Unit 3: Chemical
Processes
Section 1: Chemicals in Action

Chemistry – the study of matter, its
properties, and its changes or
transformations

Matter – anything that has mass and
takes up space

Matter has both physical and chemical
properties
Matter

Matter is either classified as a pure substance or
as a mixture

Pure substance – one in which all the particles
that make up the substance are the same
 E.g. Pure Water

Pure substances can be further classified as
elements or compounds

Elements are pure substances that cannot
be broken down into simpler substances
(made of only one type of atom)
 E.g. oxygen, hydrogen, …

Compounds are pure substances that are
made of two or more different elements in
a fixed proportion
 E.g. carbon dioxide, sodium chloride, …
Properties of Matter

Matter has both physical and chemical properties

Physical property – is a characteristic of a
substance (color, hardness, odour, solubility,
state, melting and boiling point, size)
– A physical change is a change in the size or form of a
substance, which does not change the chemical
properties of the substance

Chemical Property – is a characteristic
behaviour that occurs when a substance
changes to a new substance

A chemical change is also called a chemical
reaction. It involves changing one substance
into one or more different substances with
different properties

The starting materials in a chemical reaction are
called the reactants and the new materials are
called the products.

Chemical tests – are distinctive chemical
reactions that can be used to identify an
unknown substance
Homework

Questions 1, 2, 4, and 5 page 175.
Elements and the Periodic Table

Periodic Table – a structured
arrangement of elements

Elements on the periodic table are
grouped into chemical families – groups
of elements in the same vertical column of
the periodic table. They tend to share
similar physical and chemical properties.
Special Named Families

Alkali metals – group 1 elements

Alkaline earth metals – group 2 elements

Noble gases – group 18 elements

Halogens – group 17 elements
Elements and Atomic Structure

Atoms are composed of three subatomic particles
(protons, neutrons, electrons)

Protons – positively charge particles found in the
nucleus of the atom. Each proton has a mass of 1

Neutrons – neutral particles found in the nucleus of the
atom. Each neutron has a mass of 1

Electrons – negatively charged particles with almost no
mass and can be found in orbits surrounding the nucleus

The number of electrons in an atom is
equal to the number of protons.

Electrons that are in the outer orbit
(valence shell) have greater energy
levels and a greater probability of being
involved in a chemical reaction. These
electrons are called valence electrons.
– Think of Bohr-Rutherford diagrams…
Bohr – Rutherford Diagrams

Each orbit around the nucleus of an atom
contain a specific number of electrons.
– 1st orbit = 2 electrons
– 2nd orbit = 8 electrons
– 3rd orbit = 8 electrons

The electrons fill the orbits in order. Each
orbit must be filled completely before
electrons can be placed in a new orbit.
Sample Bohr – Rutherford
Diagrams

Create a diagram for
Nitrogen
1. Determine the number of
electrons (7)
2. Draw the first orbit with the
maximum number of electrons
allowed (2)
3. Draw the second orbit with the
remaining electrons (5)
4. Draw the nucleus with the
protons and neutrons
7p
7n
You Try It (1)

Create a Bohr – Rutherford diagram for
Magnesium
Answer
Determine the number of electrons, 12
 Draw the first orbit (2)
 Draw the second orbit (8)
 Draw the third orbit (2)
 Draw the nucleus
with the protons and
neutrons
12p

12n
You Try It (2)

Create a diagram for Sodium
Answer





Determine the number of
electrons (11)
Draw the first orbital (2)
Draw the second orbital
(8)
Draw the third orbital (1)
Draw the nucleus with
the protons and neutrons.
You Try It (3)

Create a diagram for Neon
Answer

Determine the number of
electrons (10)

Draw the first orbit (2)
10p

Draw the second orbit (8)

Draw the nucleus with the
protons and neutrons.
10n

The noble gases (column 18) are
considered to be fairly stable.
– Look at the arrangement of the electrons in
the noble gases by drawing a Bohr –
Rutherford diagram for He, Ne, and Ar. What
do you notice about the arrangement of the
electrons?

In some compounds, electrons are
transferred from one atom to another to
create a stable electron arrangement (like
the noble gases)

Create a Bohr – Rutherford Diagram for
Lithium.

Notice that lithium has two
electrons in the first orbit
and one in its outer orbit
(valence shell).

If lithium loses one electron
it will have the same
electron arrangement as
helium.

However the charge on the
atom is no longer neutral. It
has 3 protons (positive
charges) and 2 electrons
(negative charges) giving an
overall charge of 1+
P=3
N=4

An ion is a charged atom in which the
number of electrons is different from the
number of protons

The term ionic charge is used to describe
the overall charge an ion has.
 Example Li which has lost an electron is written as
Li+.

