Chemical
Bonding and Shapes of Molecules
Valence Shell and Valence Electron:
The outermost
shell of an atom is called valence shell and the electrons which belong to the
valence shell are called valence electrons. Valence electrons take part in
chemical combination. The reactivity of an atom depends up on the number of
valence electron.
Octet Theory:
The tendency
of an atom to have eight electrons in their outermost (valence shell) shell is
known as octet rule and the theory related to this rule is known as octet
theory.
The main
points of this rule are.
1. Atoms with
8 electrons in their valence shell (except helium) are chemically stable.
2. Atoms with
less than 8 electrons in their outermost shell are chemically active and take
part in chemical combination.
3. Atoms
having less than 4 electrons in outermost shell usually tend to lose electrons
while atoms having more than 4 electrons in their outermost shell usually tend
to gain electrons and having 4 electrons in their outermost have tendency to
share electrons during chemical combination.
4. Transfer
or sharing of electrons gives a stable electronic configuration of 8 electrons
in valence shell of both combining atoms.
Failure
(exception) of octet rule:
Some
compounds are stable even though constituent atoms do not follow octet rule.
Some of them are:
1. Compounds of hydrogen: Hydrogen atom doesn’t attain 8 electrons rather it attain two electrons (duplet) in its outermost shell in H2, HF, HCl etc.
2. Molecules having incomplete octet: Octet rule could not explain the formation of molecule that contain electron deficient central atoms such as BeCl2, BF3, NO, NO2, etc.
3. Molecules having super octet state:
Octet rule
could not explain the existence of molecules in which central atom has more
than eight electrons such as PCl5, ClF3, FCl3,
SF6, IF7 , etc in which central atoms have 10, 10,10, 12,
14 electrons respectively.
Chemical Bond:
The binding
force between two chemical species (atoms, ions) which is formed by either
sharing or transfer of valence electron(s) is called chemical bonding.
On the basis
of mode of formation, chemical bonds are of following three types
·
Electrovalent
or ionic bond.
·
Covalent
bond.
· Co-ordinate covalent bond or dative bond.
Electrovalent or ionic bond
The chemical bond which is formed by the complete transfer of one or more valence electrons from one atom to other atom is known as electrovalent bond or ionic bond. This type of bond is usually formed between a metal and a nonmetal atom in which metal atom changes into cation (positive ion) and nonmetal changes into anion (negative ion) .For example, bond in NaCl, KCl, MgCl2, CaCl2,NaBr,Na2S , FeCl2, FeCl3, AlCl3 etc.
Characteristics of ionic bond:
·
It is formed by complete transfer of electron between two
combining atoms.
·
It is non directional.
·
There is no fixed geometry of molecule in ionic compounds.
·
Ionic molecules do not exhibit isomerism due to non-direction
nature of ionic bond.
· The number of electrons gained or lost by an atom is known as electrovalence.
Ionic compound
The compounds
which are formed by the complete transfer of electrons from one atom to the
between combining atoms are called as ionic compound or the compounds that are
formed by existence of electrostatic force of attraction or ionic bond between
combining atoms are called ionic compounds.
For example,
NaCl, KCl, NaBr CaCl2 etc.
Characteristics
of ionic compounds
1. Nature of bond: There is existence of ionic bond in the ionic compounds.
2. Physical state: Ionic compounds usually exist in crystalline solid form at room
temperature.
3. Solubility: They are soluble in water and other polar solvent but are
insoluble in organic solvent like benzene, chloroform etc.
4. High melting and boiling point: Ionic
compounds have high M.P and B.P. due to strong electrostatic force of
attraction between cation and anion.
5. Electrical conductivity: Ionic compounds can conduct
electricity in molten state or in their aqueous solution however they can’t
conduct electricity in solid state it is because cations and anions are
strongly attracted by each other and so can’t move freely in solid state while
these ions are far apart to each other and can freely move in the aqueous
solution or in molten state due to less electrostatic force of attraction..
