UNIT 2
S PECIFIC L EARNING
O UTCOMES

Relate an element’s position on the periodic table to its
combining capacity (valence). Include: alkali metals,
alkaline earths, chalcogens, halogens, noble gases.

Explain, using the periodic table, how and why elements
combine in specific ratios to form compounds. Include:
ionic bonds, covalent bonds.

Write formulas and names of binary ionic compounds.
Include: IUPAC guidelines and rationale for their use.

Write formulas and names for covalent compounds using
prefixes. Include: mono, di, tri, tetra.

Investigate the Law of Conservation of Mass, and
recognize that mass is conserved in chemical reactions.

Balance chemical equations.

Investigate and classify chemical reactions as synthesis,
decomposition, single displacement, double
displacement, or combustion.

Perform an experiment to classify acids and bases using
their characteristic properties. Include: pH, indicators,
reactivity with metals.

Explain how acids and bases interact to form a salt and
water in the process of neutralization.
R EVIEW
OF THE
ATOM
Atoms can be broken down into subatomic
particles. The three main subatomic particles are
protons, neutrons, and electrons.
Charge
Location
Mass
Protons
+1
Nucleus
1 amu
Electrons
-
Outer shells
Negligible
Neutrons
0
Nucleus
1 amu
amu = atomic mass unit
D ETERMINING THE
N UMBER OF S UBATOMIC
PARTICLES
Number of protons
The number of protons is equal to the atomic number.
Number of electrons
In any atom, the number of electrons is equal to the atomic
number.
Number of neutrons
The number of neutrons is equal to the mass number
(rounded to a whole number) minus the atomic number.
Atomic number
Atomic mass
Protons = 6
Electrons = 6
Neutrons = 12 – 6
=6
Mass number = atomic number (number of protons) + number
of neutrons
Atomic number = mass number – number neutrons
B OHR ’ S ATOMIC M ODEL
Bohr's model of the atom proposed that electrons
occupy orbits or energy levels or shells. Bohr discovered
that the location of each shell was a certain distance
from the nucleus. He also discovered that only a specific
number of electrons populated each shell.
When electrons occupy the shells, they begin at the
closest shell to the nucleus. The first shell can contain a
maximum of two electrons, and the second and third
shells allow a maximum of eight electrons.
Nucleus
Protons
1 p+
Electrons
1 e-
Nucleus
First orbit (max 2e- )
Second orbit (max 8e- )
Third orbit (max 8e- )
L ITHIUM
Nucleus
The third electron in Li
occupies the second
energy level (shell)
F LUORINE
The first 2 electrons are
located in the first shell.
The next 7 electrons occupy
the second shell.
When an element has more than four electrons in its
shell, you arrange the dots in twos.
These electrons are important in helping to determine
how an element will react with other materials. These
electrons in the outermost energy level are called
valence electrons.
T HE P ERIODIC TABLE
Dmitri Mendeleev, a Russian scientist and professor,
placed characteristic properties of elements on pieces of
paper and arranged them in many different ways. He
discovered a certain pattern or repetition of properties.
This repetition of properties or periodicity was an
important outcome of Mendeleev's work. Mendeleev
used the periodicity of properties to create a periodic
table. This table summarized the structure and properties
of the elements.
Today's periodic law states that the properties of elements
are a periodic function of their atomic numbers.
The periodic table is an important tool for chemists. It
quickly determines some key facts about an element. The
basic information with regard to the structure of an atom
is quickly established with the periodic table.
13
+3
Al
26.98
Chemistry : Periodic Table & Elements
C OLUMNS IN THE
P ERIODIC TABLE
The periodic table arranges the elements in columns. A
single column is called a group or family. A family
contains elements that have similar but not identical
properties.
A LKALI M ETALS
(H, L I , N A , K, R B , C S , F R )

Occupies the first column in the periodic table.

Has one valence electron in its outer energy level.

