Topic 3

TOPIC 3

CHEMICAL BONDING

TYPE OF CHEMICAL BONDS

CHEMICAL BONDS

Interatomic Bonds

ionic (electrovalent)

metallic

covalent & dative

Intermolecular Forces

van der Waals forces

permanent dipole- dipole interactions

hydrogen bonds

IONIC

BONDING

THE IONIC BOND

The IONIC BOND is the electrostatic force of

attraction between two oppositely charged

ions formed as a result of the complete

transfer of one or more electrons from one

atom to another.

Metal atoms have low ionisation energies and tend to lose electrons easily to form positive ions (cations).

Non-metallic elements have high electron affinities and tend to gain electrons to form negative ions (anions).

The cations and anions formed, being oppositely charged are attracted together by strong electrostatics forces.

These forces are called IONIC BONDS.

THE IONIC BONDING

Formation of sodium chloride,NaCl

Cation (Na+)

Electronic configuration of Na = 1s22s22p63s1

Na atom can attain a full outer shell of 8 electrons by losing an electron from the 3s orbital to form a Na+

ion.

Na Na+ + e-

2.8.1 2.8

THE IONIC BONDING

Anion (Cl-)

Electronic configuration of Cl = 1s22s22p63s23p5

Cl atom can attain a full outer shell of 8 electron by gaining an electron from the Na atom to form a Cl-

ion.

Cl + e- Cl-

2.8.7 2.8.8

THE IONIC BONDING

An electron is transferred from the 3s orbital of sodium

to the 3p orbital of chlorine both species end up with

the electronic configuration of the nearest noble gas

the resulting ions are held together in a crystal lattice

by electrostatic attraction

1s22s22p63s1

1s22s22p6 1s22s22p63s23p5

1s22s22p63s23p6

2.8.1 2.8.7 2.8

2.8.8

ELECTRON

TRANSFER

Mg > Mg2+ + 2e and 2Cl + 2e > 2 Cl

2.8.2 2.8 2.8.7 2.8.8

Mg

Cl

Cl

e

e

THE IONIC BOND

Formation of Magnesium chloride

IONIC BONDING

Animations

SODIUM CHLORIDE

Cl

SODIUM ATOM

2,8,1

Na

CHLORINE ATOM

2,8,7

SODIUM CHLORIDE

Cl

SODIUM ION

2,8

Na

CHLORIDE ION

2,8,8

both species now have full outer shells; ie they have the electronic configuration of a noble gas

+

SODIUM CHLORIDE

Cl

SODIUM ION

2,8

Na

CHLORIDE ION

2,8,8

Na Na+ + e

2,8,1 2,8 ELECTRON TRANSFERRED

Cl + e Cl

2,8,7 2,8,8

+

MAGNESIUM CHLORIDE

Cl

MAGNESIUM ATOM

2,8,2

Mg CHLORINE ATOMS

2,8,7

Cl

MAGNESIUM CHLORIDE

Cl

MAGNESIUM ION

2,8

Mg CHLORIDE IONS

2,8,8

Cl

2+

PROPERTIES OF IONIC COMPOUNDS Hard, Brittle (Ionic solid)

Reason : powerful electrostatic forces that hold the ions in place

throughout the crystal.

Brittle = when the crystal is tapped sharply along a particular plane, one layer of ions is displaced relative to the next and ions of similar charge then come together and repel each other forcing apart the two portions of the crystal.

High melting point & boiling point

Reason : freeing the ions from their lattice positions (melting) and

vaporising them (boiling) requires large amounts of energy.

Do not conduct electricity in the solid state

Reason : an ionic solid consist of immobilised ions which the ions

are held in fixed positions and are not free to move.

~ when it melts or dissolves in water, the ions are free to

move and conduct an electric current.

Solubility

Insoluble in non-polar solvents but soluble in water

METALLIC

BONDING

THE METALLIC BOND

The METALLIC BOND is the electrostatic

force of attraction that two neighbouring

metal ions have for the delocalised electrons

between them.