Therefore ionic solid are bad conductor in solid but good conductors in molten
or aqueous solution.
6. Nature of reaction: Ionic compounds show ionic reaction in which ions involve in
chemical combination.
Na+ (aq.) + Cl-(aq.) + Ag+
(aq.)+ NO3- (aq.) à NaNO3
(aq.) + AgCl↓(aq.)
7. Brittleness: Generally ionic compounds are brittle in nature. When force is
applied, similar charged ions come together and repulsion force results the
breaking of crystal.
8. Non directional character: The ionic bond is non directional
in nature since an ion can attract other opposite charged ions equally from any
direction.
Isomerism:
They do not show isomerism due to non-direction nature of ionic bond.
9. Isomorphism: Ionic compounds formed from ions having similar electronic
configuration possess identity of crystalline form which is termed as
isomorphism.
Example:
i. NaF and
MgO
|
Na+ |
F- |
|
(2, 8) |
(2, 8) |
|
Mg+2 |
O-2 |
|
(2, 8) |
(2, 8) |
ii). CaCl2 and K2S
|
Cl- |
Ca+2 |
Cl- |
|
(2,8,8) |
(2,8,8) |
(2,8,8) |
|
K+ |
S-2 |
K+ |
|
(2,8,8) |
(2,8,8) |
(2,8,8) |
Covalent
Bond:
The chemical
bond which is formed by the mutual sharing of valence electrons between two
combining atoms is known as covalent bond. Each combining atom contributes
equal number of electrons for sharing.
Example: Formation of Cl2, F2, H2, O2, N2, H2O, NH3, CH4, HCl molecules etc.
Formation of H2 molecule:
Formation of Cl2 molecule
Formation of ammonia molecule:
Characteristics of covalent bond:
It is formed
by the mutual sharing of valence electrons between two combining atoms.
It is
directional.
There is
fixed geometry of molecule of covalent compound.
Covalent
compound exhibits isomerism due to directional characters of covalent bond
The number of
share electrons by an atom is called co-valency.
Covalent compound:
The compounds
which are formed by the mutual sharing of electrons between combining atoms are
known as covalent compounds. For example, CH4, CCl4, CO2,
H2S, C2H6, NH3 etc.
Characteristics
of covalent compounds
Nature of bond: Covalent compounds consists of covalent bond.
Physical state: Covalent compounds exist in all three states gas, liquid or
liquid.
Low boiling and melting point: covalent compounds usually have
low M.P and B.P since molecules are bounded by weak vanderwaal’s force.
Solubility: Covalent compounds are soluble in organic (non-polar) solvent
like benzene, chloroform but insoluble in polar solvent like water.
Electrical conductivity: covalent compounds are bad conductor since they can’t furnish
free ions but graphite can conduct electricity though it is covalent due to
presence one free electron in each carbon.
Nature of reaction: Molecules involve in chemical combination in which old bonds are
broken and new bonds are formed. They show molecular reaction.
Directional character: covalent bonds are rigid and directional since shared pair of
electrons are localized in fixed direction in the space between the nuclei of
two atoms.
Isomerism: covalent compounds possess isomerism because covalent bonds are
rigid and directional and hence can form different arrangement of atoms in the
space.
Soft and non-brittle: Covalent compounds are usually soft and non-brittle.
Lewis
structure of some covalent compounds:
Co-ordinate covalent bond or
dative bond:
The chemical
bond formed by sharing of valence electrons between two atoms in which shared
pair of electrons is contributed only by one of the two atoms and other atom
simply participates in sharing is known as co-ordinate covalent bond or dative bond.
The atom which donates shared pair of electrons is called donor and it must already
be octet while the atom which accepts electron pair in order to attain octet is
called acceptor. Co-ordinate covalent bond is represented by an arrow
(→) pointing from donor atom towards the acceptor atom.
Examples:
In SO3 molecule:
Characteristics of co-ordinate
covalent bond.