Most reactive metals

In their natural state, alkali metals are always found
combined with other substances because of their
reactivity.
A LKALINE E ARTH M ETALS (B E ,
MG, CA, S R, B A, R A)

Located in the second column of the periodic table.

Less reactive than the alkaline metal family.

Has two valence electrons in the outer energy level.
C HALCOGENS
(O, S, S E , T E , P O )

Located in the 16th column.

Slightly less reactive than the halogen family.

Have six valence electrons in their outer energy level
H ALOGENS
(F, C L , B R , I, AT )

The halogen family is the 17th family

Halogens such as fluorine and chlorine react with one
atom of hydrogen to form HF and HCI respectively.

Have seven valence electrons

Most reactive non-metals.

In their natural state, the highly reactive halogens are
found combined with another element.
N OBLE G ASES
(H E , N E , A R , K R , X E , R N )

The noble gas family is the 18th family

Called noble gases because they do not generally form
compounds with other elements.

Inert, unreactive because their outer energy levels are
completely filled with electrons.

No natural compounds formed from these gases exist.
H YDROGEN

Sometimes hydrogen behaves as a metal and sometimes
as a non-metal.

Hydrogen has one electron in its outermost energy level,
so it is reactive.

Almost all the hydrogen on Earth is combined with other
materials or with itself.
You should be able to predict with some accuracy the
properties of elements that are in the same families.
Look at your periodic table and note the number of
valence electrons in each family. State the number of
valence electrons for the families in the table shown
below.
Family
Number of
Valence Electrons
alkali metals
1
alkaline earth metals
2
chalcogens
6
halogens
7
noble gases
8
R OWS IN THE P ERIODIC
TABLE

Rows in the periodic table are called periods. Elements
in periods do not demonstrate similar properties as
they do in families. Periods, however, show trends.
Video: How to Use the Periodic Table
Video: “The Elements”
E LECTRON D OT D IAGRAM

An electron dot diagram or Lewis diagram represent
an atom and its valence electrons. The electrons in the
valence shell are shown as dots placed around the
symbol. Electron dot diagrams are a valuable tool for
describing, predicting, and explaining compound
formation.
Example:
Draw an electron dot diagram for calcium.
Ca
Example:
Draw an electron dot diagram for sulphur.
S
Example:
Draw an electron dot diagram for nitrogen.
N
Video: Valence Electrons & Lewis Diagrams
Video: How to draw Lewis Diagrams
I ONIC B ONDS

Ionic bonds result when electrons are transferred from
metal atoms to non-metal atoms. The metal atoms lose
electrons to become positive ions, while the non-metal
atoms gain electrons to become negative ions. The ions
are then held together by the action of opposite charges
in an ionic bond.

Atoms are considered neutral because the number of
protons in the nucleus and the number of electrons in
the shells around the nucleus are equal, resulting in a
zero net charge for the atom.
C ATIONS (P OSITIVE I ONS )
What happens if an electron is removed from an atom?
Sodium atom
Sodium ion
How many positive charges in
the atom?
How many positive charges in
the ion? 11
11
How many negative charges in
the atom?
How many negative charges in
the ion?
What is the residual (net)
charge in the atom? 0
What is the residual (net)
charge in the ion? 11 – 10 = +1
11
10

When an atom from the alkali family reacts with an
atom from another element, it will give off its valence
electron to the other atom. In giving away its electron,
the alkali metal atom has a filled outer shell. Also, in
giving away its electron, the atom becomes positively
charged with a 1+ charge.
Na loses this electron
Na
You do not have
to write the 1
Na

Atoms from the alkaline earth family have two
valence electrons. These atoms will give off the two
valence electrons when combining with an atom
from another element. After giving away the two
valence electrons, the alkaline earth metal atom has
a 2+ charge.
Mg loses these electrons
2+
Mg
Mg
To get charge: 0 – (-2) = +2
Electrons lost

Alkali and alkaline earth metals form positive ions
when forming an ionic bond with another element.
Always write the charge with the number first, followed
by the sign ie.) 1+
A NIONS (N EGATIVE I ONS )

