IN SUMMARY……

3 types

of bondi

ng

Ionic bondin

g

Covalent

bonding

Metallic

bonding

Pre

pare

d b

y JG

L

8/2

1/2

00

9

1

COVALENT BONDING

Pre

pare

d b

y JG

L

8/2

1/2

00

9

2

Two hydrogen atoms can share their valence electrons to attain the same electron configuration of the nearest Noble gas configuration, Helium

Covalent bond is the sharing of

two electrons between the

2 atoms

Covalent Bond

When nonmetallic elements react with other nonmetallic

elements, they share electrons in order to obtain eight valence

electrons.

Each fluorine atom has seven valence electrons. They each

require one more electron to satisfy the Octet Rule.

F F

F FThe left fluorine atom now has a total of eight electrons and the

right fluorine atom now has a total of eight electrons around it.

When nonmetallic elements react with other nonmetallic elements,

they share electrons in order to obtain eight valence electrons.

FF

The two electrons that form the covalent bond are often

Represented by a single line. The F2 molecule can be

represented using a line and dots to show the bonding pair

and the six lone pairs, respectively. This is called a Lewis dot

structure.

FF

Multiple Covalent BondSome atoms have to share more than one electron in order

to satisfy the Octet Rule.

O O

O O

Each oxygen atom has six valence electrons. They each

require two more electrons to satisfy the Octet Rule.

• The left oxygen atom now has a total of eight electrons around it. The right oxygen atom now has a total of eight electrons around it.

O O

O O

The four electrons shared by the oxygen atoms form a

double bond.

The double bond is represented by two single lines. Each line

in the Lewis dot structure represents two electrons

The element hydrogen is an exception to the Octet Rule. It

only needs two electrons, rather than eight, to be stable.

H FThe hydrogen atom has one valence electron. It requires one

more electron to be stable. The fluorine atom has seven

valence electrons. It requires one more to satisfy the Octet

rule.

H F

H FThe hydrogen atom now has a total of two electrons around

it and is stable.

The fluorine atom now has a total of eight electrons around

it and is stable.

• The Lewis dot structure of the HF molecule shows a line and 6 dots to represent the bonding pair and the 3 lone pairs of electrons, respectively.

H F H F

Rules for writing Lewis Dot structures

• Rule 1

Add together the number of valence electrons for each atom

in the molecule. For example, CF4

Carbon has four valence electrons and each fluorine has

seven valence electrons = 4 + 4(7)

= 32

Rule 2

Write out the elements of the molecule so that the least

electronegative elements is in the center surrounded by the

other elements. For example, CF4

C FF

FF

Rule 3

Place a covalent bond between the central atom and the

outside atoms. Remember each covalent bond contains two

electrons.

C F

F

F

F

The four covalent bonds use eight of the 32 valence electrons in CF4

• Rule 4There are 24 valence electrons remaining. Add electrons to

the outer atoms as lose pairs to satisfy the Octet Rule.

C F

F

F

F• This uses 24 electrons.

There Are no electrons left, so this is The Lewis dot structure for CF4

Rule 5 for example, NH3

• First apply Rules 1-4 to the molecule

• Rule 1: Count the valence electrons

• Rule 2: Place the least electronegative element at the

centre, except for H which is always an outer atom

• Rule 3: Add covalent bonds between the centre atom and

the outer atoms

• Rule 4: Add lone pairs to the outer atoms

• Rule 5: Add lone pairs to the centre atom

Rule 1

Nitrogen has 5 valence electrons and each hydrogen has 1

valence electron

The total number of valence electrons = 5 + 3 (1) = 8

Rule 2

Hydrogen is always an outer atom and is never at the centre

of a molecule

N HH

H

Rule 3

Add the bonding electrons. This uses 6 of the 8 valence

electrons.

Rule 4

The 2 remaining valence electrons are not added to the outer

atoms, because each H has its maximum of 2 valence

electrons.

• Rule 3

Add the bonding electrons. This uses 6 of the 8

valence electrons.

• Rule 4

The 2 remaining valence electrons are not added to

the outer atoms, because each H has its maximum

of 2 valence electrons.