It is formed
by the sharing of a pair of electron between combining atoms in which only one
atom contributes electron for sharing.
It is
directional.
It is semi
ionic bond.
Molecules
exhibit isomerism due to directional character of co-ordinate covalent bond.
Co-ordinate covalent compound:
The compounds
which consist of at least one or more co-ordinate covalent bond in addition
with other covalent bond are known as co-ordinate covalent compound.
Example: NO, NO2,
N2O3, H2SO4
Characteristic
of co-ordinate covalent compounds
Physical state: Generally co-ordinate covalent compounds exist in liquid and
gaseous state.
Solubility: They are sparingly soluble in water and organic solvent since they
semi polar or semi ionic in nature.
Electrical conductivity: They are poor conductor in nature.
M.P. and B.P.: they lower M.P and B.P than ionic compounds but higher than
covalent compounds.
Directional character: Co-ordinate covalent bond is rigid and direction.
Isomerism: They possess isomerism due to rigid and directional bond.
Molecular reaction: They show molecular reaction mechanism.
Bond
Characteristics
i. Bond
length:
Bond length
or bond distance is the average distance between the nuclei of two bonded atoms
in a molecule or chemical species. In general, a single bond is longer than the double bond and double is longer than
triple bond. The bond length is inversely related to bond strength and bond
enthalpy.
|
|
C-C |
C=C |
C |
|
Bond Length |
Single Bond |
>Double
Bond |
>Triple
Bond |
|
Bond
Enthalpy |
Single
Bond< |
Double
Bond< |
Triple Bond |
The bond
length between any two atoms depends upon the electronegativity of bonding
atoms, neighboring substituent, orbital hybridization and resonance. For
example, the bond length of C-H bond in methane and in chloroform is different.
Replacement of hydrogen atom by chlorine atom causes different bonding
environment resulting into different bond length.
Bond length
is measured in the order of Pico meter (pm) or Bohr radius. A Bohr radius is 1amu
length.
1 Bohr radius (1amu) =
52.91 pm
Example:
Common single
bond length of C-C = 154 pm or 2.91 Bohr radius.
ii. Ionic
Character:
Partial ionic character of
covalent compounds:
Covalent compounds which are formed by mutual sharing of electrons between dissimilar atoms having different electro negativity exhibit partial ionic character. In this case, shared pair of electrons is displaced towards more electronegative atom and it acquires partial negative charge while the less electro negative atom acquires partial positive charge and hence covalent molecule shows partial ionic character. For example, HCl, HBr, HI, NH3, H2O etc.
δ+ δ-
H-------Cl
The extent of
ionic character depends upon the difference of electro negativities of two
combining atoms. Greater the difference of electro negativity, greater is the
percentage of ionic characters. It has been observed that a bond has 50% ionic
character and 50% covalent character if the difference in electro negativities
of the participating atoms is 1.7. On the other hand if difference is more than
1.7, ionic character is dominance and if difference is less than 1.7, covalent
character is dominance. Percentage ionic
character and difference in electro negativity is given below.
|
Electro negativity difference |
3 |
2.3 |
1.7 |
1.1 |
0.6 |
0.2 |
|
% ionic
character |
99 |
75 |
50 |
25 |
10 |
1 |
Polar and non-polar covalent
compound:
Polar
covalent compound:
The covalent
compounds that are formed by sharing of electrons between dissimilar atoms of
different electro negativity in which more electronegative atom acquires
partial negative charge and less electronegative atom acquires partial positive
charge, are known as polar covalent bond.
The different poles i.e. partial negative charge and partial positive
charge are developed due to unequal distribution of bonded pair of electrons.
Some examples of polar covalent compounds are HCl, H2O, HBr, NH3,
HI etc.
Non
polar covalent compound:
Covalent
compounds that are formed by sharing of electrons between similar atoms, in
which shared pair of electrons are localized at equal distance from nuclei of
both participating atoms, are called non polar covalent compounds. Such
molecules have no any kinds of charge at atoms. Some examples of non-polar
covalent compounds are H2, Cl2, O2, N2
etc.
iii.
