This document provides an overview of different types of bonds that can form between atoms, including ionic bonds, covalent bonds, metallic bonds, and intermolecular forces. Ionic bonds form when electrons are transferred between atoms. Covalent bonds form when atoms share electrons. Metallic bonds form through delocalized electrons within metal lattices. Intermolecular forces, such as van der Waals forces, hydrogen bonding, and dipole interactions provide weaker attractions between molecules. The type of bonding between atoms influences various molecular and bulk properties.

2. CONTENTS
• Introduction
• Ionic bonding
• Covalent bonding
• Co-ordinate (dative covalent) bonding
• Metallic bonding
• Shape of molecules molecules
• Van der Waals’ forces
• Electronegativity & dipole-dipole interaction
• Hydrogen bonding
•

3. WHYDOELEMENTS
FORMABOND
• the noble gases (He, Ne, Ar, Kr, Xe and Rn) are
in Group VIII
• they are all relatively, or totally, inert
• their electronic structure appears to confer
stability
• they have filled their ‘outer shell’ of electrons
• atoms without the electronic structure of a
noble gas try to get one
• various ways are available
• the method depends on an element’s position
in the periodic table

4. Ionic
bonding
 Ionic bonds tend to be formed between
elements whose atoms need to “lose”
electrons to gain the nearest noble gas
electronic configuration and those
which need to gain electrons. The
electrons are transferred from one atom
to the other.

5. Na ——> Na+ + e¯ and
Cl + e¯ ——> Cl¯

6.  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.

7. Mg ——> Mg2+ + 2e¯ and
2Cl + 2e¯ ——> 2 Cl¯

8. PROPERTIESOF
IONICBOND
 Melting point
 very highA large amount of energy must be put in to overcome the
 strong electrostatic attractions and separate the ions.
 Strength
 Very brittle Any dislocation leads to the layers moving and similar
 ions being adjacent. The repulsion splits the crystal.
 Electrical don’t conduct when solid - ions held strongly in the lattice
 conduct when molten or in aqueous solution - the ions
 become mobile and conduction takes place.
 Solubility Insoluble in non-polar solvents but soluble in water
 Water is a polar solvent and stabilises the separated ions.

9. IONIC COMPOUNDS - ELECTRICAL PROPERTIES
SOLID IONIC
COMPOUNDS DO
NOT CONDUCT
ELECTRICITY
Na+
Cl- Na+
Cl-
Na+
Cl-Na+
Cl-
Na+
Cl- Na+
Cl-
IONS ARE HELD STRONGLY TOGETHER
+ IONS CAN’T MOVE TO THE CATHODE
- IONS CAN’T MOVE TO THE ANODE
MOLTEN IONIC
COMPOUNDS DO
CONDUCT
ELECTRICITY
Cl-
IONS HAVE MORE FREEDOM IN A
LIQUID SO CAN MOVE TO THE
ELECTRODES
SOLUTIONS OF IONIC
COMPOUNDS IN
WATER DO CONDUCT
ELECTRICITY
DISSOLVING AN IONIC COMPOUND
IN WATER BREAKS UP THE
STRUCTURE SO IONS ARE FREE TO
MOVE TO THE ELECTRODES
Na+
Na+
Cl-
Na+
Cl-
Na+
Cl-

10. COVALENTBONDS
 consists of a shared pair of electrons with one
electron being supplied by each atom either side of
the bond.
 atoms are held together because their nuclei which
have an overall positive charge are attracted to the
shared electrons

11. The important
conditions of
covalent bond:
Each bond is formed as a result of sharing of an electron pair
between the atoms.
Each combining atom contributes at least one electron to the
shared pair.
The combining atoms attain the outer-shell noble gas
configurations as a result of the sharing of electrons.
By sharing of an electron pair by two atoms covalent bond is
formed

12. SINGLE
COVALENT
 When two atoms share one electron pair they are said to be
joined by a single covalent bond
 Eg
H2, H2O etc

13. FORMATIONOF
H2

14. FORMATIONOF
WATER
MOLECULE

15. DOUBLE
COVALENT
BOND
 If two atoms share two pairs of electrons, the covalent bond
between them is called a double bond.
 Eg. O2 and CO2

16. Formationof
oxygen
molecule

17. Formationof
carbondioxide

18. TRIPLE
COVALENT
BOND
 When combining atoms share three electron pairs
Eg
as in the case of two nitrogen atoms in the 𝑁2molecule and the
two carbon atoms in the ethyne molecule, a triple bond is
formed.

