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CHAPTER 2
Chemical Bonding

Chapter Scopes
Bond energies, bond lengths & bond polarities Drawing Lewis structure and calculate the formal charge

Forms of Chemical Bonds
1) Intramolecular bond – forces hold the atoms _______ a molecule
• Ionic / Electrovalent Bond
• Covalent Bond
• Metallic Bond
2) Intermolecular bond – forces ________ the molecules • Hydrogen bonding
• Van der Waals
3) Co-ordinate / Dative Bond

FHSC1114 Physical Chemistry

Chapter Scopes
•
•
•
•

Electrovalent / ionic bonding
Covalent bonding
Co-ordinate / dative covalent bonding
Intermolecular bonding (including hydrogen bonding, Van der Waals)
• Metallic bonding
• Electronegativity

Chemical Bond
• Chemical bond − the force of attraction that binds atoms together in a chemical compound • When atoms react to form chemical bonds, only the electrons in the ______
___________ are involved

Electrons In An Atom
2 groups of electrons:
a) Valence electrons
• outermost shell electrons
• chemical properties
• = group number in periodic table
b) Core electrons
• inner shell electrons
• not involved in chemical behaviour

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Valence & Core Electrons
B (1s22s22p1)
Core e- = [He]

Bond & Lone Pairs Electrons
• Valence electrons are distributed as shared (bond pairs) & unshared (lone pairs) Valence e- = 2s22p1

Core e- = [Ar]3d10
Valence e- = 4s24p5

Lewis Symbol &
Lewis Structures

Cl

••

Br ([Ar]3d104s24p5)

lone pair e-

••

H

••

Shared or bond pair eThis is called a LEWIS ELECTRON DOT structure (Lewis Structure)

Building A Lewis Electron
Dot Structure

• Lewis symbol — a chemical symbol to represent the nucleus & core electrons of an atom, together with dots placed around the symbol to represent the valence electron • For example: Si
•
Electron configuration:
• Si •
[Ne]3s2 3p2
•

Example 1: Ammonia, NH3
1) Decide on the central atom;
Never H
Central atom is atom of lowest affinity for electrons.
Therefore, N is central.
2) Count valence electrons
H = 1, N = 5
Total = (3 x 1) + 5 = 8 electrons / 4 pairs

3) Form a single bond between the central atom and each surrounding atom

Example 2: Carbon Dioxide, CO2
1) Central atom = _______
2) Valence electrons = __ e- or __ pairs e3) Form bonds
O C O 6 pairs of electrons are now left

4) Remaining electrons form LONE PAIRS to complete octet (8 e-) as needed
••

4) Place lone pairs on outer atoms

3 BOND PAIRS and 1 LONE PAIR

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Example 3: Sulfur Dioxide, SO2
1) Central atom = _____
2) Valence electrons = __ e- or ___ pairs e3) Form bonds
Leave 14 e4) Remaining pairs become lone pairs, first on outside atoms and then on central atom ••
••
••

5) So that C has an octet, we shall form
DOUBLE BONDS between C & O

Each atom is surrounded by an octet of electrons

5) Form double bond so that has an octet — but note that there are 2 ways of doing this bring in left pair
••

O
••

••

S

••

O
••

••

••

or bring in right pair

S

O
••

••

••

O

••

Resonance Structures
• 2 or more Lewis structures can be drawn for a molecule or polyatomic ion
• For example:
a) Sulfur dioxide, SO2

This leads to the following structures
••

O
••

••

S

••

••

O
••

••

••

••

S

••

••

O

O
••

b) Nitrogen dioxide, NO2

These equivalent structures are called
RESONANCE STRUCTURES

Octet Rule

Exceptions To Octet Rule

• Octet rule — the tendency of molecules & polyatomic ions to gain, lose or share electrons until they are surrounded by __ valence electrons
• Each atom has a share in 4 pairs of electrons, so each has achieved a stable noble gas configuration

Exceptions when molecules & ions have:
a) fewer than 4 pairs of electrons (8 valence electrons) on a central atom (BF3)
b) more than 4 pairs of electrons (8 valence electrons) on a central atom (PCl5, SF6)
c) have an odd number of electrons (NO,
NO2)

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Formal Atom Charges

Example:
Carbon Dioxide, CO2

• The formal charge for an atom in a molecule / ion is the charge calculated for that atom based on the Lewis structure of the molecule / ion
• Formal charge
= Group number – 1/2 (no. of bonding electrons) – (no. of lone pair electrons)
= No. valence electrons in free atom –
1/2 (no. of bonding electrons) – (no. of lone pair electrons)