METALLIC BONDING

Involves a lattice of positive ions surrounded by

delocalised electrons

Metal atoms achieve stability by off-loading electrons to attain the electronic structure of the

nearest noble gas. These electrons join up to

form a mobile cloud which prevents the newly-

formed positive ions from flying apart due to

repulsion between similar charges.

METALLIC BONDING

Involves a lattice of positive ions surrounded by delocalised electrons

Metal atoms achieve stability by off-loading electrons to attain the electronic structure of the nearest noble gas. These electrons join up to form a mobile cloud

which prevents the newly-formed positive ions from flying apart due to repulsion

between similar charges.

Atoms arrange in regular

close packed 3-dimensional

crystal lattices.

The outer shell electrons of each

atom leave to join a mobile cloud or sea of electrons which can roam throughout the metal. The

electron cloud binds the newly-

formed positive ions together.

METALLIC BOND STRENGTH

Depends on the number of outer electrons donated to the cloud and the size of the metal atom/ion.

The strength of the metallic bonding in

sodium is relatively weak because each

atom donates just one electron to the cloud.

The metallic bonding in potassium is weaker

than in sodium because the resulting ion is

larger and the electron cloud has a bigger

volume to cover so is less effective at

holding the ions together.

The metallic bonding in magnesium is

stronger than in sodium because each atom

has donated two electrons to the cloud. The

greater the electron density holds the ions

together more strongly.

Na

Mg

K

METALLIC PROPERTIES

MOBILE ELECTRON CLOUD ALLOWS THE CONDUCTION OF ELECTRICITY

For a substance to conduct electricity it must have mobile ions or electrons.

Because the ELECTRON CLOUD IS MOBILE, electrons are free to move throughout

its structure. Electrons attracted to the positive end are replaced by those entering

from the negative end.

Metals are excellent conductors of electricity

MALLEABLE CAN BE HAMMERED INTO SHEETS

DUCTILE CAN BE DRAWN INTO RODS AND WIRES

As the metal is beaten into another shape the delocalised electron cloud

continues to bind the ions together.

Some metals, such as gold, can be hammered into sheets thin enough to

be translucent.

METALLIC PROPERTIES

Metals can have their shapes changed relatively easily

HIGH MELTING POINTS

Melting point is a measure of how easy it is to separate the individual particles. In

metals it is a measure of how strong the electron cloud holds the positive ions.

The ease of separation of ions depends on the...

ELECTRON DENSITY OF THE CLOUD

IONIC / ATOMIC SIZE

PERIODS Na (2,8,1) < Mg (2,8,2) < Al (2,8,3)

m.pt 98C 650C 659C

b.pt 890C 1110C 2470C

METALLIC PROPERTIES

Na+ Al3+ Mg2+

MELTING POINT INCREASES ACROSS THE PERIOD

THE ELECTRON CLOUD DENSITY INCREASES DUE TO THE

GREATER NUMBER OF ELECTRONS DONATED PER ATOM. AS A

RESULT THE IONS ARE HELD MORE STRONGLY.

HIGH MELTING POINTS

Melting point is a measure of how easy it is to separate the individual particles. In

metals it is a measure of how strong the electron cloud holds the positive ions.

The ease of separation of ions depends on the...

ELECTRON DENSITY OF THE CLOUD

IONIC / ATOMIC SIZE

GROUPS Li (2,1) < Na (2,8,1) < K (2,8,8,1)

m.pt 181C 98C 63C

b.pt 1313C 890C 774C

METALLIC PROPERTIES

MELTING POINT DECREASES DOWN A GROUP

IONIC RADIUS INCREASES DOWN THE GROUP. AS THE IONS GET

BIGGER THE ELECTRON CLOUD BECOMES LESS EFFECTIVE

HOLDING THEM TOGETHER SO THEY ARE EASIER TO SEPARATE.

Na+ K+ Li+

COVALENT

BOND

HYDROGEN

H H

Another hydrogen atom

also needs one electron to

complete its outer shell

Hydrogen atom needs

one electron to

complete its outer shell

atoms share a pair of electrons to

form a single covalent bond

A hydrogen MOLECULE is formed

H H

H H

WAYS TO REPRESENT THE MOLECULE

COVALENT BOND

A Covalent Bond is the electrostatic force of

attraction that two neighbouring nuclei have

for a localised pair of electrons shared

between them.