Create a Bohr – Rutherford diagram for
calcium.

How many electrons would calcium have
to lose in order to become stable like a
noble gas?

What would the ionic charge be on Ca?

Metals tend to lose their electrons and
nonmetals tend to gain electrons.
YOU DO IT
What would most likely to happen to the
electrons in fluorine if it was to become
stable like a noble gas?
 What would be the ionic charge on
fluorine if this happened?


Fluorine would gain one electrons
(nonmetal).

Fluorine would have a net ionic charge of
“1-” and the atom would be represented
as “F-” (fluoride ion)
 The ending “ide” is added to nonmetal ions
YOU DO IT

What would most likely to happen to the
electrons in sulfur if it was to become
stable like a noble gas? What would be
the ionic charge on sulfur?

Sulfur would gain two electrons

S2- (sulfide ion)
Homework

1. Create a Bohr – Rutherford diagram for:
–
–
–
–
Boron
Chlorine
Nitrogen
Beryllium

2a. Create a Bohr – Rutherford diagram for the
stable ion formed by each of the above atoms

b. State the ionic charge on each of the ions,
and write the name of the atom in ionic form.
How Elements Form Compounds

Compounds are made by combining elements
together.

Ionic Compound – is formed when positive
and negative ions join together. The ions
combine in a fashion that generally leaves the
ionic compound with a neutral charge.

Molecular Compound – are formed when
nonmetals combine with other nonmetals.
Chemical Formula

A chemical formula is a combination of
symbols that represent a compound
– Example: magnesium chloride (MgCl2)
describes a compound with one magnesium
ion to two chloride ions
Homework

Answer questions 2 and 3 on page 189.
For part b construct a Bohr – Rutherford
Diagram instead of a Bohr diagram.
Ionic Compounds

Metals and nonmetals combine to form
compounds by sharing electrons (metals
lose electrons to form positive ions and
nonmetals gain electrons to form negative
ions)
– Example: Aluminum chloride (AlCl3) Al = 3+
and Cl = 1- (3+) + 3(1-) = 0
Writing Formulas For Ionic Compounds

What is the formula for the ionic compound formed by
calcium and iodine?

Steps:
1.
Write the symbols, with the metal first.
2.
Write the ionic charge above each symbol to indicate the stable ion
that each element forms.
Ca
2+
Ca
3.
I
1I
Determine how many ions of each type you need so that the total
ionic charge is zero.
One Ca2+ ion will balance the charge of two I- ions.
4.
Write the formula using subscripts to indicate the number of ions of
each type
CaI2
YOU TRY IT

What is the formula for the ionic
compound formed by aluminum and
sulfur?
ANSWER

Al2S3
YOU TRY IT 2

What is the formula for the ionic
compound formed by aluminum and
sulfur?
ANSWER

Al2S3
YOU TRY IT 3

What is the formula for the ionic
compound formed by nickel and oxygen?
ANSWER

NiO
Naming Ionic Compounds

The name of the metal is first followed by
the name of the nonmetal. However, the
ending of the name for the nonmetal
changes to “ide”
 Example: Calcium and Iodine become Calcium
iodide
 Example: Aluminum and Sulfur become Aluminum
sulfide
Transition Metal Charges

Sc
Ti
V
Cr
Mn
 Fe
 Co
 Ni
 Cu
 Zn




3+
4+
5+
3+
2+
3+
2+
2+
2+
2+
3+
4+
2+
4+
2+
3+
3+
+
Pd
Ag
 Cd
 Sn
 Sb
 Pt
 Au
 Hg
 Tl
 Pb
 Bi
 Po


2+
4+
+
2+
4+
3+
4+
3+
2+
+
2+
3+
2+
2+
5+
2+
+
+
3+
4+
5+
4+
Ce
 Pr
 Nd
 Pm
 Sm
 Eu
 Gd
 Tb
 Dy
 Ho
 Er
 Tm

3+
3+
3+
3+
3+
3+
3+
3+
3+
3+
3+
3+
Yb
 Lu
 Ac
 Th
 Pa
 U
 Np
 Pu
 Am
 Cm
 Bk
 Cf

2+
3+
3+
3+
4+
5+
6+
5+
4+
3+
3+
3+
3+
4+
6+
Examples

Name or write the chemical formula for the
following molecules.
–
–
–
–

1) Tin (IV) oxide
2)PbO
Platinum (II) sulfide
Ag3N
Answers
–
–
–
–
1) SnO2
Lead (II) oxide
PtS
Silver (I) nitride
Homework

