CHAPTER 6
Ionic Bonding
© 2013 Marshall Cavendish International (Singapore) Private Limited
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Chapter 6
Ionic Bonding
6.1 The Stable Electronic Configuration of
a Noble Gas
6.2 Forming Ions
6.3 Ionic Bond: Transferring Electrons
6.4 Chemical Formulae of Ionic Compounds
6.5 Structure and Physical Properties of Ionic
Compounds
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6.1
The Stable Electronic Configuration
of a Noble Gas
Learning Outcome
At the end of this section, you should be able to:
• describe the stable electronic configuration of
a noble gas.
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6.1
The Stable Electronic Configuration
of a Noble Gas
What are Noble Gases?
• Elements that belong to Group 0 of the
Periodic Table
• Examples: He, Ne, Ar, Kr and Rn
• Atoms of noble gases are stable and
unreactive.
• They exist in nature as single atoms.
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6.1
The Stable Electronic Configuration
of a Noble Gas
What is the Noble Gas Structure?
• Noble gases have full or complete outer shells.
Helium has a duplet configuration
(2 outer electrons).
All other noble gases have an octet configuration
(8 outer electrons).
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6.1
The Stable Electronic Configuration
of a Noble Gas
Why Do Atoms React?
• Atoms of most other elements are reactive
because they do not have the noble gas structure
(i.e. their outer shells are not fully-filled).
• Atoms of these elements lose, gain or share
outer electrons to attain the noble gas
configuration and form compounds.
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6.1
The Stable Electronic Configuration
of a Noble Gas
Chemical Bonding
Atoms gain or lose
electrons to attain
noble gas configuration
Ionic bonding
Atoms share electrons
to attain noble gas
configuration
Covalent bonding
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Chapter 6
Ionic Bonding
6.1 The Stable Electronic Configuration of a
Noble Gas
6.2 Forming Ions
6.3 Ionic Bond: Transferring Electrons
6.4 Chemical Formulae of Ionic Compounds
6.5 Structure and Physical Properties of Ionic
Compounds
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6.2 Forming Ions
Learning Outcome
At the end of this section, you should be able to:
• describe the formation of positive ions (cations)
and negative ions (anions) to achieve the noble
gas configuration.
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6.2 Forming Ions
What is an Ion?
Recall:
Atoms have an equal number of
protons and electrons. They are
electrically neutral.
• An atom loses or gains electrons to form ions.
• Ions are charged particles.
No. of electrons ≠ No. of protons
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6.2 Forming Ions
What is an Ion?
URL
•
Ions can be positively- or negatively-charged.
•
Positively-charged ions are called cations.
•
Negatively-charged ions are called anions.
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6.2 Forming Ions
Formation of Cations
•
Atoms of metals lose electrons to form
positively-charged ions called cations.
•
In this way, they achieve the noble gas
configuration.
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6.2 Forming Ions
Example 1: Formation of sodium (Na+) ion
Na atom
Electronic configuration: 2, 8, 1
Number of protons = 11
Number of electrons = 11
The Na atom loses one outer electron to form the
Na+ ion. Why?
To achieve stable octet (noble gas) configuration.
Neon (2, 8)
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6.2 Forming Ions
Example 1: Formation of sodium (Na+) ion
2, 8
2, 8, 1
+
sodium atom loses one
outer electron
Na atom: 11p, 12n, 11e
Charge = 11p + 11e
= (+11) + (–11)
=0
Neutral
Na atom
Na+ ion: 11p, 12n, 10e
Charge = 11p + 10e
= (+11) + (–10)
= +1
Positively-charged
Na+ ion
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6.2 Forming Ions
Example 2: Formation of calcium (Ca2+) ion
2, 8, 8
2, 8, 8, 2
calcium atom loses
two outer electrons
Ca atom: 20p, 20n, 20e
Charge = 20p + 20e
= 20(+1) + 20(–1)
= (+20) + (–20) = 0
Neutral
Ca atom
2+
Ca2+ ion: 20p, 20n, 18e
Charge = 20p + 18e
= 20(+1) + 18(–1)
= (+20) + (–18) = +2
Positively-charged
Ca2+ ion
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6.2 Forming Ions
Common Cations and Their Charges
Metal
Ion
Formula of ion
sodium
sodium ion
Na+
potassium
potassium ion
K+
calcium
calcium ion
Ca2+
magnesium
magnesium ion
Mg2+
aluminium
aluminium ion
Al3+
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6.2 Forming Ions
Formation of Anions
• Atoms of non-metals gain electrons to form
negatively-charged ions called anions.
• In this way, they achieve the noble gas
configuration.
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6.2 Forming Ions
Example 1: Formation of chloride (Cl–) ion
Cl atom
Electronic configuration: 2, 8, 7
Number of protons = 17
Number of electrons = 17
What happens in the formation of a chloride ion?
The chlorine atom gains one electron in its outer shell to
achieve a stable octet (noble gas) configuration.