Formal Charge (O) = 6 – 1/2(4) – 4 = 0
••

• Transferring of electrons from 1 atom to another, creating _______ & _______ ions
• Electrostatic attraction between 2 oppositely charged ions
• For example:
••

••

••
Ionic
compound

• The mutual attraction of 2 nuclei for the same electrons
• Sharing of valence electrons between atoms A +

H

B

H

A

H

• For the formation of NaCl:
a) Na (1s22s22p63s1)
Na+ (1s22s22p6)
≅ [Ne] noble gas configuration
b) Cl ([Ne]3s23p5)
Cl- ([Ne]3s23p6)
≅ [Ar] noble gas configuration

1) When 2 atoms approach each other, the electrons repel one another due to like charges Electron-electron repulsion

B

• A region of high electron density resulting from the overlapping atomic orbitals between bonded atoms

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Valence Electron Configurations
& Ionic Compound Formation

Formation of Covalent Bond
Example: H• •H

Covalent Bonding

H

–

••

••

••
••
nonmetal e- transfer from atom reducing agent to oxidizing agent

Na+ Cl
••

Na • • Cl

••

metal atom ••

+ • Cl

••

O

Formal Charge (C) = 4 – 1/2(8) – 0 = 0
∴ Sum of formal charge = 0 + 0 = 0

Ionic Bonding

Na •

C

••

••

O

2) Similarly, their nuclei also repel each other Nucleus-nucleus repulsion

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4) The nuclei & the electrons of neighboring atoms, however, attract each other
Electron-nucleus attraction
5) When the attraction forces > powerful than repulsive forces, a bond is formed

Attraction is greater than repulsion

Polar Covalent Bond in Polar
Molecule
H

F

Non-uniform distribution of electron density between the two atoms
Due to F is more electronegative than H, the F atom tends to exert a stronger attraction on the bonding e– compared to the H atom

Coordinate Covalent Bonding

Electron density of the bonding electrons to be higher around the F atom than the H atom ‘F end’ of molecule acquires a partial negative charge while the ‘H end’ acquires a partial positive charge
The covalent bond in the H–F molecule is polarised • Dative bond
• In some molecules or ions, a single atom contributes both of the electrons (lone pair e-) to be a shared pair of electrons
• Electron-pair ______ ― the atom which donates the lone pair electrons
• Electron-pair ______ ― the atom accepting the pair of electrons

Metallic Bonding

Dative bond
A dative bond is represented by an arrow pointing from the donor of the electron pair
(N atom) to the acceptor of the electron pair
(H atom)

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• Metallic bonding
– bonding in solid metals
– Electrostatic attraction between the positively charged metal ions & the ‘sea’
/‘cloud’ of delocalised (mobile) electrons
• For example: Cu, Fe, Al….
• Electron Sea Model – describe the metallic bonding

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Metallic Bonding
Electron Sea Model
+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

• e-

+

Positive metal ion

• In metals, the valence electrons are associated with a particular metal atom but are free to move (mobile) throughout the solid piece of metal
• Electrons are free to move away from a – ve electrode to a +ve electrode
(delocalised) when an electrical potential is applied
• With valence electrons now delocalised, the metal atoms are effectively ionised

Delocalised valence electrons

Effect of Bonding on Physical Properties
1) Metallic Bonding

high electrical conductivity high thermal conductivity ductility and malleability grey, black, brown/yellow in appearance not soluble melting point (depend on bond strength) solid at room temperature & non volatile

3) Covalent bond
• no electrical and thermal conductivity
• no ductility and malleability
• water – colourless
• halogen – colour
• nonpolar molecules – insoluble in water but soluble in organic solvents
• polar molecule – soluble in water
• giant structure – insoluble in all solvents

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2) Ionic bonding
• poor conductor in solid state (electrical conductivity only in aqueous solution)
• no thermal conductivity
• no ductility and malleability
• most are colourless
• soluble in polar solvents but insoluble in nonpolar solvents
• high melting points
• solid at room temperature

Intermolecular
Forces of Attraction
• In addition to the covalent bonds within a molecule, there are also other forces of attraction between molecules
• Intermolecular forces of attraction ______ than covalent / ionic bonds
• 2 types of intermolecular forces:
a) Van der Waals forces
b) Hydrogen bond