Read page 195

Questions 1, 3 – 9 on page 195
Polyatomic Compounds

Are formed when metals combine with
polyatomic ions (refer to Table 2 on page
196 for a list of polyatomic ions and their
charge)

Polyatomic Ions – groups of atoms that
tend to stay together and carry an overall
ionic charge
Writing Formulas for Polyatomic
Compounds

Rules
1. Write the symbols of the metal and of the
polyatomic group.
2. Write the ionic charges
3. Choose the number of ions to balance the
charge.
4. Write the formula using subscripts
Common Polyatomic Ions












H3O+
NH4+
HCO3HSO4CH3CO2ClO4NO3ClO3NO2ClO2MnO4ClO-
hydronium
ammonium
bicarbonate
hydrogen sulfate
acetate
perchlorate
nitrate
chlorate
nitrite
chlorite
permanganate
hypochlorite













CNOHCO32O22SO42CrO42SO32Cr2O72S2O32HPO42PO43AsO43BO33-
cyanide
hydroxide
carbonate
peroxide
sulfate
chromate
sulfite
dichromate
thiosulfate
hydrogen phosphate
phosphate
arsenate
borate
YOU TRY IT 1

What is the formula for the ionic
compound formed by sodium and sulfate?
ANSWER

Na2SO4
YOU TRY IT 2

What is the formula for the ionic
compound formed by lead (IV) and
carbonate?
ANSWER

Pb(CO3)2
Naming Polyatomic Compounds

The name of a polyatomic compound is
simply a combination of the name of the
metal and the name of the polyatomic ion.
– Example: Potassium ion and carbonate ion
form Potassium carbonate.
Oxyacids

Are compounds formed when hydrogen
combines with polyatomic ions that
contain oxygen.

Table 3 on page 198 list some common
oxyacids.
Naming Acids

Acids containing ions ending with ide often
become hydro -ic acid
– Cl– F– S2-
chloride
fluoride
sulfide
HCl hydrochloric acid
HF hydrofluoric acid
H2S hydrosulfuric acid

Acids containing ions ending with ate usually become -ic
acid
–
–
–
–
–
–
–
–
–
–
CH3CO2CO32BO33NO3SO42ClO4PO43MnO4CrO42ClO3-
acetate
carbonate
borate
nitrate
sulfate
perchlorate
phosphate
permanganate
chromate
chlorate
CH3CO2H
H2CO3
H3BO3
HNO3
H2SO4
HClO4
H3PO4
HMnO4
H2CrO4
HClO3
acetic acid
carbonic acid
boric acid
nitric acid
sulfuric acid
perchloric acid
phosphoric acid
permanganic acid
chromic acid
chloric acid

Acids containing ions ending with ite
usually become -ous acid
–
–
–
–
ClO2NO2SO32ClO-
chlorite
nitrite
sulfite
hypochlorite
HClO2
HNO2
H2SO3
HClO
chlorous acid
nitrous acid
sulfurous acid
hypochlorous acid
Homework

Questions 3 – 4, 6 – 7 on page 198
Molecular Compounds

Most of the compounds you encounter every day
do not contain ions but neutral groups of atoms
called molecules

These compounds form covalent bonds – a
shared pair of electrons between two nonmetal
atoms that holds the atoms together in a
molecule.
– Example: Hydrogen gas is a molecule formed when
two hydrogen atoms combine and share their electron
 Make a diatomic molecule (refer to table 1 on page 202 for a
list of diatomic molecules)
Writing Formulas for Molecular
Compounds

The combining capacity of a nonmetal is a
measure of the number of covalent bonds that
it will need to form a stable molecule

Rules for writing molecular formulas
1. Write the symbols with the left-hand element first
with the combining capacity.
2. Crisscross the combining capacities to produce
subscripts.
3. Reduce the subscripts if possible.
YOU TRY IT 1

How would you write the formula for a
compound formed between carbon and
sulfur?
Answer

CS2
Naming Molecular Compounds

The names of molecular compounds often
contain prefixes that are used to count the
number of atoms when the same two elements
form different combinations.
Prefixes in Molecular Compounds
Prefix
Number
Example (formula)
Mon(o)-
1
Carbon monoxide (CO)
Di-
2
Carbon dioxide (CO2)
Tri-
3
Sulfur trioxide (SO3)
Tetra-
4
Carbon tetrafluoride (CF4)
Pent(a)-
5
Phosphorus pentabromide (PBr5)
Homework

Questions 4 – 6 on page 204

Read pages 205 – 210

What is the difference between a natural
product and a synthetic product?

What is a polymer?