Dipole Moments:
The degree of
polarity or ionic character in a covalent bond is expressed in term of dipole
moment. Dipole moment is defined as the product of the magnitude of charge of
any one of atoms and the distance between them. It is denoted by Greek letter µ
(meu) and it unit is Debye or D.
Mathematically, it can be expressed as:
µ= q × d
Where,
q= magnitude
of charge on any one of atoms (order of 10-10 esu)
d=distance
between both (order of 10-8
cm)
Hence, 1D =
10-10 ×10-8 = 1 ×10-18 esu cm
For HCl
µ=1.03 × 10-18
esu cm
= 1.03D
Application of dipole moment
Determination of polar and non-polar molecules: Molecules
having some value of dipole moment are polar while molecules having zero value
of dipole moment are non-polar. For, e.g. H2, N2, Cl2
etc. are non-polar as their dipole moment is zero and HCl, HBr, NH3 etc.
are polar as they have some value of dipole moment.
Determination of ionic character in covalent compound: The dipole
moment provides idea about ionic character of covalent bond. Higher the value
of dipole moment, larger is the ionic character of covalent compound.
Determination of geometry of molecules: The dipole
moment gives the idea about geometry and shape of molecule.
Molecules
having linear geometry: The resultant dipole
moment of tri atomic molecules like CO2, CS2, etc. is
zero. Therefore, they have linear geometry. (This means each polar bond in such
molecules lies opposite direction so as to cancel dipole moment of each other and
hence give linear geometry).
Special
bonds
1.
Hydrogen bond:
The weak force of attraction between hydrogen covalently bonded with more electronegative element and more electronegative elements like N, O, F etc. in the same or different molecule is known as hydrogen bond.
Conditions for hydrogen bonding
Hydrogen atom
should already be covalently bonded with more electro negative element
The
electronegative element should have smaller size.
Type of hydrogen bond
i.
Intermolecular hydrogen bond:
The weak attraction force between hydrogen (already covalently bonded with more electronegative elements) and more electronegative elements like N, O, F, etc. of different molecules of either same or different compound is called intermolecular hydrogen bond.
Some points:
It results
the association of molecules. So, it has effect on viscosity and surface
tension.
It affects
the boiling point and melting point.
It affects
the solubility of compounds.
Effect of intermolecular H-bond:
i.
It results the association of molecule. So, it increases viscosity and surface
tension
ii.
It increases B.P and M.P of compounds.
iii. It
increases solubility of compounds in water.
ii.
Intra molecular Hydrogen bond:
The hydrogen bond which is formed between hydrogen (already covalently bonded with more electro negative elements) and more electronegative element within the same molecule is known as intra molecular hydrogen bond.
Some points:
It does not
aggregate molecules.
It does not
affect the melting point, boiling point and solubility.
Application of hydrogen bond:
Increase in viscosity and surface tension:
Molecules of
compounds get associated due to intermolecular hydrogen bonds which results
increase in viscosity surface tension of liquid compound. For example, glycerol
is more viscous than alcohol because there is large number of intermolecular
hydrogen bonds between glycerol molecules than in alcohol.
Increase in boiling point and melting:
Boiling and
melting point of compounds increases due to presence of intermolecular hydrogen
bonds. It requires extra amount of heat energy to break down intermolecular
hydrogen bond.
Effect on physical state:
H2O is liquid at room temperature but H2S is gas because H2O can form inters molecular hydrogen bonds between water molecules but H2S can’t form intermolecular hydrogen bonds rather H2S molecules are attracted by very weak Vander Waal’s force of attraction. Therefore, H2O is liquid but H2S is gas at room temperature.
Solubility of covalent compounds:
The covalent
compounds which can form inter molecular hydrogen bonds with water molecules
are soluble in water but the covalent compounds which are unable to form
intermolecular hydrogen bonds with water
molecules are insoluble in water.
