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CHAPTER 6Ionic Bonding
© 2013 Marshall Cavendish International (Singapore) Private Limited
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
Chapter 6 Ionic Bonding
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Learning Outcome
6.1 The Stable Electronic Configuration of a Noble Gas
• describe the stable electronic configuration of a noble gas.
At the end of this section, you should be able to:
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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
All other noble gases have an octet configuration
(2 outer electrons).
(8 outer electrons).5
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
Ionic bonding Covalent bonding
Atoms share electrons to attain noble gas
configuration
Atoms gain or lose electrons to attain
noble gas configuration
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6.1 The Stable Electronic Configuration of a Noble Gas
Chapter 6 Ionic Bonding
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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
• describe the formation of positive ions (cations) and negative ions (anions) to achieve the noble gas configuration.
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Learning Outcome
At the end of this section, you should be able to:
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 protons10
6.2 Forming Ions
• 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
What is an Ion?
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
The Na atom loses one outer electron to form the Na+ ion. Why?
To achieve stable octet (noble gas) configuration.
Example 1: Formation of sodium (Na+) ion
Na atom
Electronic configuration: 2, 8, 1
Number of protons = 11
Number of electrons = 11
Neon (2, 8) 13
6.2 Forming Ions
Na atom: 11p, 12n, 11e
2, 8, 1 2, 8
sodium atom loses one outer electron
+
NeutralNa atom
Positively-chargedNa+ ion
Charge = 11p + 11e = (+11) + (–11) = 0
Na+ ion: 11p, 12n, 10e
Charge = 11p + 10e = (+11) + (–10) = +1
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6.2 Forming Ions
Example 1: Formation of sodium (Na+) ion
Charge = 20p + 20e
Ca2+ ion: 20p, 20n, 18e
2, 8, 8, 2 2, 8, 8
Ca atom: 20p, 20n, 20e
2+calcium atom loses two outer electrons
NeutralCa atom
Positively-chargedCa2+ ion
= 20(+1) + 20(–1) = (+20) + (–20) = 0
Charge = 20p + 18e
= 20(+1) + 18(–1)
= (+20) + (–18) = +2
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6.2 Forming Ions
Example 2: Formation of calcium (Ca2+) ion
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+
Common Cations and Their Charges
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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
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) 18
6.2 Forming Ions
Example 1: Formation of chloride (Cl–) ion
chlorine atom gains one electron
2, 8, 7
Cl atom: 17p, 18n, 17e Cl– ion: 17p, 18n, 18e
NeutralCl atom
Negatively chargedCl– ion
Charge = 17p + 17e = (+17) + (–17) = 0
Charge = 17p + 18e = (+17) + (–18) = –1
2, 8, 8
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6.2 Forming Ions
Example 1: Formation of chloride (Cl–) ion
oxygen atom gains two electrons
O atom: 8p, 8n, 8e O2– ion: 8p, 8n, 10e
NeutralO atom
Negatively chargedO2– ion
Charge = 8p + 8e = (+8) + (–8) = 0
Charge = 8p + 10e = (+8) + (–10) = –2
2, 82–
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2, 6
6.2 Forming Ions
Example 2: Formation of oxide (O2–) ion
Non-metal Ion Formula of ion
chlorine chloride ion Cl–
bromine bromide ion Br–
oxygen oxide ion O2–
sulfur sulfide ion S2–
Common Anions and Their Charges
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6.2 Forming Ions
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
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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
Learning Outcome
• describe how an ionic bonds are formed between metals and non-metals.
At the end of this section, you should be able to:
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6.3 Ionic Bond: Transferring Electrons
Examples:Group VII: Fluorine, chlorineGroup VI: Oxygen, sulfur
Examples:Group I: Sodium, potassiumGroup II: Magnesium, calcium
Ionic Bonding
• Ionic bonds are formed between metals and non-metals.
• This is done through the transfer of electron(s) from metals to non-metals.
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6.3 Ionic Bond: Transferring Electrons
Metallic atom Non-metallic atom
loses electron(s) gains electron(s)
Positive ion(cation)
Negative ion(anion)
electrostatic forces of attraction
Ionic Bonding
(hold oppositely charged ions together) 26
6.3 Ionic Bond: Transferring Electrons
Step 1: Formation of Positive Ions
Each sodium atom (Na) loses its single outer electron to form a positively-charged sodium ion (Na+).
