HONORS CHEMISTRY UNIT TWO
MATTER : States, Changes, Mixtures &
Solutions


Matter
Matter – anything that has a mass and takes up space.
Law of Conservation of Mass/Matter - Matter cannot be created or
destroyed in a chemical reaction but only rearranged to form
different substances
PROPERTIES
CHEMICAL
PHYSICAL
INTENSIVE
EXTENSIVE
Part A: Characteristics of Matter
 Physical Properties - characteristics that can be observed without the
production of a new substance.
Examples:
color, smell, taste, hardness, density, texture,
melting/boiling/freezing points, magnetic attraction, solubility,
electrical conductivity, temperature, state or phase

Types of physical properties:

Extensive: depends on how much there is
Examples: volume, mass, weight, etc

Intensive: depends on the type of matter (not the amount)
Examples: density, melting/boiling pts, color (sometimes)
Part A: Characteristics of Matter
 Chemical Properties - describes how a substance reacts or fails to
react with other substances to produce new substances
Examples:
oxidation, corrosion, hydrolysis, combustion, flammability, reaction
to acid or base.
Part B: Kinetic Theory & Energy
 Kinetic Theory - all matter is made of tiny particles (atoms and
molecules) that are in constant motion

Potential Energy (PE) - energy due to position or condition
 At the atomic level it refers to the distance between particles
 Closer = lower PE
Farther = higher PE

Kinetic Energy (KE) – energy of motion
 Faster = higher KE Slower = lower KE
Part C: States of Matter
State or
Phase
Particle level
picture
Particles Description
Keep
Volume?
Solid
Arranged in orderly
pattern
Yes
Liquid
Touching, but not
tightly packed
No
Far apart and
rarely touching
No
No
Gas
Yes
Keep
Shape?
Yes
Part C: States of Matter
State or
Phase
Particle
Movement
Solid
Vibrational
only
Low
Low
Ice
Liquid
Vibrational &
translational
Moderate
Moderate
Water
Gas
Move freely
High
Vapor

Amount PE
Amount KE
High
Example
Plasma - 4th state of matter (extraordinary), consists of electrically
charged, high energy particles
Examples:
fluorescent lights, stars, lightning, lasers
PART C: CHANGE IN MATTER

Types of Changes
 Physical Change – an alteration of a substance that only
changes the physical properties of a substance * Does not
change the chemical composition!

Chemical Change – an alteration of the chemical composition
of a substance that results in the formation of a new
substance.
 Always forms a new substance with new physical and
chemical properties
 Also known as a chemical reaction
 Temperature change, color change, formation of a
precipitate, formation of a gas, and odor change
Matter
Pure Substances
Elements
Compounds
Mixtures
Homogeneous
Heterogeneous
Part D: Classification of Matter
 Pure Substance - made of only one type of matter
 Element - made of only one type of atom
 cannot be changed into simpler substances under normal ordinary
conditions
 Compound - atoms of two or more elements, chemically combined in
a definite ratio.

Mixture – a physical combination of two or more substances
no definite ratio of particles
 Homogeneous mixture - atoms of two or more elements, physically
combined in no definite ratio.
- The same throughout.
- Must be a SOLUTION
 Heterogeneous mixture - two or more types of atoms, physically
combined in no definite ratio.
-Different throughout
Part D: Classification of Matter
- Types of mixtures:
1.
2.
3.
Solution – Particles are Very tiny, will not separate by filtering,
will not settle out when allowed to stand, particles too small to
scatter light, (-) Tyndall effect. Ex. Kool Aid, brine
Colloid – Particles are tiny, will not separate by filtering, will not
settle out when allowed to stand, particles will scatter light, (+)
Tyndall effect. Ex. Milk, whipped cream, aerosols
Suspension – Particles visible with unaided eye, will separate when
filtered, will settle out if allowed to stand, particles will scatter
light, (+) Tyndall effect. Ex. Muddy water, snow globe
- Homogeneous mixture
 Solution
 a solute dissolved in a solvent
 The solvent is the part in greater quantity (usually water)
solute
solvent
- Homogeneous mixture
 Solution
Example: salt water
- Homogeneous mixture
 Solution
 Electrolyte – a solution that conducts electricity in water
 Example: salt water
 Non-example: sugar water
- Homogeneous mixture
 Solution
 Soluble – able to dissolve
 Salt dissolves in water (opposites attract)

Water is a polar molecule and a universal solvent
Dipole – dipole forces: attraction between oppositely
charged regions of polar molecules
 “Like dissolves like” – polar dissolves polar (salt & water), non-polar dissolves
non-polar (oil in gasoline)

- Homogeneous mixture
 Solution
1) The solution is well stirred during formation.
2) Particles will not settle out
3) It is clear and transparent
4) Considered to be in one phase and will not filter unless extremely
fine filter paper is used