19. FormationofN2
molecule

20. Formationof
Ethyne
molecule

22. PROPERTIES
 Bonding Atoms are joined together within the molecule by covalent bonds.
 Electrical Don’t conduct electricity as they have no mobile ions or electrons
 Solubility Tend to be more soluble in organic solvents than in
water;
 some are hydrolysed
 Boiling point Low - intermolecular forces are weak;
 they increase as molecules get a larger surface area
 e.g. CH4 -161°C C2H6 - 88°C C3H8 -42°C
 as the intermolecular forces are weak, little energy is required to
 to separate molecules from each other so boiling points are low
 some boiling points are higher than expected for a given mass
 because you can get additional forces of attraction

23. POLARITYOF
COVALENT
COMPOUNDS
 Electrons are shared between different nonmetal atoms
 Example: HF
 In HF molecule the electrons shared between the two atoms
are attracted
 more towards fluorine; since the electronegativity of
fluorine is greater
 than that of hydrogen

25. DATIVE(
COORDINATE)
BOND
 A dative covalent bond differs from covalent bond only
in its formation
 Both electrons of the shared pair are provided by one
species (donor) and it shares the electrons with the
acceptor
 Donor species will have lone pairs in their outer shells
 Acceptor species will be short of their “octet” or
maximum.
 Lewis base a lone pair donor
 Lewis acid a lone pair acceptor

26. Boron
trifluoride-
ammonia
NH3BF3

27. Al2Cl6

29. Shapesof
molecules
Electron-pair
repulsion theory
The shape of a molecule depends upon the number of
valence shell electron pairs (bonded or unbonded) around
the central atom.
Pairs of electrons in the valence shell repel one another since
their electron are negatively charged.
These pairs of electrons tend to occupy such positions in
space that minimize repulsion and thus maximize distance
between them
The order of repulsion is lone pair–lone pair (most
repulsion) > lone pair–bond pair > bond pair–bond
pair (least repulsion).

30. SHAPEOF
METHANE
Methane has four bonding pairs of electrons
surrounding the central carbon atom. The equal
repulsive forces of each bonding pair of electrons
results in a tetrahedral structure with all H C H bond
angles being 109.5°

31. SHAPEOF
AMMONIA Ammonia has three bonding pairs of electrons and one
lone pair. As lone pair–bond pair repulsion is greater
than bond pair–bond pair repulsion, the bonding pairs
of electrons are pushed closer together. This gives the
ammonia molecule a triangular pyramidal shape. The
H N H bond angle is about 107°.

32. SHAPEOF
WATER
Water has two bonding pairs of electrons and two lone
pairs. The greatest electron pair repulsion is between
the two lone pairs. This results in the bonds being
pushed even closer together. The shape of the water
molecule is a nonlinear V shape. The H O H bond angle
is 104.5°.

33. Moremolecular
shapes
Borontrifluoride
 Boron trifluoride is an electron-deficient molecule. It
has only six electrons in its outer shell. The three
bonding pairs of electrons repel each other equally, so
the F B F bond angles are 120° We describe the shape
of the molecule as trigonal planar.

34. Phosphorus
pentafluoride
 Phosphorus pentafluoride has five bonding pairs of
electrons and no lone pairs. The repulsion between the
electron pairs results in the most stable structure
being a trigonal bipyramid . Three of the fluorine
atoms lie in the same plane as the phosphorus atom.
The bond angles FPF within this plane are 120°. Two
of the fluorine atoms lie above and below this plane at
90° to it

35. Sulfur
hexafluoride Sulfur hexafluoride has six bonding pairs of
electrons and no lone pairs. The equal repulsion
between the electron pairs results in the
structure shown above . All F S F bond angles
are 90°. We describe the shape as octahedral.

36. σ bondsandπ
bonds
 The covalent bond is classified into two types depending upon
the types of overlapping. They are
Sigma(σ) Bond
Pi(π) Bond

37. σbonds
 This bond is formed by the end to end (head-on) overlap of
bonding orbitals along the internuclear axis. This is called as
head on overlap or axial overlap.