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Van der Waals Forces

3 Types of Dipoles

• Van der Waals forces are weak intermolecular forces that contributed by permanent dipole – permanent dipole & temporary dipole – induced dipole attraction • The larger the molecule size, or the larger the number of electrons in a molecule, the larger the Van der Waals forces, & the higher the melting point/boiling point of the molecule

1) Permanent dipole-permanent dipole attractions The force of attraction between the negative end of a ______ molecule & the positive end of another ______ molecule

2) Dipole-induced dipole attractions

3) Induced dipole-induced dipole attractions
(London Dispersion Forces)
• Interaction between ________ molecules
(O2, N2, CO2 & noble gases)
• This is caused by the random movement of the e- in an atom or molecule
• For example:
Argon has 18 e- in its atom. Since it is non-polar, the arrangement of e- is symmetrical on the average

• Interaction between a _____ molecule
& ________ molecule
+

neutral

-

The nonpolar is polarized by the polar molecule polarized
+

+

-

-

• When these e- revolve around the nucleus, the e- density might be higher at one end than the other & cause a temporary dipole
Electron cloud δ+ dipole – dipole attraction

• The positive end of the temporary dipole will distort the e- cloud of the neighboring atoms giving rise to induced dipoles temporary dipole then attract one another δ- δ+

δ-

δ+

δ-

δ+

δ-

temporary dipole

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Temporary dipole-induced dipole attractions

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Hydrogen Bonding in
Hydrogen Fluoride (HF)

Hydrogen Bond
• A special type of permanent dipole – permanent dipole attraction between a H atom, which is bonded to:

δ– δ+ (a) a small & highly electronegative atom
(O, N or F) of a neighbouring molecule
(b) a lone pair of electrons of another very electronegative atom

δ–
••

δ+

δ+

••

δ+

δ+ δ– δ+

δ–

δ+

If 2 molecules of HF are close to another, the H atom of 1 molecule will be attracted to the F atom of another molecule
Electrostatic attraction between the partial
+ve charge on the H atom & the partial – ve charge on F atom

Hydrogen Bonding in
Ammonia (NH3) δ+ δ+

δ–

Bond Properties
•
•
•
•

Bond Polarity
Bond Order
Bond Length
Bond Energy

• In NH3 molecule, the N atom has 1 lone pair of electrons
• Each NH3 molecule can form 1 hydrogen bond Bond Polarity
• Polar bond in polar molecules – the bond between the two atoms has a partial +ve end & a partial -ve end (dipole moment)
• The unequal _______ of electrons leads to:
a) a partial -ve charge on the more electronegative element (δ −)
b) a partial +ve charge on the less electronegative element (δ +)
• A covalent molecule is polar if the covalent bond is polarised

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• Electronegativity, χ
a) a measure of the ability of an atom in a molecule to attract bonding electrons in a covalent bond to itself
b) decide whether a bond is polar, which atom of the bond is -ve & which is +ve
c) increase from left to right across a period & decrease down a group
Electronegativity trend: F > O > N > C > H
F > Cl > Br > I

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Bond Order
• Bond order
= number of bonding electron pairs shared by 2 atoms in a molecule

Bond Order

Number of shared pairs linking X – Y
Number of X-Y links

Type of bonding Bond order

C−C
−

Single

1

C=C
=

• Bond order (BO) =

Bond

Double

2

C≡C
≡

Triple

3

Bond Length

Bond Order
• Bond order is depends on to 2 important bond properties:
a) Bond length
– the distance between the centers of 2 atoms joined by a bond
b) Bond energy
– the energy required to break a bond

Bond Length
• Bond length depends on bond order

>

>

• The bond lengths become ______ as the bond orders increase
• Bond length:
Single bond > double bond > triple bond

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• Bond length depends on size of bonded atoms • Size of bonded atoms ↑, bond length __

<
H—F

<
H—Cl

H—I

Bond Energy
Type of bond

Typical bond energy
(kJ mol-1)
Covalent bonds
200 – 400
Hydrogen Bonds
20 – 80
Van der Waals bonds
Less than 20
The relative strengths of chemical bonds: covalent bonds > hydrogen bonds > Van der
Waals bonds

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Relation of Bond Length & Bond
Strength with Bond Order
• The GREATER the number of bonds
(bond order), the HIGHER the bond strength (energy required)

Summary
• Draw Lewis structure & calculate formal charges • Differentiate all types of bonding
• Relate bonding to their properties

• The GREATER the number of bonds
(bond order), the SHORTER the length of bond FHSC1114 Physical Chemistry

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