Download Unit 2: Exploring Matter - Fort McMurray Composite High School

Unit 2: Matter and Chemical Change
Lesson 1: Properties of Matter
All chemicals should be handled very
carefully.
WHMIS (Workplace Hazardous Materials
Information System) has been developed to
provide guidelines for handling, storage and
disposal of reactive materials (chemicals).
Some materials are CAUSTIC - will burn,
corrode or destroy organic tissue
MATTER: Matter is anything that has mass.
All matter is described in terms of its
properties
PROPERTIES - are characteristics you can
use to describe or identify different
substances. Ex. Color, luster, state, taste,
melting point, and behavior
Properties can be classified as:
1. QUALITATIVE PROPERTIES
- Those properties, which describe a quality of
matter that has no numerical value (no
number)
- Usually involves one of the senses
Ex. Taste, odor, texture, color, luster
2. QUANTITATIVE PROPERTIES
-Those properties, which describe a quantity
of matter
- Have a number associated with the property
Ex. Melting point of water is 00 C
 Freezing point, boiling point, number of
legs, density etc.
CHEMICAL PROPERTIES – those properties
that involve the formation of a new substance.
Chemical properties cannot be tested without
destroying the substance.
Ex. Rust is formed when iron reacts with
oxygen
Magnesium burns to produce a white
powder
Paper burns to produce carbon
dioxide and water
PHYSICAL PROPERTIES – do not involve the
formation of a new substance. Physical
properties can be tested without destroying
the substance.
Ex. Melting point – temperature at which
something changes from a solid to a
liquid
Boiling point – temperature at which
something changes from a liquid to a
gas
Malleability – can be pounded or rolled
into sheets
Ductility – can be stretched into a wire
Density – the amount of mass in a given
volume of a substance
Conductivity – ability of a substance to
conduct electricity
Solubility – ability to be dissolved in
another
The Particle Model of Matter states that:
1. All matter is made up of extremely tiny
particles.
2. Each pure substance has its own kind of
particle, different from the particles of other pure
substances.
3. Particles attract each other.
4. Particles are always moving.
5. Particles at a higher temperature move faster
on average than particles at a lower temperature.
All matter can be classified according to its state
as a
1. Solid – has a definite shape and a
definite volume
- Particles are close together and the
forces between the particles are
strongest
2. Liquid
- has a definite volume but
takes the shape of its container
- Are said to be fluid because they
flow
- Particles are further apart and the
forces between the particles are
weaker
3. Gas - has neither a definite shape nor a
definite volume
- Expands to fill the container
- Gases are fluids because they also
flow
- Particles are furthest apart and the
forces between the particles are very
weak
CHANGES OF STATE
1.
2.
3.
4.
5.
Melting
Freezing
Boiling/Vaporizing
Condensation
Sublimation
Classification of Matter
Matter can also be classified according to its composition:
Matter
Pure Substances
Mixtures
Elements Compounds
Mechanical
Mixtures
(heterogeneous)
Suspension
Colloids
Solutions
(homogeneous)
PURE SUBSTANCES have properties that
are always the same.
Ex.
melts
Table salt is a white solid that
at 801 °C and boils at 1465 °C.
Water, vinegar, sugar etc.
Elements
-Pure substances that cannot be broken
down into any simpler substances - most
elements are solids but several are gases
and two are liquids
-All of the elements have been arranged on
the Periodic Table according to certain
properties.
- Contains only one type of atom
Ex. Silver, Oxygen, Iron, Carbon, Mercury
Compounds
- Pure substances that contain two or more
different elements, combined in a definite
fixed proportion.
- Can be broken down chemically into
different substances since it is made up of
different kinds of atoms
Ex.
Water - 2 hydrogen + 1 oxygen ---> H2O
Salt - 1 sodium + 1 chlorine ---> NaCl
MIXTURES contain at least two different
substances. Properties are variable
Solutions -Homogeneous are a mixture in
which one substance is dissolved in another.
It is a uniform mixture that appears the same
throughout.