Argon (2, 8, 8)
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6.2 Forming Ions
Example 1: Formation of chloride (Cl–) ion
2, 8, 7
2, 8, 8
chlorine atom gains
one electron
Cl atom: 17p, 18n, 17e
Charge = 17p + 17e
= (+17) + (–17)
=0
Neutral
Cl atom
Cl– ion: 17p, 18n, 18e
Charge = 17p + 18e
= (+17) + (–18)
= –1
Negatively charged
Cl– ion
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6.2 Forming Ions
Example 2: Formation of oxide (O2–) ion
2, 6
2, 8
oxygen atom gains
two electrons
2–
O atom: 8p, 8n, 8e
O2– ion: 8p, 8n, 10e
Charge = 8p + 8e
= (+8) + (–8)
=0
Charge = 8p + 10e
= (+8) + (–10)
= –2
Neutral
O atom
Negatively charged
O2– ion
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6.2 Forming Ions
Common Anions and Their Charges
Non-metal
Ion
Formula of ion
chlorine
chloride ion
Cl–
bromine
bromide ion
Br–
oxygen
oxide ion
O2–
sulfur
sulfide ion
S2–
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Why do metals lose electrons to
form positive ions (cations) but
non-metals gain electrons to form
negative ions (anions)?
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Chapter 6
Ionic Bonding
6.1 The Stable Electronic Configuration of a
Noble Gas
6.2 Forming Ions
6.3 Ionic Bond: Transferring Electrons
6.4 Chemical Formulae of Ionic Compounds
6.5 Structure and Physical Properties of Ionic
Compounds
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6.3 Ionic Bond: Transferring Electrons
Learning Outcome
At the end of this section, you should be able to:
• describe how an ionic bonds are formed
between metals and non-metals.
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6.3 Ionic Bond: Transferring Electrons
Ionic Bonding
• Ionic bonds are formed between metals and
non-metals.
Examples:
Group VII: Fluorine, chlorine
Group VI: Oxygen, sulfur
Examples:
Group I: Sodium, potassium
Group II: Magnesium, calcium
• This is done through the transfer of electron(s)
from metals to non-metals.
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6.3 Ionic Bond: Transferring Electrons
Ionic Bonding
Metallic atom
loses electron(s)
Non-metallic atom
gains electron(s)
Positive ion
Negative ion
(cation)
(anion)
electrostatic forces of
attraction
(hold oppositely charged
ions together)
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6.3 Ionic Bond: Transferring Electrons
Formation of Ionic Compound
Example 1: Sodium chloride
Step 1: Formation of Positive Ions
Each sodium atom (Na) loses its single outer electron
to form a positively-charged sodium ion (Na+).
Na
2, 8, 1
Na+ + e−
2, 8
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6.3 Ionic Bond: Transferring Electrons
Step 2: Formation of Negative Ions
Each chlorine atom gains an electron from a
sodium atom to form a negatively-charged
chloride ion (Cl−).
+
Cl
2, 8, 7
e−
Cl –
2, 8, 8
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6.3 Ionic Bond: Transferring Electrons
Step 3: Formation of Ionic Bonds
Loses
one electron
Sodium
atom
2, 8, 1
Gains
one electron
Chlorine atom
2, 8, 7
Electrostatic forces
of attraction
Sodium ion Chloride ion
2, 8, 8
2, 8
Sodium and chlorine react in the ratio of 1 : 1 to form
sodium chloride (NaCl).
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URL
6.3 Ionic Bond: Transferring Electrons
Example 2: Magnesium chloride
Magnesium
atom loses two
electrons.
Chlorine atoms
gain one electron each.
Chloride ion Magnesium Chloride ion
2, 8, 8
2, 8, 8
ion
2, 8
Magnesium reacts with chlorine in the ratio of 1 : 2
to form magnesium chloride (MgCl2).
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Chapter 6
Ionic Bonding
6.1 The Stable Electronic Configuration of a
Noble Gas
6.2 Forming Ions
6.3 Ionic Bond: Transferring Electrons
6.4 Chemical Formulae of Ionic Compounds
6.5 Structure and Physical Properties of Ionic
Compounds
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6.4 Chemical Formulae of Ionic Compounds
Learning Outcome
At the end of this section, you should be able to:
• deduce the chemical formula of an ionic compound
from the charges on the ions and vice versa.
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6.4 Chemical Formulae of Ionic Compounds
Chemical Formulae of Ionic Compounds
• The formula of an ionic compound is
constructed by balancing the charges on the
positive and negative ions.
• All the positive charges must equal all the
negative charges in an ionic compound.
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6.4 Chemical Formulae of Ionic Compounds
Example: Magnesium oxide
Oxygen forms
O2− ions.
Magnesium forms
Mg2+ ions.
Mg2+
Charge: +2
O2−
Charge: −2
Since 1 × (+2 charge) balances out 1 × (−2 charge),
The formula is MgO.