HYDROGEN CHLORIDE

Cl H

Hydrogen atom also

needs one electron

to complete its outer

shell

Chlorine atom

needs one electron

to complete its

outer shell

atoms share a pair of

electrons to form a

single covalent bond

H Cl H Cl

WAYS TO REPRESENT THE MOLECULE

METHANE

C

Each hydrogen

atom needs 1

electron to

complete its

outer shell

A carbon atom needs 4

electrons to complete

its outer shell

Carbon shares all 4 of

its electrons to form 4

single covalent bonds

H

H

H

H

H C H

H

H

H C H

H

H

WAYS TO REPRESENT

THE MOLECULE

AMMONIA

N Each hydrogen atom needs

one electron to

complete its

outer shell

Nitrogen atom needs 3 electrons

to complete its outer shell

Nitrogen can only share 3 of its

5 electrons otherwise it will

exceed the maximum of 8

A LONE PAIR REMAINS

H

H

H

H N H

H

H N H

H

WAYS TO REPRESENT

THE MOLECULE

WATER

O Each hydrogen

atom needs

one electron to

complete its

outer shell

Oxygen atom needs 2 electrons


Oxygen can only share 2 of its 6

electrons otherwise it will


2 LONE PAIRS REMAIN

H

H

H O

H

H O

H

WAYS TO REPRESENT

THE MOLECULE

HYDROGEN

H

H H

H H H

H H

both atoms need one electron

to complete their outer shell

atoms share a pair of electrons

to form a single covalent bond

DOT AND

CROSS

DIAGRAM

METHANE

C

H H

H H

C

H

H

H

H

H C H

H

H

H C H

H

H

each atom needs one

electron to complete

its outer shell

atom needs four


its outer shell

Carbon shares all 4 of its

electrons to form 4 single

covalent bonds

DOT AND

CROSS

DIAGRAM

AMMONIA

N H H

H

N

H

H H

H N H

H

H N H

H

each atom needs one


its outer shell

atom needs three


its outer shell

Nitrogen can only share 3 of

its 5 electrons otherwise it will


A LONE PAIR REMAINS

WATER

O H

H O

H

H

each atom needs one


its outer shell

atom needs two


its outer shell

Oxygen can only share 2 of its

6 electrons otherwise it will


TWO LONE PAIRS REMAIN

H O

H

H O

H

OXYGEN

O

each atom needs two electrons


each oxygen shares 2 of its

electrons to form a

DOUBLE COVALENT BOND

O O O

O O

Dative bond (or Co-ordinate Bond)

a covalent bond (a shared pair of electrons) in which both electrons come from the same

atom.

COVALENT BOND

Sigma bond

Formed when the orbitals of two atoms have head to head overlapping.

COVALENT BOND

Pi bond

Formed when the p orbitals of two atoms have side to side overlapping.

COVALENT BOND

Shape of Simple Molecules

Valence shell electron pair repulsion (VSEPR)

1.Electrons pairs (bond pairs and lone pairs) repel each

other and move as far apart as possible.

2. Lone pairs of electrons repel more than boding pairs.

3. The repulsion between electron pairs is increased by

the increase in electronegativity of the central atom.

The ability of an atom to attract the pair of electrons in a covalent bond to itself.