Na Na+ + e−
2, 8, 1 2, 8
Formation of Ionic Compound
Example 1: Sodium chloride
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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 –Cl
2, 8, 7 2, 8, 8
+ e−
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6.3 Ionic Bond: Transferring Electrons
Sodium and chlorine react in the ratio of 1 : 1 to form sodium chloride (NaCl).
Sodium atom2, 8, 1
Chlorine atom2, 8, 7
Sodium ion2, 8
Chloride ion2, 8, 8
Electrostatic forces of attraction
Step 3: Formation of Ionic Bonds
Gains one electron
Loses one electron
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6.3 Ionic Bond: Transferring Electrons
Chlorine atoms gain one electron each.
Chloride ion2, 8, 8
Chloride ion2, 8, 8
Magnesium ion2, 8
Magnesium atom loses two
electrons.
Magnesium reacts with chlorine in the ratio of 1 : 2 to form magnesium chloride (MgCl2).
Example 2: Magnesium chloride
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6.3 Ionic Bond: Transferring Electrons
Chapter 6 Ionic Bonding
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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
Learning Outcome
• deduce the chemical formula of an ionic compound from the charges on the ions and vice versa.
At the end of this section, you should be able to:
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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
Since 1 × (+2 charge) balances out 1 × (−2 charge),
Example: Magnesium oxide
Magnesium forms Mg2+ ions.
Oxygen forms O2− ions.
Mg2+ O2−
The formula is MgO.
Charge: +2 Charge: −2
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6.4 Chemical Formulae of Ionic Compounds
Since 1 × (+2 charge) balances out 2 × (−1 charge),
Copper ion Hydroxide ion
Cu2+ OH−
The formula is Cu(OH)2.
Charge: +2 Charge: −1
To balance the charges, multiply the smaller charge (−1) by 2 to make it equal to +2.
Example: Copper(II) hydroxide
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6.4 Chemical Formulae of Ionic Compounds
Example 1
Write the chemical formula of aluminium oxide.
aluminium ion oxide ion
Al 3 + O 2 −
Al2O3
Charge: +3 Charge: −2
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Therefore, the formula is Al2O3.
6.4 Chemical Formulae of Ionic Compounds
Since ‘2’ is a common factor, it can be removed. Therefore, the formula is CaCO3.
CaCO3
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Example 2Write the chemical formula of calcium carbonate.
calcium ion carbonate ion
Ca 2 + CO3 2 −
Ca2(CO3)2
Charge: +3 Charge: −2
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6.4 Chemical Formulae of Ionic Compounds
Chapter 6 Ionic Bonding
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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
Learning Outcomes
• state that ionic compounds form giant lattice structures;
At the end of this section, you should be able to:
• deduce the formulae of ionic compounds from their lattice structures;
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• relate the physical properties of ionic compounds to their lattice structures.
6.5 Structure and Physical Properties of Ionic Compounds
Ionic compounds form giant ionic structures.
Structure of Ionic Compounds
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
Sodium ions and chloride ions alternate with each other.
Structure of NaCl
Three-dimensional arrangement of sodium ions and chloride ions
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Sodium chloride crystal
6.5 Structure and Physical Properties of Ionic Compounds
Strong forces of attraction between ions in crystal lattice
A large amount of energy is required to overcome these forces of attraction between ions.
Structure of NaCl
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–
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6.5 Structure and Physical Properties of Ionic Compounds
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Each sodium ion is surrounded by six chloride ions.
Each chloride ion is surrounded by six sodium ions.
Cl− ion
Na+ ion
The ratio of sodium ions to chloride ions is 1 : 1.Hence, the formula unit of sodium chloride is NaCl.
Structure of NaCl
6.5 Structure and Physical Properties of Ionic Compounds
• High melting and boiling points
• Non-volatile
• Exist as solids at room temperature
Melting and Boiling Points of Ionic Compounds
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Na+
Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–Cl–
Cl–Cl–
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6.5 Structure and Physical Properties of Ionic Compounds
• Usually soluble in water
Solubility of Ionic Compounds
• Usually insoluble in organic solvents E.g. ethanol, turpentine, petrol
Water molecules
dissolve in water
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Cl–Na+ Na+
Cl–Na+ Na+
Cl–Na+Cl–
Na+
Cl–
Na+
Cl–
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6.5 Structure and Physical Properties of Ionic Compounds
aqueous NaClsolid NaCl
molten NaCl
Electrical Conductivity of Ionic Compounds
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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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6.5 Structure and Physical Properties of Ionic Compounds
Concept Map
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Chapter 6 Ionic Bonding
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