The oxygen family and the halogen family behave
differently. These two families readily accept
electrons to fill their valence shells so that their
electron configuration also resembles a noble gas.
The halogens need only one electron to fill their
valence shells, so they accept only one electron.
Once the extra electron is accepted, a 1- ion is
formed.
O XYGEN
To get charge: 0 + (-2) = -2
As 2 electrons are placed
into the valence shell,
the oxygen atom
becomes a 2- ion.
8p
8n

Oxygen has a combining capacity of two. The
chalcogen family all have the same combining
capacity.
F LUORINE
To get charge: 0 + (-1) = -1
As the electron is placed
into the valence shell,
the fluorine atom
becomes a 1- ion.
9p
10 n

Fluorine has a combining capacity of one. The
halogen family all have the same combining capacity.

When a positively charged ion comes near a negatively
charged ion, they attract each other and form a bond
called an ionic bond. An ionic bond will hold the two
ions together to form a compound.

The formation of compounds often takes place
vigorously when metals and non-metals are placed
together. If a sample of sodium metal is placed in a
container of chlorine gas, an explosive reaction takes
place and the sodium combines with the chlorine to
form sodium chloride. The Bohr model below provides
an explanation for this reaction.
Electron from Na has been transferred to Cl
Sodium
Chlorine
11 p
12 n
17 p
18 n
An electrostatic attraction holds the ions together.
The cation and anion are attracted to each other.
This is what forms an ionic bond.
+
11 p
12 n
17 p
18 n
-

Notice that atoms when combine to form ionic
compounds, they always gain or lose enough
electrons to have a valence shell like its closest noble
gas neighbour.
Ionic Bonds
S TRUCTURE
OF
N AC L
M ETALS AND N ON - METALS
Metals
o Generally have three or fewer valence electrons
o Alkali family of elements has the strongest metallic
properties
o The next strongest metallic properties are found in the
alkaline Earth family.
o Metallic properties exist but in decreasing amounts
through the boron family.
M ETALS AND N ON - METALS
Non-metals
o Have five or more electrons in the valence shell.
o Begin at the chalcogen family and end at the noble gas
family.

There are exceptions to this general classification and
some families have members that behave as both
metals and non-metals (e.g., silicon). These elements
are called metalloids. Many periodic tables have stair
steps across families at the right side. These steps show
the dividing line between metals and non-metals. The
elements on the dividing line are metalloids.
State
Metals
Appearance
o solids at room temp., o shiny lustre
except for Hg (liquid)
Non-metals o some gases at room
temp.
o some solids
o one liquid (bromine)
o not very shiny
Metalloids o solids at room
temperature
o can be shiny or
dull
Conductivity
Metals
o good conductors of
heat and electricity
Malleability and
ductility
o malleable
o ductile
Non-metals o poor conductors of
heat and electricity
o brittle
o not ductile
Metalloids o may conduct
electricity
o poor conductors of
heat
o brittle
o not ductile
R EVIEW OF VALENCE E LECTRONS
Family
Valence
Electrons
Electrons
Lost/Gained
Charge
of Ions
Alkaline metals
1
Lose 1
1+
2
Lose 2
2+
3
Lose 3
3+
Carbon
4
Lose or gain 4
4+/-
Nitrogen
5
Gain 3
3-
Chalcogen
6
Gain 2
2-
Halogen
7
Gain 1
1-
Noble gas
8
-----------
0
Alkaline earth
metals
Boron

When ionic compounds are formed, elements with a
positive valence number will combine with elements
having a negative valence number.

In general terms, metals (families 1 and 2) combine
with non-metals (families 16 and 17).
C OVALENT B ONDS

Many compounds do not form ionic bonds. These
compounds contain two or more non-metallic atoms.
For example, C02 is made of two different non-metals,
carbon and oxygen. These compounds are formed
through the sharing of valence electrons. A covalent
bond is formed when two or more non-metallic atoms
share valence electrons.

Two hydrogen atoms form a covalent bond by sharing
electrons to produce a hydrogen molecule.

