Chemical Bonding
The Basic Principles
• The Law of Definite Proportions (Joseph Louis
Proust, 1799) and Dalton’s development of
atomic theory (1803) lead to the recognition that
atoms of an element had a characteristic
combining ability with other atoms, which came
to be called valence.
• Existence of atoms suggested that compounds
were composed of collections of atoms bound
together by chemical bonds.
Introduction
• Atoms combine to form molecules.
• The combining power of atoms to form
molecules is called valency.
• All atoms having unstable or incomplete outer
shell have a tendency to gain or lose electrons
 tendency of atoms to complete and hence
stabilize their outermost orbit of electrons which
is mainly responsible for chemical combination
between the atoms.
• According to ‘electronic theory of valency’ ‘a
chemical bond is formed as a result of
electronic interactions.
• However, a molecule is formed only when
electrons of the constituent atoms interact in
such a way that the potential energy is
lowered.
• A chemical bond as an effect that causes the
energy of two atoms close together to be
markedly lower (by about 100 kJ per mole or
more) than when they are far apart.
• The forces that hold bonded atoms together are
basically just the same kinds electrostatic
attractions that bind the electrons of an atom to
its positively-charged nucleus.
Lewis electron – dot formulas
• Lewis originated the idea of the electron pair
bond.
• In 1902, Lewis developed the concept of
valence electrons and realized that all elements
known to form simple ions by losing or gaining
whatever number of electrons is needed to leave
eight in the valence shell of each.
• In 1916, Lewis published shared electron-pair
theory.
• Lewis Symbol/Structure --- representing atom
singly or in combination
• Only the valence electrons are shown.
• A Lewis dot symbol consists of the symbol of an
element and one dot for each valence electron
in an atom of the element.
• Note: except for He, the number of valence
electrons in each atom is the same as the group
number of the element.
Type of Bonds
• Electrovalent or ionic bond
Electropositive elements + Electronegative elements
• Covalent bond
Electronegative elements + Electronegative elements
• Coordinate bond
Electropositive elements + Electropositive elements
Ionic or Electrovalent Bonds
• The ionic bond is formed due
to the “electrostatic attraction
between stable ions formed by
the complete transfer of
electrons from one atom to
another”.
• The atom which loses
electrons acquires a positive
charge, whereas the one
which gains electrons
becomes negatively charged.
• Ionic bonds: anions + cations
Ionic Bonds
• These charged atoms are called ions and held
together by electrostatic attraction forces.
• Such a mode of combination of atoms is called
electrovalency and the bond formed between
the atom is called electrovalent or Ionic bond.
• The compound thus formed is called an
electrovalent or ionic compound.
Ionic Bonds
• Atoms having a tendency to lose electrons are
called electropositive whereas the atoms which
gain electrons are called electronegative.
• The number of electrons gained or lost by an
atom in order to acquire an inert gas
configuration gives numerical value of the
electrovalency of the atom.
• Example 1: NaCL
Ionic Bonds
• The electronic arrangement of Na and Cl
atoms are:
Na (11) 2, 8, 1 (1s2, 2s2, 2p6, 3s1)
Cl (17) 2, 8, 7, (1s2, 2s2, 2p6, 3s2, 3p5)
• The electron from the outermost orbit of Na is
completely transferred to the outermost orbit of
Cl atom. As a result of this transfer, both atoms
acquire inert gas structure.
Ionic Bonds
• The Na becomes Na+ (2, 8) and Cl become Cl–
(2, 8, 8).
• Example 2:
• The formation of lithium flouride (LiF)
Li
1s2 2s1
+
F
1s2 2s2 2p5
Li
F
+ e-
Li +
1s2
F
1s2 2s2 2p6
Li + + e F
(or LiF)
Ionic Bonds
• Electrovalent compounds exhibit following
properties:
– Electrovalent compounds are generally hard
solids.
– They have high melting and boiling point.
