for strength)
Colloids
• A homogeneous mixture.
• Mixed together but not dissolved.
• Appears cloudy. Scatters light.
•Particles are 2 – 100 nm in size
•Particles do not settle
•Particles pass through filter paper
•Usually not transparent (cloudy)
• Examples: milk, cool whip, toothpaste, lotions,
fog, smoke, Jello.
Suspension





Heterogeneous Mixture
Particles settle to the bottom unless
constantly stirred
Particles are greater than 100 nm
Not transparent, must be mixed
Ex: Paint, Chocolate Milk,
Composition of Matter
Matter
Pure Substance
Element
Compound
Mixture
Homogeneous
Solution
Colliod
Heterogeneous
Suspension
2.3 Changing Matter
• Describe how chemical potential energy relates to
heat and work.
• Describe the law of conservation of energy and
how heat flows between system and surroundings
during both endothermic and exothermic
processes.
• Use the specific heat equation to perform
calculations that relate mass, specific heat, change
in temperature, and the amount of heat absorbed
or released.
Thermochemistry
Chemical Potential Energy
• the energy stored in the
chemical bonds of a
substance
• the kinds of atoms and
their arrangement in the
substance determine the
amount of energy stored in
the substance.
If 2 objects remain in contact,
heat will flow from a warmer
object to the cooler object until
the temperatures are equal.
AKA:
HEAT
(A) Object A starts with a higher temperature than object B. No
heat flows when the objects are isolated from each other.
(B) When brought into contact, heat flows from A to B until the
temperatures of the two objects are the same.
Heat
• Symbol = q
• energy that transfers from one object to another
because of temperature difference
• Heat always flows from a warmer object to a cooler
object.
An Ice Cold Spoon
energy
transfer
A Hot Spoon
Thermochemistry
• The study of energy changes that occur
during chemical reactions and changes of
state.
Law of
Conservation of Energy
• In any chemical or physical
process, energy is neither created or
destroyed
• If energy of the system decreases
the energy of the surroundings must
increase by the same amount so that
the total energy of the universe
remains unchanged.
The Universe
•
•
•
•
Can be divided into 2 “parts”
System- the part you are investigating
Surroundings- the rest of the universe
In a thermo-chemical experiments the
region in immediate vicinity of the
system are the surroundings.
Units of Heat
calorie
joule
• the amount of energy
contained within food
• SI Unit of energy
• Quantity of heat required to to
raise the temp of 1g of water
1° C
• Heat changes in
chemical reactions
are typically
measured in joules.
Calorie=Kilocalorie=1000calories
Joule and calorie
Conversion Factors
• 1J = 0.2390cal
• 1 cal = 4.184 J
HEAT
HEAT
HEAT
HEAT
Different type of materials may have the same
temp, same mass, but different conductivity.
•They are affected by the potential energy
stored in chemical bonds or the IMFs holding
molecules together
•It is possible to be at same temp (same KE) but
have very different thermal energies.
•The different abilities to hold onto or release
energy is referred to as the substance’s
heat capacity
Heat Capacity
The measure of how well a material absorbs
or releases heat energy
•Physical property unique to a particular material
•The heat capacity depends on both its mass and its chemical
composition.
•The greater the mass; the greater the heat capacity.
•It can be thought of as a reservoir to hold heat, how much
it holds before it overflows is its capacity
•Water takes 1 calorie of energy to raise temp 1 °C
•Steel takes only 0.1 calorie of energy to raise temp 1 °C
Specific Heat
The amount of heat
required to raise
the temperature of
one gram of
substance by one
degree Celsius.
q = mCT
Calculations Involving Specific Heat
q  c m  T
p
c = Specific Heat
q = Heat lost or gained
T = Temperature change
m = Mass
Table of Specific Heats
Calorimetry
• The precise measurement of heat flow into or
out of a system for chemical and physical
purposes
• The heat released by the system is equal to the
heat absorbed by its surroundings.
• Conversely, the heat absorbed by a system is
equal to the heat released by it’s surroundings.
Calorimeter
There are three methods used to
transfer heat energy.
•
•
•
Conduction – transfer of heat through
direct contact
Convection – transfer of heat through
a medium like air or water
Radiant – transfer of heat by
electromagnetic radiation