A molecule is the smallest unit of a covalent compound.

A molecule has different characteristic properties from
the atoms that form it.

The Bohr model for hydrogen shown below illustrates a
covalent bond.
H H
Each line is a bond
which shares 2e-
Covalent bond

The electrons are shared in the outer shells of both
atoms. This covalent bond forms a molecule of
hydrogen (H2). The two hydrogen atoms form a
diatomic molecule.
Covalent Bonds
See white board for example
A basic rule in chemistry is that an atom with eight
electrons in its outer shell is particularly stable. This
need for eight electrons in a covalent bond is called the
octet rule.
Isn’t it Ionic Song
Chemical Bonds Song
Covalent vs Ionic
How to Determine Ionic & Covalent Bonds
D IATOMIC M OLECULES
Name of
Element
Symbol for One Atom of
the Element
Formula for One
Molecule
Hydrogen
H
H2
Nitrogen
N
N2
Oxygen
O
O2
Fluorine
F
F2
Chlorine
Cl
Cl2
Bromine
Br
Br2
I
I2
Iodine

The elements forming diatomic gases are unstable as
single atoms and combine almost instantaneously to
form stable molecules.

Diatomic molecules are still classified as elements
even though they are molecules. Remember that
diatomic molecules are made of only one kind of
atom.
C HEMICAL F ORMULAS

Chemistry has its own language. Chemists
communicate in this language to describe the millions
of known compounds. This communication depends
on a standard system of naming and writing the
formulas for compounds. Chemists formed a group to
standardize the system of naming and called
themselves the International Union of Physical and
Applied Chemists, or IUPAC.

A chemical formula is a shorthand method to represent
compounds that uses the elements' symbols and
subscripts. The chemical formula gives the following
information:

The different elements in the compound.

The number of atoms of each element in the
compound.
element
symbols
H2O
subscript
•Subscript tells you amount of each element.
•Water contains:
•2 Hydrogens
•1 Oxygen
* No subscript indicates only 1 atom is present*
Na2SO4
•Contains:
•2 sodium atoms
•1 sulphur atom
•4 oxygen atoms
Types of Chemical Formulas
N AMING I ONIC
C OMPOUNDS
When naming an ionic compound from its formula follow
the rules below:
1.
The cation (positive ion) is named first, followed by the anion
(negative ion).
2.
Write the full name of the metallic element (positive ion).
3.
Write the name of the non-metallic element (negative ion) and
change the ending to "-ide".
Example:
Write the name of NaCl.
Step 1: Name the first element.
Na = sodium
Step 2: Name the second element and change the ending
to "-ide".
Cl = chlorine  chloride
sodium chloride
The name of the compound is ______________________.
Example:
Write the name of Mg3P2.
Step 1: Name the first element.
Mg = magnesium
Step 2: Name the root of the second element and add
"-ide".
P = phosphorus  phosphide
magnesium phosphide
The name of the compound is _____________________.
W RITING I ONIC F ORMULAS
The following must occur, when writing the formula for
ionic compounds.
1. The formula must have the cation first, followed by
the anion.
2. The sum of the charges of the ions must be zero.
That is, the number of positive charges must equal
the number of negative charges.
3. You may not change the charge of the ions to make
the ion charges equal zero.
T HE “C ROSS -O VER ”
M ETHOD
1. Write the ions and their charges side by side.
2. Make the number of the charge of one ion the
subscript of the other ion (omitting the + or – sign).
Remember we do not write the number one as a
subscript.
3. Reduce all subscripts to their simplest form, if
necessary.
Example:
Write the formula for aluminum oxide.
Step 1: Write the ions and their charges.
Al3+ and O2Step 2: Make the number of the charge of one ion the
subscript of the other ion.
Al3+ O2Al2O3
Example:
Write the formula for barium fluoride.
2+
Ba
1F
BaF2
The charge on the fluoride ion is 1-. Since IUPAC rules do
not write the number one as a subscript, we leave the
barium without a subscript.
Example:
Write the formula for magnesium chloride.
See white board
Example:
Write the formula for calcium oxide.
See white board
Naming Ionic Compounds
P OLYATOMIC I ONS

Some ions are composed of several atoms joined
covalently. These are called polyatomic ions
(poly = many).