Ex. Salt water, apple juice, and air
Mechanical - Heterogeneous are mixtures that
do not appear the same throughout. The
different components are visible.
Ex. Soil, chocolate chip cookies,
chicken noodle soup
Suspensions
– Heterogeneous mixture made of large
particles that are uniformly mixed but will
settle if left undisturbed
Ex. Sand in water, powdered chalk in water
Colloids
- Heterogeneous mixture composed of fine
particles evenly distributed throughout a
second substance
Ex. Hair gel
Emulsions
- Type of colloid in which liquids are
dispersed in liquids
- Many will separate quickly to form layers
of the original liquids (oil and vinegar)
Lesson 2: Chemical and Physical Changes
Physical Changes
- No new substance produced
- Change of size, shape or state
- Ex. Cutting, freezing, molding, boiling, dissolving
Chemical Changes
- Starting material is used up
- New substance formed with different properties
- Atoms are rearranged to form new molecules
- Changes cannot be reversed
- Ex. Cooking an egg, rusting, burning
Evidence of a Chemical Change
(a) Color Change
(b) Gas Formed – bubbling
(c) Solid material, called a precipitate is formed.
- two solutions are combined and a solid is formed
(d) Energy Change - energy is the ability to do work
- Ex. Light, heat, mechanical, sound, electrical
There are 2 types of energy change:
1. Endothermic – energy is required.
- Energy is added to the starting materials
- Ex. cooking
2. Exothermic – energy is released.
- Ex. Burning
Identify the following substances as pure
substances (element / compound) or as a
mixture (homogeneous / heterogeneous):
1. zinc
6. vinegar
2. carbon dioxide
7. tossed salad
3. orange juice
8. aluminum
4. nitrogen
9. kool-aid
5. sugar
10. windex
Lesson 3: History of the Atomic Theory
Aristotle: (350 BC) – Greek philosopher
- Believed that everything was made of
1. Earth (dry and cold)
2. Air (wet and hot)
3. Fire (dry and hot)
4. Water (wet and cold)
Robert Boyle: (1660’s) – England
Recognized that elements could be combined to form
compounds
Lavoisier: (1770-1780) – France
1. Defined elements as pure substances that cannot be
decomposed (broken down into simpler substances)
2. Developed a system for naming chemicals, so that all
scientists could use the same words
3. Identified 23 pure substances as elements
4. Discovered that in a chemical change, “the mass of
the new substances is always the same as the mass of
the original substances” – LAW OF CONSERVATION
OF MASS.+
John Dalton: (1808) – England
Atomic Theory:
1. All matter is made up of small particles called atoms
2. Atoms cannot be created, destroyed or divided into
smaller particles
3. All atoms of the same element are identical in mass
and size. Atoms of one element are different in mass
and size from the atoms of other elements
4. Compounds are created when atoms of different
elements link together in definite proportions.
Dalton’s Theory led to the current definitions:
Element – a pure substance made up of one type of particle,
or atom.
Compounds – pure substances that are made up of 2 or more
elements chemically combined together. Compounds can
be broken down into elements again by chemical means.
J.J. Thompson: (1897) England
- Raisin bun model (plum pudding)
- Atom is a sphere, which is positive, with negative
electrons embedded in it like raisins in a bun
Ernest Rutherford: (1911) McGill University, Canada
- Atoms have a nucleus which is positive
- Most of the atom is empty space occupied by the
moving negatively charged electrons
- Proposed the existence of protons in a nucleus
Neils Bohr: (1913) – Danish
-Electrons move in circular orbits around the nucleus
- Like a miniature solar system
James Chadwick: (1932)
- showed that the nucleus must contain heavy neutral
particles to account for all of the atom's mass
- proposed the existence of neutrons
Lesson 4: Element Symbols
- All elements have been given an atomic symbol
(a) A single capital letter – O – oxygen
(b) Capital letter & a lower case letter – Co – cobalt
(c) Capital letter & 2 lower case letters – Uun – ununnilium
- In the 1860’s Dmitri Mendeleev, a Russian chemist arranged
the elements in order of increasing ATOMIC MASS and
created the PERIODIC TABLE
- ATOMIC MASS is the average mass of an atom of an
element
Ex. Oxygen = 16.00 g/mol
- Mendeleev found that the properties of the elements
repeated at definite, or periodic intervals (ex. Lithium,
sodium and potassium have similar properties so he placed
them in the same family or vertical row)
- He left blanks in the table where he predicted elements
should be and predicted what their properties would be,
based on where they were on his table
- After the development of atomic theory, the periodic table
was rearranged in order of increasing ATOMIC NUMBER
- ATOMIC NUMBER is the number of protons an element has
in its nucleus
Ex. Fluorine – atomic number = 9, therefore it must have
9 protons in its nucleus.