Rule 5

Place the remaining 2 Valence

electrons on the central

nitrogen atom

Rule 6

Check all atoms in the molecule to ensure that each has 8electrons(2 for hydrogen). If an atom has fewer than 8electrons, create double or triple bonds. (Note: Doublebonds only exist between C,N,O and S atoms)

N HH

HThis is the Lewis structureFor NH3

Apply rule 6 to the following; CH4, CF4,

C H

H

H

H

• Hydrogen : 1 bond = 2 electrons (stable)

• Carbon : 4 bonds = 8 electrons (stable)

C F

F

F

F• Fluorine : 1 bond + 3

lone pairs = 2 + 3 (2)

= 8 electrons (stable)• Carbon : 4 bonds = 8

electrons (stable)

Example; CH2O

Apply Rules 1-5 to the molecule

Rule 1: Count the valency electrons

Rule 2: Place the least electronegative element at the

centre, except for H, which is always an outer

atom

Rule 3: Add covalent bonds between the centre and the

outer atoms

Rule 4: Add lone pairs to the outer atoms

Rule 5: Add lone pairs to the centre atom

Rule 1

Carbon has 4 valence electrons, each hydrogen has 1 valence

electron, and oxygen has 6 valence electrons.

Total number of valence electrons : 4 + 2(1) + 6 = 12

Rule 2

Carbon is at the centre of the molecule because it is less

electronegative than oxygen. Hydrogen is always an outer

atom and is never at the centre of the molecule.

CH

O

H

• Rule 3

Add the bonding electrons.

This uses 6 of the 12 valence

electronsC

H

O

H

CH

O

H • Rule 4

Add the remaining 6 lectrons to

the outer atom. Hydrogen does

not need any more electrons, but

Oxygen needs 6 to complete its

octet.

CH

O

HC

H

O

H

• Rule 6Oxygen shares one of its

lone pairs with C and O and give the desired 8 electron total

Rule 5 There are no valence electrons left to add to the centre

CH

O

HThis is the Lewis dotStructure for CH2O

Exceptions to the Octet Rule

The Octet Rule applies to Groups IVA through VIIA in the

second row of the Periodic Table, but there are exceptions

to the rule among some other elements. The following two

cases are an example

Example BF3

Rule 1

Boron has 3 valence electrons and each Fluorine has 7

valence electrons

Total number of electrons = 3 + 3 (7) = 24

BF

F

F Rule 2

Boron is at the centre of

the molecule because it is

less electronegative than

fluorine

Rule 3

Add the bonding electrons.

This uses 6 of the 24 valence

electrons

BF

F

F

Rule 4

Add the remaining electrons

to the outer atoms. Each

Fluorine has the required 8

electrons

BF

F

F

Rule 5

This uses the remaining

electrons leaving none to add

to the Boron central atom BF

F

F

Rule 6

Check the number of electrons around each atom. Each

Fluorine atom has 8 electrons, but the Boron Atom has only

6. This is an exception to the Octet Rule. A B=F bond is not

an option, because double bonds exist only between C,N,O,

and S atoms

BF

F

FThis is the Lewisdot structure BF3

METALLIC BONDING

Pre

pare

d b

y JG

L

8/2

1/2

00

9

31

However, metals behave differently.

Metallic bonding is similar to both covalent and ionic bonding

The valence (outermost) electrons are loosely held

by the metal ions, so much so that they move away from the atom to

form a positively charged ION.

The electrons are free to move from one positively charged ION

to the next (i.e. They are DELOCALISED) and are shared (just like in covalent bonding among the various metallic

positively charged ions

The number of electrons = the number of protons.

The metal is therefore electrically NEUTRAL

Source: www.daviddarling.info/images/metallic_bond.jpg

COMPARE AND CONTRAST TYPES OF BONDING P

rep

are

d b

y JG

L

8/2

1/2

00

9

32

Metallic and ionic bonding involve electrostatic attractions between positive and negatively charged particles.

Metallic bonding shares electrons among the ions in a similar manner to how electrons are shared in covalent bonding.

Covalent bonding shares electrons rather than having electrostatic charges.

Similarities Differences