The charge for polyatomic ions is for the whole group
of atoms not just for the atom written last. DO NOT
change the subscripts of polyatomic ions; if you
change the subscripts you change the identity of these
ions.

When indicating the presence of more than one
polyatomic ion in a compound, we use parenthesis
around the polyatomic ion, followed by its subscript.
For example, the compound Al(CH3COO)3 has an
aluminum ion and 3 acetate ions. Placing the acetate
ion in parenthesis and following it with the subscript 3
indicates there are 3 acetate ions.
Example:
Write the name for KNO3.
Step 1: Identify the cation.
K+  potassium ion
Step 2: Identify the anion.
NO3-  nitrate ion
Step 3: Write the name of the cation first, followed by the
anion. potassium nitrate
Example:
Write the name of Na3PO4.
Step 1: Identify the cation.
Na+  sodium ion
Step 2: Identify the anion.
PO43-  phosphate ion
Step 3: Write the name of the cation first, followed by the
anion.
sodium phosphate
Writing the formula for polyatomic ions is the same
as writing the formula for ionic compounds. You will
use the cross-over method and your polyatomic
table.
Example:
Write the name of sodium sulfate.
Step 1: Identify the cation.
sodium  Na+
Step 2: Identify the anion.
sulfate  SO42Step 3: Cross-over the charges to write the formula.
+
Na
2SO4
Na2SO4
Example:
Write the name of ammonium thiocyanate.
See white board
S TOCK N AMING S YSTEM

Most of the transition metals have more than one
possible ion charge. They are often referred to as being
multivalent. For example,
Ion
Possible Ion Charges
Copper
Iron
Cobalt
Chromium
1+, 2+
2+, 3+
2+, 3+
2+, 3+
Lead
Tin
2+, 4+
2+, 4+

In 1919, Alfred Stock (1876 – 1946), a German chemist,
suggested using numbers to indicate the charge of the
ions. Prior to this the ions were given different names
based upon their charge.

The Cu+ ion was called cuprous and the Cu2+ ion was
called cupric. However, the Fe2+ ion was ferrous and the
Fe3+ ion was ferric.

Since the charges were not always the same, the "–ic"
and "–ous" suffixes caused some confusion.

Today, the Stock naming system uses Roman numerals
following the metal ion's name to indicate the ion's
charge.
Example:
Copper (I) = Cu+
Copper (II) = Cu2+
cuprous
cupric
Iron (II) = Fe2+ ferrous
Iron (III) = Fe3+ ferric
As a general rule, all metals are multivalent (have more
than one ion charge) except group one and two metals,
silver, cadmium, zinc, and aluminum. Unless the metal is
one of these use the Roman numeral.
Example:
Write the formula for iron (III) chloride
Step 1: Write out the ions.
Fe3+ and ClStep 2: Cross-over the charges.
FeCl3
Example:
Write the formula for lead (IV) sulfide.
Step 1: Write the ions.
Pb4+ and S2Step 2: Cross-over the charges.
Pb2S4
Step 3: Reduce the subscripts.
Pb2S4 ÷ 2  PbS2
*divide by GCF*
N AMING M ULTIVALENT
C OMPOUNDS

We name in a very similar manner as those ions with a
single ion charge, except we must determine the
charge on the metal ion.
See overhead for examples
N AMING C OVALENT
C OMPOUNDS

Non-metals tend to combine chemically by sharing
electron pairs. These bonds are known as covalent
bonds. Neutral compounds made of atoms joined
covalently are called molecular or covalent
compounds.