The Periodic Table contains a lot of information
about the different atoms. For example:
Atomic number
4
Name
Symbol
Be
Beryllium
9.01
Atomic Mass
The horizontal rows on the periodic table are called Periods.
The vertical rows on the table are called Groups or
Families.
Elements in the same family have similar properties (behave
in a similar manner)
There are 18 Groups or Families.
The key on the periodic table will indicate the state of each
element. Ex. White box – solid
Grey box – gas
Black box - liquid
- All elements can be classified as metals, non-
metals or metalloids depending on their
properties
Metals
- Found to the left of the
staircase line
- 80% of all elements
- Lustrous (shiny)
- Ductile
(stretched into wire)
- Malleable (hammered/shaped)
- Conduct electricity
- All solids, except mercury
liq.
- Ex. Sodium, iron
Non-metals
- located to the right of
of the staircase line
- 20 % of all elements
- dull
- non-ductile
- brittle
- non-conductors
- Mostly gases, some metals, 1
Ex. Oxygen, bromine
Metalloids - these elements have properties of
both metals and non-metals
Ex. Silicon – shiny like a metal, poor
conductor like a non-metal
There are 4 special named groups in the table:
Group 1 – Alkali Metals
- Most reactive metals
- Never found in pure form in nature
Ex. Lithium, sodium, potassium
Group 2 – Alkaline Earth Metals
- React fairly vigorously with some substances
Ex. Magnesium, calcium, barium
Group 17 – Halogens
- Most reactive non-metals
Ex. Fluorine, Chlorine
Group 18 – Noble Gases
- Most non-reactive elements
- Used to be called “Inert” gases until 1963 when a
Canadian chemist at UBC, made some of them react
- Different noble gases produce different colors
Ex. Argon – blue
Helium – yellow-white
Lesson 5: ATOMIC STRUCTURE

Atom - the smallest part of an element (which
retains the chemical and physical properties of
the element). Atoms are made up of 3 subatomic particles
1. Electron (e)
-Smallest particle in an atom
-Has a negative charge
-Located in the extra nuclear region of the
atom - outside the nucleus
2. Proton (p)

-Has a large mass
-Has a positive charge
-Located inside the nucleus
3. Neutron (n)
-Same mass as a proton
-Has a neutral charge (no charge)
-Located inside the nucleus
Nuclear Notation
- Atomic number is the number of protons in the
nucleus
- The number of protons equals the number of
electrons in a neutral atom (#p = #e)
-
Atomic Mass Number is the total number of
protons and neutrons in the nucleus
Number of neutrons = mass # - atomic #
And,
Atomic # = #p = #e
Example:
Find the number of protons, electrons and neutrons
for iron and sodium.
Fe
Atomic # = 26
Atomic mass = 55.85 = 56 (round the mass)
Therefore:
# of p = 26
# of e = 26
# of n = 56 – 26 = 30
Note: when finding the number of neutrons we round the
atomic mass to the nearest whole number.
Na
Atomic # = 11
Mass # = 22.99 = 23
p = 11
e = 11
n = 23 – 11 = 12
Au
Atomic # = 79
Atomic mass = 196.96 = 197
p = 79
e = 79
n = 197 – 79 = 118
Lesson 6: Bohr’s Model of the Atom
- Bohr’s model states that electrons can be found
only in certain energy levels or orbits around
the nucleus
- He also stated that only a certain maximum
number of electrons are allowed in each orbit.