Various
organic compounds such as carboxylic acids, alcohols, amines, etc. are soluble
in water as they can form intermolecular hydrogen bonds with water molecules.
Fig:
electron sea model
The strong
electrostatic force of attraction between positive kernels (metallic ion) and
mobile free electrons which binds metal atoms together resulting hardness and
solid properties is known as metallic bond.
According to
Drude (1900) and Lorentz (1916), metal consists of positive charged metallic
ions (kernels) in which valence electrons are moving on the surface of metal as
like gas molecule. So this model is called electron sea model or electron gas
model.
Properties of metal
Metallic luster: When light falls on the metal surface, the free electron absorb
light and get vibrated. These vibrating electrons immediately emit radiation
due to which metal surface appears shinning.
Electrical and thermal conductivity: When the
electricity is applied, the free electrons in the metal start moving from
negative terminal to the positive terminal.
Thermal conductivity: when an end of a metal is heated, the free mobile electrons move
rapidly towards cooler end and transfer heat to the adjacent electron by
colliding with it.
Malleability and ductility: when the stress is applied on
metal, the lattice of kernels slips on each other layers and the free mobile
electron immediately occupy their positions without destroying lattice of
crystal and metal become sheet or wire.
High tensile strength: the property by virtue of which metal can be stretched without
breaking is known as tensile. This property is due to the strong electro static
force of attraction between positive kernel and mobile electrons.
Hardness: Metals are very hard due presence of strong electro static force of attraction between positive kernels and mobile electrons.
Vander Waal’s
force:
The very weak
force of attraction between instantaneous dipoles and induced dipoles developed
in covalent molecules are called Vander Waal’s force or London force.
The rapid movement of electrons results unsymmetrical distribution of electron density in atoms of molecule which leads separation of negative and positive poles temporary. This temporary dipole is called instantaneous dipole which influences the electron distributions of other close molecules. As a result dipoles appear on the molecule which is termed as induces dipole.
Some important points
It is an instantaneous
force of attraction.
Vander Waal’s
force of attraction is weaker than hydrogen bond.
Vander Waal’s
force of attraction increases with increasing surface area of molecules and polarity
in the molecule.
Vander Waal’s
force increase the melting point and boiling point of molecules.
Example:
The order of
boiling and melting point of halogen molecules: I2 > Br2
> Cl2
Order of
surface area: I2 > Br2
> Cl2
Order of Vander
Waal’s force: I2 > Br2 > Cl2
Molecular
solids:
The solids
which consist of molecules as constituent held by weak Vander Waal’s force of
attraction are known as molecular solids. For e.g. ice, sugar, iodine solid,
dry ice, etc.
Properties
i. They have
low m.p. and b.p.
ii. They are
generally soft.
iii. They are
bad conductor of electricity.
iv. They are volatile and have low heat of vaporization.
Resonance:
The
phenomenon in which single molecule can be represented by two or more
structures as a result of delocalization of pi (π) electrons, lone pair of
electrons and odd electron in the molecule is known as resonance and the
different structures of a molecule represented by resonance phenomenon are
known as resonance structure or canonical structure. The two resonating
structures are denoted by double headed arrow and intermediate structure among
the resonance structure is called resonance hybrid.
Example
Resonance in
O3
Resonance in
SO3
Resonance of
SO2
Resonance in
CO3 - -
Resonance of
SO4- -
Resonance in
NO3-
Resonance of
PO4- -
Hybridization:
The
phenomenon of intermixing of the two or more different orbitals of comparable
energies of an atom to produce new hybrid orbitals of same shape, size and
equivalent energy is known as hybridization.
Conditions
of hybridization
·
The orbital should have comparable energy undergo in
hybridization.
·
Orbital of valence shell take part in hybridization.
·
The orbitals of same atom take part in hybridization
·
Both half-filled as well as completely filled orbitals can take
part in hybridization.