We name covalent compounds differently than ionic
compounds. We must indicate the number of each
element by adding a prefix in front of the element's
name.
Subscript
Prefix
one
two
three
four
five
six
seven
eight
nine
ten
mono
di
tri
tetra
penta
hexa
hepta
octa
nona
deca

There are three exceptions to the naming rules. Here
the common names for the compounds are used:
H2O = water
NH3 = ammonia
CH4 = methane
Example:
Write the name for CO2.
Step 1: Name the first atom with prefixes.
There is only 1 carbon. We omit “mono” for the first element.
carbon
Step 2: Name the second element using prefixes and end in "-ide".
There are 2 oxygens, so we use the di prefix
dioxide
Step 3: Write the name of the compound writing the substance
found more to the left on the periodic table first.
carbon dioxide
Example:
Write the name for N2O4.
dinitrogen tetraoxide
Example:
Write the name for SF6.
sulfur hexafluoride
W RITING C OVALENT
F ORMULAS
Writing formulas for covalent compounds involves the
following rules:
1.
Write the symbol for the first element followed by the
subscript indicated by the prefix.
2.
Write the symbol of the second element followed by the
subscript indicated by its prefix.
DO NOT REDUCE THE SUBSCRIPTS!!!
Example:
Write the formula for dinitrogen monoxide.
Step 1: Write the symbol and subscript for the first
element.
dinitrogen  N2
Step 2: Write the symbol and subscript for the second
element.
1 as a subscript is
monoxide  O
not needed
Step 3: Combine
N2O
Example:
Write the formula for sulphur hexachloride.
SCl6
Example:
Write the formula for carbon tetrachloride.
CCl4
Naming Covalent Compounds 1
How to Name Ionic & Covalent Compounds
Naming Covalent Compounds 2
B ALANCING E QUATIONS

The Law of Conservation of Mass tells us that every
chemical equation must have equal numbers of
atoms of each element on each side of the equation.

This means atoms cannot be created or destroyed in a
chemical reaction.
Four atoms of
hydrogen
Two atoms of
oxygen
Four atoms of
hydrogen and two
atoms of oxygen.
As you balance equations, there are several rules to
remember
1. You cannot change the formula of any reactant or
product to change the numbers of atoms.
2. You can only change the coefficients in front of the
reactants and products.
a. Coefficients can be placed only in front of the formula,
not somewhere inside it
b. Coefficients apply to the whole molecule. If no
coefficient is shown in front of a molecule, it means
the molecule has a coefficient of 1.
c. Note the difference between a coefficient and a
subscript.
i.
A coefficient tells us the number of molecules in an
equation and can be changed to balance an equation.
ii. A subscript tells us the number of atoms in a
molecule and cannot be changed.

There are some rules that you can use when you are
given an unbalanced
H2(g) + O2(g)  H2O(l)
1. Determine the number of atoms for each element in
the molecules for reactants and products.
The number of atoms on each side of the
equation are not the same.
2.
If the numbers of atoms on both sides of the equation
are equal at this point, the equation is already
balanced and you are finished. In this example, they
are not the same and you go to the next step.
3.
Choose the substance that has the most influence on
the equation and insert coefficients in the formulas as
needed.
4.
Inspect the equation and recalculate the numbers of
atoms on both sides of the equation. If they are equal,
the equation is balanced. If they are not equal, change
the coefficients until the equation is balanced.
A Beginner’s Guide to Balancing Equations
How to Balance Equations
How to Write, Balance & Classify a Chemical Reaction
How to Write Chemical Equations
C HEMICAL R EACTIONS
In chemistry, a reaction happens when two or more
molecules interact and the molecules change.
S INGLE D ISPLACEMENT R EACTIONS
In single displacement reactions, one element replaces
another element in a compound. There are two
possible reactions:
a) One positive ion replaces another
Zn + HCI  ZnCl2 + H2
b) One negative ion replaces another
Cl2 + 2NaBr  2NaCl + Br2
A + BC  B + AC
S YNTHESIS R EACTIONS
Synthesis reactions involve joining atoms to make a
molecule, or joining elements to form a compound.
Two or more simple elements or compounds combine
to form a more complex compound.
2Mg + O2  2MgO
A + B  AB
D ECOMPOSITION R EACTIONS
During decomposition, one compound splits apart into
two (or more) pieces.
sodium chloride  sodium + chlorine
2HgO 2Hg + 02
AB  A + B
D OUBLE D ISPLACEMENT R EACTIONS
In double displacement, elements in different
compounds replace each other (split up and exchange
partners).
AC + BD  AD + BC
C OMBUSTION R EACTIONS
Combustion reactions involve the burning of a
chemical substance in oxygen.
Metals burn in oxygen to produce a metal oxide.
Magnesium + oxygen  magnesium oxide
Nonmetals burn in oxygen to produce a dioxide.
Carbon + oxygen  carbon dioxide
Hydrocarbon + O2  CO2 + H20
Bill Nye Chemical Reactions
Chemical Reactions
Exothermic and Endothermic Reactions
7 Chemical Reactions that will Facinate
C OMMON A CIDS AND
B ASES