Orbit #
1st
2nd
3rd
4th
5th
6th
Max. # of electrons
2
8
8
18
18
32
When one orbit is filled the remaining electrons go to the
next orbit – you cannot exceed the maximum allowed.
We can draw the Bohr diagram for any element. It must
have a nucleus showing the number of protons
and neutrons and circles outside the nucleus
showing the number of electrons.
Reminder: # of protons = # of electrons = atomic #
e.g. Draw the Bohr model for the following elements:
a) Lithium
Step 1 – Look up the atomic number
It’s 3.
So,
# of p = 3
# of e = 3
Step 2 – Look up the atomic mass.
It’s 6.94 = 7 (round to the nearest whole #)
Find the number of neutrons.
Reminder: # of n = atomic mass – atomic #
So, # of n = 7 – 3 = 4
Step 3 – Draw the diagram.
1st orbit
#p=3
#e=3
#n=4
Atomic # =3
Atomic Mass=7
P=3
n=4
The Bohr model diagram can be simplified – we can use lines
instead of circles – this is called Electron Energy Level
Representation.
Electron Energy Level Representations (EELR)
Ex. Zinc - Atomic # = 30
Mass # = 65
Therefore, #p = 30
#e = 30
#n = 65 – 30 = 35
12
8
8
2e
P=30
n =35
These must add up to 30.
Lesson 7: Molecules
and Compounds
Molecule
– a particle formed when two or more atoms combine
* can be 2 atoms which are the same ex. H2, or O2
OR * can be 2 or more different atoms ex. CuS, NaCl, CO
Compound is a pure substance made of 2 or more
different elements
- Compounds can be broken down into simpler substances
- ELECTROLYSIS – use of electricity to separate a chemical
compound into its elements
Ex. Water –broken down into hydrogen and oxygen
- CHEMICAL BONDS hold elements together
- If elements SHARE electrons to form a bond it is
called a MOLECULAR BOND -- MOLECULAR
ELEMENT or MOLECULAR COMPOUND –
Non-metals share electrons to form molecular
compounds.
E.g. CO2
If atoms transfer electrons from one atom to another to
form a bond it is called an IONIC BOND -- IONIC
COMPOUND – metals transfer electrons to non-metals to
form ionic compounds.
E.g. NaCl - Sodium chloride
- CHEMICAL FORMULAS use symbols and numbers
- If only one atom is represented, no numbers are used
- if there is more than one of that type of atom present a
small number written below the line is used to tell us the
number of that type of atom. This is called a SUBSCRIPT.
E.g. H2O - One water molecule is made up of
2 atoms of Hydrogen and 1 Oxygen.
NaCl
- sodium chloride 1 – sodium
1 – chlorine
H2
- hydrogen
C12H22O11
- sucrose
12- carbon
22 – hydrogen
11 – oxygen
Cu(NO3)2
- copper nitrate
1 – copper
2 – nitrogen
6 – oxygen
2 – hydrogen
All pure substances can be identified in two ways:
* Element or compound
* Atom or molecule
NaCl
H2
C6H12O6
Zn
Cu(NO3)2
- compound/molecule
- element/molecule
- compound/molecule
- element/atom
- compounds/molecule
Lesson 8: Molecular Elements and Compounds
Diatomic Elements
– Molecules made of 2 atoms of the same element
- All of the halogens, plus oxygen, hydrogen and
nitrogen are diatomic elements
H2, O2, F2 Cl2, N2, Br2
Molecular Compounds
- Formed when 2 non-metals share electrons
- Most molecular compounds have low melting and
boiling points; therefore they are found as solids,
liquids and gases at room temperature
- They are poor conductors
Naming Binary Molecular Compounds
IUPAC – International Union of Pure and Applied
Chemistry
– Determines how compounds are named followed by chemists around the world
Step 1. Write the entire name of the first element
Step 2. Change the ending on the name of the
second element to – ide
Step 3. Use a prefix to indicate the number of
each type of atom in the formula.