·
The hybrid orbitals have equal energies and identical shape.
· The number of hybrid orbitals is equal to the number of orbital taking part in hybridization.
Types
of hybridization.
Depending
upon the types of orbitals involved in the hybridization process, hybridization
is classified into following three types.
1. sp hybridization:
The hybridization in which one s-orbital and one p-orbital of comparable
energies of an atom intermix to produce two sp-hybrid orbitals of same shape,
size and equivalent energy is known as sp hybridization.
Feature
of sp-hybridization
·
Molecule has linear geometry
·
Bond angle of such molecule is 1800
·
S and P characters are 50-50% respectively
·
Atom with sp hybridization has relative high electro negativity
Examples:
Formation of
BeF2 molecule
Be=1s2 2s2
Two half-filled
hybrid orbital of Be undergo head to head overlapping with half-filled p-orbital of two fluorine atoms to form linear BeF2 molecule.
Orbital Geometry of BF2
SP2 hybridization:
The process of intermixing of one s-orbital and two p-orbitals to form three sp2 hybrid orbitals of same shape, size and equivalent energies is known as sp2 hybridization.
Features
of sp2 hybridization
Molecule has
trigonal planar geometry
Bond angle in
such molecule is 1200
S and p
characters are 33.33% and 66.67% respectively
Example
Formation of
BF3 molecule:
B: 1S2 2S2 2P1
Three half-filled hybrid orbitals of B undergo head to head overlapping with half-filled orbital of three fluorine atoms to form trigonal planar molecule.
Trigonal planar
|
F |
|
|
|
F |
|
|
|
|
|
B |
|
|
|
120 |
3. SP3 hybridization:
The
phenomenon which involves intermixing of one s-orbital with three p-orbitals of
the same atom to form four sp3 hybrid orbitals is known SP3.
These four hybrid orbitals lie in tetrahedral arrangement.
Features
of sp3 hybridization
Molecule has
tetrahedral structure.
Bond angle of
the molecule is 109.5o
3s and p
characters are 25% and 75% respectively.
Examples: CCl4, CH4, H2O, NH3
molecules.
Formation of CH4 molecule
Carbon atom
undergoes sp3 hybridization at excited state to form four sp3
hybrid orbitals. Each hybrid orbital contains single electron and combines with
1s orbital of H atoms to form CH4 molecule. The geometry of molecule
is tetrahedral with bond angle 109.5o between each C-H bond.
Deviated
geometry of molecules with lone pair of electrons:
i.
Molecule of water
Oxygen atom undergoes sp3 hybridization to form four sp3 hybrid orbitals. Two hybrid orbitals contain single electron in each which combine with 1s orbital of H atoms to form two O-H sigma bonds and other two hybrid orbitals contain non bonded pair of electrons (lone pair) in each. Presence of two lone pair of electrons causes repulsion between them which reduce bond angle between two O-H bonds to 104.5o from 109.5o. Hence, water molecule is angular (V- shape) shape.
ii. Molecule
of H2S:
Sulphur atom undergoes sp3 hybridization to form four sp3 hybrid orbitals. Two hybrid orbitals contain single electron in each which combine with s orbital of H atoms to form two S-H sigma bonds and other two hybrid orbitals contain non bonded pair of electrons (lone pair) in each. Presence of two lone pair of electrons causes repulsion between them which reduce bond angle between two S-H bonds to 920 from 109.5o. Hence, H2S molecule is angular (V- shape) shape.