Types of solutions
1) Gas-Gas - Carbon dioxide, Nitrogen, Oxygen (air)
2) Liquid-Gas - Water Vapor in Air (moist air)
3) Gas-Liquid - Carbon dioxide in Water (soda water)
4) Liquid-Liquid - Acetic acid in Water (vinegar)
5) Solid-Liquid - Sodium chloride in Water (brine or salt water)
6) Solid-Solid - Copper in Silver (Sterling Silver)
Part E: The Mole
Atomic Mass
 A single atom has a very small mass (on the order of 10-23 grams per
atom)
 Because this mass is so small, we use a unit called amu to describe the
mass of a single atom (for example, 1 atom of carbon has a mass of
12.011 amu)
 Average atomic mass - the weighted average of the masses of all
isotopes of an element
 atomic mass units are not practical for use in the lab (too small a
quantity)
 grams are the preferred unit of mass
 Therefore, scientists needed a way to determine the number of
atoms in a given mass of an element
The Mole
 Carbon-12 was selected as the standard. The number of atoms in 12.0 g
of C-12 was determined experimentally using sophisticated equipment
 The number came out to be 6.022 x 1023  this number is called a
mole (mol)
 The amount of a pure substance that contains 6.022 x 1023 particles of
that substance
 Try it this way…
There are 6.022 x 1023 carbon atoms in 12 grams of carbon
 There are 6.022 x 1023 hydrogen atoms in 1.0 grams of hydrogen
 There are 6.022 x 1023 oxygen atoms in 16.0 grams of oxygen
 There are 6.022 x 1023 gold atoms in 197 grams of gold

The Mole (cont.)
 The mole establishes a relationship between the atomic mass unit
and the gram.
 The mole is used to describe a huge amount of any extremely small
particle. A mole of gold, a mole of salt and a mole of water each
contain 6.022 x 1023 individual units.


How would you feel about inheriting a mole of pennies???
$6.02 x 1021 or $6,020,000,000,000,000,000,000
If you gave $1million a day to every person on Earth, it would
take you > 3000 yrs to run out of money
Calculating Molar Mass
 The mass in grams of 1 mole of a substance.
 Example: H2O
H 2 x 1.01 = 2.02 g
O 1 x 16.00 = 16.00 g
18.02 g
The molar mass of H2O = 18.02 g/mol
 Keep two digits after the decimal for molar mass!
Converting Moles  Grams
 How many grams are in 3.6 moles of NaCl?
3.6 moles NaCl x
=
Converting from Grams  Moles
 How many moles are in 1.75 g of BeF2?
1.75 g BeF2 x
=
Part F: Concentration of Solutions
 Molarity (M) – The number of moles of solute dissolved in each liter of
solution
Molarity = moles of solute
liters of solution



Problems:
What is the molarity when 0.75 mol is dissolved in 2.50 L of solution?
If you had 75.0 g of NaCl and you dissolved it in exactly 2.00 L of
solution. What would be the molarity of the solution?
Part F: Concentration of Solutions
 Problems:
 If you wanted to make a 3.0 L of 4.0 M HCl solution. How much HCl
would you need in grams?
Part G: Solution Solubility
 Solution
 Solubility - the amount of solute that dissolves in a given quantity of
solvent at a given temperature to produce a saturated solution
 Expressed in g/100 g solvent

Solution concentrations:
 Saturated solution – Contains the maximum amount of solute for a
given amount of solvent at a constant temperature.
 Unsaturated solution - Contains less solute than a saturated solution.
 Supersaturated solution – Contains more solute than is should be able
to hold at a given temperature. (Unstable)

Factors that affect solubility:
 Nature, temperature and pressure, amount of solute already dissolved
- Homogeneous mixture
 Solution

Increase solubility by…

Stir the solution – increase of collisions and the breaking of solute attraction

Decrease the particles size of the solute – increases surface area, smaller particles –
more area for contact with the solvent

Increase temperature – in crease the temperature of the solvent increases kinetic
energy and the number of collisions between particles
- Homogeneous mixture
o Solubility curve
- Homogeneous mixture
 Solubility curve questions
1) Generally, what happens to the solubility of a substance as its
temperature rises?
2) A solution is made by dissolving 40 g in 100 g of water at 80°C.
What is the substance?
3) How many grams of this substance would be required to make a
100 g solution at 20°C?
4) A solution is made with 120 g of NaNO3 at 70°C. Is this solution
saturated, unsaturated or supersaturated?
5) What happens to a solution of NaNO3 when it is cooled from
80°C to 60°C?
6) What mass of NaNO3 is needed to make a saturated solution at
20°C with 250 g of water?