Prefixes are:
1 – mono (only used for the second element)
CO – carbon monoxide
2 – di
3 – tri
4 – tetra
5 – penta
6 – hexa
7 – hepta
8 – octa
9 – nona
10 – deca
Ex. P2O5 –
SiO SCl2 -
diphosphorous pentaoxide
silicon monoxide
sulphur dichloride
NO2 – Nitrogen dioxide
N2O – Dinitrogen monoxide
N2O4 – Dinitrogen tetraoxide
Writing Molecular Formulas
1. Write the symbol for the elements in the same
order as they appear in the name.
2. Use subscripts to indicate the numbers of each
type of atom.
Ex. Carbon tetrabromide
- CBr4
Triarsenic hexasulphide - As3S6
We use small symbols in parentheses after the
formula for each compound to indicate the state of
matter
(s) - solid - NaCl(s)
(l) - liquid - H2O(l)
(g) - gas
- CO2 (g)
Prefixes:
1 = mono
2 = di
3 = tri
4 = tetra
5 = penta
6 = hexa
7 = hepta
8 = octa
9 = nona
10 = deca
Assignment 11:
Name or give the formula:
MEMORIZE
Name_________________________
1. Silicon dioxide ____________________________________
2. Sulphur monoxide_________________________________
3. OF2 _________________________________________
4. SiBr4 ______________________________________
5. PH3 _______________________________________
6. N2O _______________________________________
7. CO ________________________________________
8. NBr3 _______________________________________
9. P2I3 _________________________________________
10. SO3 _________________________________________
11. N2O4 _____________________________________________
12. Tetraphosphorous hexaoxide _________________________________
13. Dinitrogen tetraoxide _______________________________________
14. Heptasilicon monobromide _____________________________________
15. Octaboron decaiodide _________________________________________
16. B2O3 ______________________________________________
17. BrF7 ______________________________________________
18. N3O6 _____________________________________________
19. H2Cl5 _____________________________________________
20. Triselenium diastatide _________________________________
21. Diarsenic pentaoxide _____________________________________
22. Sulphur trioxide _________________________________________
23. C3O2 ___________________________________________________
24. C2H6 ___________________________________________________
25. As3Br7 __________________________________________________
26. SO2 _____________________________________________________
27. Selenium monoxide __________________________________________
28. Diboron trioxide _____________________________________________
29. PF3 ________________________________________________________
30. P2O5 __________________________________________
__________________________________________________
31. P4O10 ________________________________________
32. Arsenic trifluoride ______________________________
33. BrF7 __________________________________________
Lesson 9:
Ionic Compounds
- Atoms that gain or lose electrons to become stable are
called IONS.
- If they gain an electron they have more negative charges
than positive charges so they have a slight negative charge
- Non-metals gain electrons.
- If they lose an electron they have less negative charges
than positive charges so they have a slight positive charge
- Metals lose electrons.
Ionic Compounds
- Made up of a metal bonded to a non-metal – electrons
transfer from the metal to the non-metal.
- All solids at room temperature
- Separate into positive and negative ions when they
dissolve in water
- The ions conduct electricity
Naming Ionic Compounds
1. The name includes both elements in the compound, with
the name of the metallic element first.
2. The non-metallic element is second. Its ending is
changed to – ide.
3. No prefixes used in naming.
Ex. CaCl2
- calcium chloride
(1 calcium/2 chlorine)
Na2S
- sodium sulphide
(2 sodium/1 sulphur)
Fe2O3 - iron(iii) oxide
(rust)
Lesson 10:
Chemical Reactions
Chemical Reactions: formation of a new
substance
Chemical bonds are broken and new bonds
formed
Reactants
Starting materials
Products
become
End materials
Reactants: Any substance that is used up in the reaction
Products: Any substance that is produced in the
reaction
Word equation:
-Gives the names of all the reactants (separated by a +
sign)
-Arrow points to the names of all the products (separ. by +
sign)
Write Word Equations for the following reactions:
(a) When sodium reacts with chlorine, sodium chloride (salt)
is produced.