– Electrovalent compounds are generally
sparingly soluble in organic solvents.
Ionic Bonds
– Electrovalent compounds in solid state are
poor conductors of electricity. But when
dissolved in solvents of relatively high
dielectric constant, they exhibit a strong
electrical conductivity. They also conduct
electricity in the molten state.
Lattice energy of ionic compounds
• The stability of solid ionic compounds depends
on the interactions of all cations and anions, and
not merely on the interaction of a single cation
with a single anion.
• A quantitative measure of the stability of any
ionic solid is its lattice energy --- the energy
required to completely separate one mole of a
solid ionic compound into gaseous state.
Lattice energy of ionic compounds
• Lattice energy cannot be measured directly. It can be
determined by using Coulomb’s Law if structure and
composition of an ionic compound known.
• Determine indirectly by assuming that the formation of
an ionic compound takes place in a series of steps
known as the Born-Haber cycle.
• The Born-Haber cycle relates lattice energies of ionic
compounds to ionization energies, electron affinities,
and other atomic and molecular properties.
• Lattice energy is an indication of the stability of ionic
compounds, its value can help us rationalize the
formulas of these compounds.
Covalent Bond
• Formation of molecules by the
sharing of electrons between
combining atoms is called
covalency and the bond
formed is called covalent
bond or covalent linkage.
• Compounds containing this
type of linkage are called
covalent compounds.
• In covalent bond formation the
inert gas configuration of the
two concerned atoms is
achieved by sharing equal
number of electrons.
Covalent Bond
• The sharing of electrons to form chemical bond
between two atoms is described by showing
pairs of electrons between the bonded atoms.
• If one pair of electrons is shared, the bond
formed is called single bond, whereas sharing of
two or three pairs of electrons leads to the
formation of double or triple bonds respectively.
• Due to sharing of electrons, both the bonding
atoms acquire inert gas configuration.
Covalent Bond
• Example 1: the formation of hydrogen molecule
from 2 hydrogen atoms.
• The covalent bond is represented by (–).
Covalent Bond
• Example 2: the formation of CCl4
• Example 3: the formation of CH4
Some other example of the formation of
covalent bonds:
• Single bond
• Double bonds
• Triple bonds
• In covalent compounds the numerical value of
covalency of any element in the molecule is the
number of electron pairs shared between the atoms.
Thus the valency of hydrogen in H2 is one, oxygen in
O2 is two, nitrogen in N2 is three and carbon in CH4 is
four.
• Two types of covalent bonds:
– Polar covalent bonds
When a bond is formed between unlike atoms,
the bonding electrons will not be equally shared.
Example: HCl
The shared electrons will be shifted more
towards the atom having higher electronegativity
and this will result in the accumulation of a –ve
charge on it. The other atom will carry an
equivalent +ve charge.
• Example HCl:
The chlorine atom acquires small amount of –ve
charge because of its higher electronegativity and
hydrogen has an equivalent +ve charge. The δ+
and δ – represents respectively the small +ve and
–ve charge.
Electronegativity:
The ability of an atom to attract toward itself the
electron in a chemical bond.
– Non-polar covalent bonds.
• When a bond is formed between atoms of the
same element, the bonding electrons are equally
shared on account of equal electronegativity of
the atoms.
• In case of such a bond, the centre of +ve charge
coincides with the centre of –ve charge in the
molecule.
• For example, bonds involved in the formation of
H2, Cl2, O2, N2 etc. are non-polar bonds.
• Covalent compounds possess following general
characteristics:
– Covalent compounds possess definite geometrical
shapes. They exhibit isomerism because covalent
bonds are rigid and possess directional
characteristics.
– Covalent compounds are mostly liquids and gases.
The solid compounds are generally volatile.
– They are highly soluble in organic solvents but slightly
soluble in water. Some compounds like HCl and NH3
readily dissolve in water because they react with
water.