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6.4 Chemical Formulae of Ionic Compounds
Example: Copper(II) hydroxide
Copper ion
Hydroxide ion
Cu2+
Charge: +2
OH−
Charge: −1
To balance the charges, multiply the smaller charge (−1) by
2 to make it equal to +2.
Since 1 × (+2 charge) balances out 2 × (−1 charge),
The formula is Cu(OH)2.
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6.4 Chemical Formulae of Ionic Compounds
Example 1
Write the chemical formula of aluminium oxide.
oxide ion
aluminium ion
O 2−
Al 3 +
Charge: −2
Charge: +3
Al2O3
Therefore, the formula is Al2O3.
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6.4 Chemical Formulae of Ionic Compounds
Example 2
Write the chemical formula of calcium carbonate.
calcium ion
Ca 2 +
carbonate ion
CO3 2 −
Charge: +3
Ca2(CO3)2
Charge: −2
CaCO3
Since ‘2’ is a common factor, it can be removed.
Therefore, the formula is CaCO3.
URL
37
Chapter 6
Ionic Bonding
6.1 The Stable Electronic Configuration of a
Noble Gas
6.2 Forming Ions
6.3 Ionic Bond: Transferring Electrons
6.4 Chemical Formulae of Ionic Compounds
6.5 Structure and Physical Properties of
Ionic Compounds
38
6.5
Structure and Physical Properties of
Ionic Compounds
Learning Outcomes
At the end of this section, you should be able to:
• state that ionic compounds form giant lattice
structures;
• deduce the formulae of ionic compounds from their
lattice structures;
• relate the physical properties of ionic compounds to
their lattice structures.
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6.5
Structure and Physical Properties of
Ionic Compounds
Structure of Ionic Compounds
Ionic compounds form giant ionic structures.
Also known as giant lattice structures or
crystal lattices
Consist of an endlessly repeating three-dimensional
lattice of positive and negative ions
Ions are closely packed, arranged in an orderly
manner and held in place by ionic bonds
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6.5
Structure and Physical Properties of
Ionic Compounds
Structure of NaCl
Sodium chloride
crystal
Three-dimensional
arrangement of sodium
ions and chloride ions
Sodium ions and chloride ions
alternate with each other.
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6.5
Structure and Physical Properties of
Ionic Compounds
Structure of NaCl
Strong forces of attraction
between ions in crystal
lattice
A large amount of energy
is required to overcome
these forces of attraction
between ions.
Na+
Cl–
Cl– Na+
Na+
Cl– Na+
Cl– Na+
Cl– Na+
Cl– Na+
Na+
–
Cl Na+
Cl– Na+
–
Cl–
Cl Na+
–
Cl Na+
Cl– Na+
+
Na Cl–
Na+ Cl–
Na+
Cl–
Cl– Na+
–
Cl Na+
Cl– Na+
+
Na Cl–
Na+ Cl–
Na+ Cl–
Cl–
Na+
Cl–
Na+
Cl–
Na+
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6.5
Structure and Physical Properties of
Ionic Compounds
Structure of NaCl
Cl− ion
Each chloride ion is
surrounded by six sodium
ions.
Na+ ion
Each sodium ion is
surrounded by six chloride
ions.
The ratio of sodium ions to chloride ions is 1 : 1.
Hence, the formula unit of sodium chloride is NaCl.
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6.5
Structure and Physical Properties of
Ionic Compounds
Melting and Boiling Points of
Ionic Compounds
• High melting and boiling
points
• Non-volatile
• Exist as solids at room
temperature
Na+
Cl–
Cl– Na+
Na+
Cl– Na+
Cl– Na+
Cl– Na+
Cl– Na+
Na+
–
Cl Na+
Cl– Na+
–
Cl–
Cl Na+
–
Cl Na+
Cl– Na+
+
Na Cl–
Na+ Cl–
Na+
Cl–
Cl– Na+
–
Cl Na+
Cl– Na+
Na+ Cl–
Na+ Cl–
Na+ Cl–
Cl–
Na+
Cl–
Na+
Cl–
Na+
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6.5
Structure and Physical Properties of
Ionic Compounds
Solubility of Ionic Compounds
• Usually soluble in water
Cl–
Na+ Cl– Na+
Na+
dissolve in water
Cl– Na+ Cl–
Na+ Cl– Na+
Na+
Water
molecules
•
URL
Cl–
Usually insoluble in organic solvents
E.g. ethanol, turpentine, petrol
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6.5
Structure and Physical Properties of
Ionic Compounds
Electrical Conductivity of Ionic Compounds
solid NaCl
aqueous NaCl
molten NaCl
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6.5
Structure and Physical Properties of
Ionic Compounds
Electrical Conductivity of Ionic Compounds
•
Ionic compounds conduct electricity in the molten
and aqueous states.
•
They do not conduct electricity in the solid state.
•
In the molten and aqueous states, mobile ions are
present.
•
Mobile ions conduct electricity.
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Chapter 6
Ionic Bonding
Concept Map
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