Sodium + Chlorine
Sodium chloride
(b) Hydrogen gas and zinc chloride are produced when a
piece of zinc metal is dropped into a beaker of hydrochloric
acid.
Zinc + Hydrochloric acid
Hydrogen + Zinc chloride
(c) Potassium iodide is decomposed to produce potassium
metal and iodine.
Potassium iodide
Potassium + Iodine
Law of Conservation of Mass:
In a chemical reaction the total mass of the reactants is
always equal to the total mass of the products.
In a chemical reaction mass is neither gained nor lost.
Molecules may be broken apart and new ones may be
formed, but the atoms in the products are the same ones
that were in the reactants
Ex. Start with 100 grams of reactants - end up with 100
grams of product.
Start out with 20 Hydrogen and 10 oxygen - end up
with 20 hydrogen and 10 oxygen
Using numbers called COEFFICIENTS in front of the
elements and compounds in the reaction balances chemical
reactions.
2 H2O --- 2 H2 + 1 O2
Reaction Rate
– A measure of how fast a reaction occurs
- Some reactions are naturally fast others are slow
We can influence the rate in several ways:
1. Temperature
– The higher the temperature the faster the rate
- Molecules move faster at higher temperatures
- Molecules collide more often and form new substances
more quickly
- The faster the rate, the less time needed for the
reaction
2. Concentration
- Refers to the amount of solute present in a specific
amount of solution - the higher the concentration, the
faster the
reaction
3. Surface Area
- For 2 substances to react, they must come into close
contact - the greater the surface area, the more contact
the two substances have therefore the faster the
reaction
- Can increase surface area by grinding up a chemical
4. Stirring
- Increases the chances of collisions, therefore
speeds up the reaction rate
Lesson 11: Catalysts and Inhibitors
Catalyst
- A substance that speeds up the rate of a
reaction (without being changed itself).
Ex. saliva – acts as a catalyst to break
down starch
Inhibitors
- Substances that slow down chemical reactions
Ex. Added to some foods and medicines to slow down
their decomposition
Corrosion
Corrosion is a chemical reaction. It is the
“eating away” of a metal as it reacts with
other substances in the environment.
Corrosion of iron is called rusting
4 Fe(s) + 3 O2(g)
iron + oxygen
2 Fe2O3(s)
produces rust
Different metals corrode at different rates. Iron
corrodes quickly. Gold does not corrode at all.
Aluminum and copper corrode only on the
surface.
Corroded materials lose their strength.
Rusting is sped up by high temperature,
surface area, and the presence of water, air,
salt or acid.
If the metal is totally protected from air or
water, rusting cannot occur.
Rust Protection
1. Paint.
2. Rust Check – spray with an oil film to keep
air and water away
3. Galvanization - coat it with zinc
- Zinc is more resistant to corrosion
Ex. Galvanized nails
Electroplating – covering a metal with
another metal by using electrolysis
-Ex. Bumpers are coated with a thin layer
of chromium - chromium improves the
hardness, stability, and appearance.
4.
Combustion (Burning)
- Chemical reaction that occurs when oxygen
reacts with a substance to form a new
substance
- Combustion is exothermic
- Oxygen is always one of the reactants – no
oxygen no combustion (no fire).
- Carbon dioxide and water vapor are the
products of combustion when one of the
reactants contains “carbon”
Ex. Methane - CH4
CH4 + 2 O2
Methane + oxygen
CO2 + 2 H2O
Carbon dioxide + water
Identification Tests for the Products of
Combustion
1. Test for carbon dioxide
- Bubble the gas through limewater solution
- If the limewater turns milky, the gas is carbon dioxide
2. Test of water
- Touch the cobalt (II) chloride paper to the liquid
- If the paper turns pink, the liquid is water