– The melting and boiling points of covalent
compounds are relatively low because the forces
involved in covalent compounds are less strong
than those involved in electrovalent or ionic
compounds.
– These compounds do not contain ions. Therefore,
when dissolved, they do not conduct electricity.
They even do not conduct electricity in the molten
state.
Ionic vs Covalent Compounds
Coordinate Bonds
• When, an atom having a complete octet donates a pair
of free valence electrons to another atom which is short
of two electrons, the resulting bond is known as
coordinate bond. Thus both the atoms acquire inert gas
configurations.
• The atom which donates a pair of electrons is called
‘donor’ and the other atom which accepts the electrons
is called ‘acceptor’. The coordinate bond is similar to
covalent bond except that both the shared electrons are
donated by one atom.
• The formation of covalent and coordinate bond is illustrated
below:
• When one atom furnishes both electrons for the
formation of a covalent bond as described above,
the process is called ‘coordination’.
• Since one atom donates an electron pair and the
other accepts, the molecule acquires polarity.
These bonds, therefore, are also known as ‘semi
polar bonds’ or ‘dative bond’.
• Coordinate bond is represented by the symbol
().
• Example:
• Following are the main characteristics of coordination
compounds:
– Coordinate compounds, like covalent compounds,
exhibit space isomerism. This is due to directional
characteristics possessed by coordinate linkage.
– The B.P. and M.P. of these compounds have
intermediate value between electrovalent and
covalent compounds.
– They are only slightly soluble in water and most of
them are soluble in organic solvents.
Writing Lewis Structure
•
•
•
•
Write the skeletal structure of the compound showing
what atoms are bonded to what other atoms.
Count the total number of valence electrons present,
referring, if necessary to Lewis dot symbols.
Draw a single covalent bond between the central atom
and each of the surrounding atoms. Complete the
octets of the atoms bonded to the central atom.
If the octet rule is not met for the central atom, try
double or triple bonds between the surrounding atoms
and the central atom, using the lone pairs from the
surrounding atoms.
• Example:
Write the Lewis structure of nitric acid, HNO3, in which
the three O atoms are bonded to the central N atom
and the ionizable H atom is bonded to one of the O
atoms.
• Answer:
Step 1: The skeletal structure of HNO3 is
O
N
O
H
O
• Step 2: The outer-shell electron configurations of N, O,
and H are
N: 2s2 2p3
O: 2s2 2p4
H: 1s1
Thus, there are: 5 + (3 x 6) + 1 = 24, valence electrons to
account for in HNO3.
• Step 3: Draw a single covalent bond between N and
each of the three O atoms and between one O atom and
the H atom. Then we fill in electrons to comply with the
octet rule for the O atoms.
O
N
O
O
H
• Step 4: We see that this structure satisfies the octet rule
for all the O atoms but not for the N atom. Therefore we
move a ione pair from one of the end O atoms to form
another bond with N.
O
N
O
O
H
Formal Charge & Lewis Structure
• Electrons are shared in a bond, we must divide the electrons
in a bonding pair equally between the atoms forming the bond.
The difference between the valence electrons in an isolated
atom and the number of electrons assigned to that atom in a
Lewis structure is called that atom’s formal charge.
• The equation for calculating the formal charge on an atom in a
molecule is given by:
• FC = total number of valence electrons in the free atom – total
number of nonbonding electrons - ½ (total number of bonding
electrons)
• Example:
Write formal charges of the carbonate ion!
• Answer: The Lewis structure for carbonate ion
O
O
C
2-
O
• The FC on the atoms can be calculated as follows:
– The C atom
: FC = 4 – 0 – ½ (8) = 0
– The O atom in C=O: FC = 6 – 4 – ½ (4) = 0
– The O atom in C-0 : FC = 6 – 6 – ½ (2) = -1
– Thus the Lewis formula for CO32- with FC is
O
O
C
O
– Note that the sum of the FC is -2, the same as the charge on
carbonate ion