Acids are used in many industrial processes. Production
of paper, steel, and many other products requires the
use of acids. Sulphuric and hydrochloric acids are
commonly used in such industrial processes.

Bases are important household chemicals used for
cleaning and disinfecting. Bases are also used in
hairdressing when a "permanent" change is needed to
change straight hair to curly hair.
The names and formulas of some common acids and
bases are shown below.
Acids
Bases
hydrochloric acid HCl
sodium hydroxide
NaOH
sulphuric acid
H2SO4
Ca(OH)2
nitric acid
HNO3
calcium hydroxide
ammonium
hydroxide
NH4OH
Notice that acids tend to have H in their formula and bases
often have OH (hydroxide) as the second component in
their formula.
C HARACTERISTICS OF
A CIDS AND B ASES
Taste
Feel
Reaction to
Litmus Paper
Other
Properties
Acids
Bases
sour
burns
turns litmus
paper red
bitter
slippery
turns litmus
paper blue
•neutralizes basic
•neutralizes acidic
solutions
solutions
•corrosive to metal •conduct electricity
•conduct electricity
Acid + bases  water +
salt
Reactions Acid + metals  H2 gas
Acid + carbonates 
CO2 + H2O + salt
Bases + acids  water +
salt
•HCl – gastric juice
•citric acid –lemons &
oranges
•acetic acid – vinegar
•ascorbic - vitamin C
Examples acid
•acetylsalicylic - Aspirin
acid
•NaOH – sodium
hydroxide
•Ca(OH)2 – limewater
•Mg(OH)2 – milk of
magnesia
•bleaches
•soaps
•toothpaste

A substance that changes colour when added to an
acid or base is called an indicator.
Indicator
Acid
Base
Litmus
red
blue
Phenolphthalein
colourless
pink
Methyl Orange
orange
Bromothymol
Blue
blue
N EUTRALIZATION

An acid and a base, when combined, will neutralize
each other. Acids will lose their acid properties and
bases will lose their base properties.

When an acid reacts with a base, a salt and water are
produced.
acid + base  salt + water

A salt is a compound composed of the negative ion of
an acid and the positive ion of the base. The water is
formed when the hydrogen ion (H+) of the acid
combines with the hydroxide ion (OH-) of the base.
acid
base
salt
water
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
M EASURING THE S TRENGTH
OF ACIDS AND B ASES

The strength of an acid or a base is measured using a
scale called pH.
pH 0-6.9
acidic
pH 7 (water)
neutral
pH 7.1-14
basic (alkaline)
Measures the strength of acid where a smaller
number equals stronger acid and a larger
number equals weaker acid
Neutral - neither acid nor base
Measures the strength of base where smaller
number equals weaker base and larger
number equals stronger base
strong
weak
weak
strong
Question:
Use the pH values described below to determine whether
the substance is neutral, a weak or strong acid, or a weak
or strong base.
pH = 12
strong base
pH = 7
neutral
pH = 6
weak acid
pH = 9
weak base
pH = 2
strong acid
Acids, Bases and pH