38. S-S Overlapping: Overlap of two half filled s-orbitals
S-P Overlapping: Overlap occurs between half filled s-orbitals of one atom
and half filled p-orbitals of another atom
P–P Overlapping: Overlap takes place between half filled p-orbitals of the
two approaching atoms

39.  π bond is formed when the atomic orbitals overlap in such a
way that their axes remain parallel to each other and
perpendicular to the internuclear axis
Pi(π) Bond

40. Strength of Sigma
and pi Bonds
The strength of a bond depends upon the extent of
overlapping.
In case of sigma bond, the overlapping of orbitals takes place
to a larger extent. Hence, sigma bond is stronger as compared
to the pi bond where the extent of overlapping occurs to a
smaller extent.
Pi bond between two atoms is formed in addition to a sigma
bond.
It is always present in the molecules containing multiple
bond

41. Theshapeof
someorganic
molecules
 We can explain the shapes of molecules in terms of the
patterns of electron density found in σ bonds and π
bonds

42. Ethane

43. Ethene
Each carbon atom in ethene uses three of its four outer
electrons to form σ bonds. Two σ bonds are formed with
the hydrogen atoms and one σ bond is formed with the
other carbon atom. The fourth electron from each
carbon atom occupies a p orbital, which overlaps
sideways with a similar p orbital on the other carbon
atom. This forms a π bond.

44. 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 strong
electrostatic force of attraction between the cation
and electrons.

45.  increasing positive charge on the ions in the metal
lattice
 decreasing size of metal ions in the lattice
 increasing number of mobile electrons per atom.
METALLICBOND
STRENGTH

46. Na
The strength of the metallic bonding in
sodium is relatively weak because each
atom donates one electron to the cloud.
The metallic bonding in potassium is
weaker than in sodium because the
resulting cation is larger
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 cations together more strongly.
Mg
K

47. PROPERTIES
 Metals have high melting and boiling point ( strong
electrostatic force of attraction ) ,it increases down
the group.
 Metals are excellent conductors of heat and
electricity ( due to delocalised mobile electrons)
 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.

48. Intermolecular
forces
 The forces between molecules are much weaker. We
call these forces intermolecular forces.
 There are three types of intermolecular force:
 ■■ van der Waals’ forces (which are also called
dispersion forces and temporary dipole–induced dipole
forces)
 ■■ permanent dipole–dipole forces
 ■■ hydrogen bonding.

49. Strengthof
differentbonds

50. VAN DER
WAALS
FORCES
 Bromine is a non-polar molecule that is liquid at room
temperature. The weak forces of attraction are keeping
the bromine molecules together at room temperature.
These very weak forces of attraction are called van der
Waals’ forces. van der Waals’ forces exist between all
neutral atoms or molecules.

51. causes
• electrons in atoms or molecules are moving at high
speeds in orbitals, it is possible for more electrons to
be on one side of an atom/molecule. A dipole forms
when one side is slightly negative and the other slightly
positive, a dipole in one atom/molecule can then
induce a dipole in a neighboring one.
van der Waals’ forces increase with
■■ increasing number of electrons (and protons) in the
molecule
■■ increasing the number of contact points between the
molecules – contact points are places where the
molecules come close together.

53. Permanent
dipole-dipole
interactions
The forces between two molecules having permanent dipoles are called
permanent dipole–dipole forces. The attractive force between the δ+ charge on one
molecule and the δ– charge on a neighbouring molecule causes a weak attractive
force between the molecules.

54. Hydrogen
bonding
 Hydrogen bonding is the strongest type of
intermolecular force. For hydrogen bonding to occur
between two molecules we need: ■■ one molecule
having a hydrogen atom covalently bonded to F, O or N
(the three most electronegative atoms) ■■ a second
molecule having a F, O or N atom with an available
lone pair of electrons. eg

55. HYDROGEN BONDING - HF

56. Howdoes
hydrogen bonding
affectboiling
point?
Some compounds may have higher boiling points than
expected. This can be due to hydrogen bonding

57. each water molecule is hydrogen-bonded to 4
others in a tetrahedral formation
ice has a “diamond-like” structure
volume is larger than the liquid making it
when ice melts, the structure collapses
slightly and the molecules come closer; they
then move a little further apart as they get
more energy as they warm up
this is why…
a)water has a maximum density at 4°C
b)ice floats.
The peculiar properties of water

58. Surfacetension
andviscosity
 Water has a high surface tension and high viscosity.
Hydrogen bonding reduces the ability of water
molecules to slide over each other, so the viscosity of
water is high. The hydrogen bonds in water also exert
a significant downward force at the surface of the
liquid. This causes the surface tension of water to be
higher than for most liquids.

59. THANK YOU