Pauling Scale - a scale for measuring electronegativity

- values increase across periods

- values decrease down groups

- fluorine has the highest value

H

2.1

Li Be B C N O F

1.0 1.5 2.0 2.5 3.0 3.5 4.0

Na Mg Al Si P S Cl

0.9 1.2 1.5 1.8 2.1 2.5 3.0

K Br

0.8 2.8

INCREASE INC

RE

AS

E

Electronegativity

Polar Covalent Bond Non-polar bond - similar atoms have the same

electronegativity they will both pull on the electrons to the same extent the electrons will be equally shared

Polar bond - different atoms have different electronegativities

One will pull the electron pair closer to its end it will be slightly

more negative than average, - The other atom will be slightly less negative, or more positive, +

a dipole is formed and

The bond is said to be polarised (has a dipole) Greater the electronegativity difference, the greater the polarity

Example: hydrogen fluoride, HF, F(4.0), H(2.1)

Occurrence - not all molecules containing polar bonds are polar overall

- if bond dipoles cancel each other out the molecule wont be polar

- if there is a net dipole the molecule will be polar

HYDROGEN CHLORIDE TETRACHLOROMETHANE WATER

POLAR MOLECULES

NET DIPOLE - POLAR NON-POLAR NET DIPOLE - POLAR

A molecule is polar (has a dipole moment) if

its bond is polarised & it is not symmetrical

Determine whether the following molecules are polar or non-polar:

NH3 AlCl3 CH3Cl

CO2

Intermolecular

Forces

Introduction

Other than the covalent bonds in the molecule, there are also forces holding the

molecules together

Intermolecular Forces between

Covalent Bond

There are 3 general intermolecular forces:

Van der Waals forces

Dipole-Dipole Attraction

Hydrogen Bonds

Van der Waals Forces

forces of attraction between non-polar molecules, which arise due to induced

(temporary) dipole attraction.

Van der Waals Forces

Also known as Temporary Dipole-Dipole Forces, Dipole-Induced Dipole Attraction, London Forces or

Dispersion forces

Momentary dipoles occurring due to uneven electron distributions in neighbouring molecules as they

approach one another

This gives rise to a temporary dipole (induces dipole in neighbouring molecules)

Van der Waals Forces are very weak forces of attraction

The strength of van der Waals depends on:

- Number of electrons in the molecule

(No. of e- , stronger forces)

- Shape of the molecule

( Branching reduces the strength of the forces)

e.g penthane vs 2,2-dimethylpropane

The strength of the van der Waals forces increases with the total number of electrons in the molecule.

Boiling points of hydrides

Mr Boiling point

C

CH4 16 -161

SiH4 32 -117

GeH4 77 -90

SnH4 123 -50

Boiling point

increase when

the number of

electrons

increases

Dipole-Dipole Attraction

Due to the difference in the electronegativity

occur between molecules that have permanent dipoles

- polar molecules

partial positive charge on one molecule is electrostatically attracted to the partial negative charge

on a neighbouring molecule

stronger intermolecular forces than van der Waals

Example: HCl

Hydrogen Bonding

When a hydrogen atom is covalently attached to a very electronegative atom (N, O and F), the hydrogen atom can

form a hydrogen bond with another very electronegative atom, which has a lone pair of electron

The strongest intermolecular force

The molecules which have this extra bonding are:

Hydrogen Bonding in Ammonia

Hydrogen Bonding in Water

http://www.northland.cc.mn.us/biology/Biology1111/animations/hydrogenbonds.html

Hydrogen bonding

Covalent bonding

. ______

HYDROGEN BONDING in ICE

ice has a diamond-like structure

hydrogen bonding

Hydrogen Bonding in Hydrogen

Fluoride

F H

F H

H F

H F

d +

d

d +

d

d + d

d + d

hydrogen bonding

BOILING POINTS OF HYDRIDES

GROUP IV

GROUP V

GROUP VI

GROUP VII

Mr

BO

ILIN

G P

OIN

T / C

100

0

-160

140 50 100

H2O

HF

NH3

The higher than expected boiling

points of NH3, H2O and HF are due to

intermolecular HYDROGEN BONDING

Properties of Hydrogen Bonding Solubility in water

- form H bond with water soluble

- more H bond more soluble

Unusually high boiling point

- more energy is needed to overcome the strong H bond

Anomalous relative molecular masses

- Mr found by measurement in the vapour phase or in

organic solvents, Mr are twice large than expected.

Reason : Dimerisation take place

Intermolecular & intramolecular H bond

e.g 2-nitrophenol (214C) vs 4-nitrophenol (279C)

Topic 3

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