Molecule of ammonia
In ammonia molecule, nitrogen atom has four sp3 hybrid orbitals. Among
them, three hybrid orbitals contain single electron in each which combine with
1s orbital of three H atoms to form three N-H sigma bonds and remaining one
hybrid orbital contains non bonded pair of electrons (lone pair). Presence of
lone pair of electrons causes repulsion between lone pair and bond pair
electrons. Since repulsion of electrons follow the order of L.PL.P>
L.P.-B.P>B.P-B.P, therefore bond
angles between N-H bonds reduces to 107.5o from 109.5o. Hence, ammonia molecule
has trigonal pyramidal structure.
iii. Molecule
of ethene (CH2=CH2):
Both carbon atoms of ethene have three sp2 hybrid orbitals. One sp2 hybrid orbital of each carbon atom combine together to form C-C sigma bond and other two hybrid orbitals of both carbon atoms combine with 1s orbital of two H atoms to form two C-H sigma bonds to each carbon. The remaining un-hybridized p-orbital of both carbons overlap laterally above and below the plane of sigma bonds to form one pi bond between two carbon atoms.
Fig.:
Orbital geometry of ethene
iv. Molecule
of ethyne:
Both carbon atoms of ethyne molecule have two sp hybrid orbitals. One sp hybrid orbital of each carbon atom combine together to form C-C sigma bond and other one hybrid orbital of each carbon atom combine with 1s orbital one H atom to form one C-H sigma bond to each carbon atom. The remaining two un-hybridized p-orbitals of both carbon atoms overlap laterally above and below the plane of sigma bonds to form two pi bonds between two carbon atoms.
Fig.:
Orbital geometry of ethyne
Types of
Covalent bond:
On the basis
of types of overlapping of orbitals, covalent bonds are of two types
1. Sigma
(σ) bond:
The covalent
bond formed by head –to-head overlapping of two half-filled orbitals along
inter nuclear axis is called sigma bond.
Feature of sigma bond.
·
It is stronger bond than pi bond.
·
It may or may not be polar.
·
It doesn’t involve in resonance.
·
It can be rotated.
Formation of sigma bond
Sigma bond is
formed in three different ways by overlapping of half-filled orbitals. Both s
and p orbitals involve in formation of sigma bond.
i. S-S overlap: Half-filled
S-orbital of one atom overlaps with half-filled s-orbital of another atom to
form Sigma bond.
Example: Formation of H2 molecule
ii. S-P overlap: Half-filled
s-orbital of one atom overlaps with half-filled P- orbital of another atom
along inter nuclear axis to form sigma bond.
Example: Formation of HCl molecule
iii. P-P overlapping:
One
half-filled P-orbital of an atom overlaps with half-filled P-orbital of another
atom above and below the inter-nuclear axis to form sigma bond.
Example: Formation of F2 molecule, Cl2 molecule
2. Pi (π)
bond:
The covalent bond formed by sideways (lateral) overlapping of both lobes of two half-filled p-orbitals above and below the inter-nuclear axis is known as Pi (π) bond.
Feature of Pi bond
·
It is weaker bond than the sigma bond.
·
It is formed by lateral overlapping of two half-filled p-orbitals.
·
Overlapping occurs above and below inter nuclear axis.
·
It can’t be rotted.
·
It involves in resonance.
Difference
between pi and sigma bonds:
|
Sigma bond |
Pi bond |
|
It occurs head to head overlapping of two half-filled orbitals along
nuclear axis. |
It occurs
laterally overlapping of two half filed p orbital perpendicular to the inter-nuclear
axis. |
|
It is stronger than the pi bond due to maximum overlapping. |
It is
weaker than the sigma bond due to low extent of overlapping. |
|
It is formed by S-S, S-P and P-P overlapping. |
It is
formed by p-p overlapping only and s orbital can’t form pi bond. |
|
Electrons of sigma bond are located between nuclei of bonded atoms. |
Electrons
of pi bond are located above and below the inter-nuclear axis. |
|
It can be rotted. |
It can’t be
rotated. |
|
It does not involve in resonance. |
It involve
in resonance. |
Valence
Shell Electron Pair Repulsion (VSEPR) theory :
This theory
was proposed by Sidgwick and Powell in 1940 and further developed by Gillespie
and Nyholm in 1957.This theory explains molecular shape and bond angles more
exactly on the basis of electrostatic attraction between electron pairs around
the central atom.
VSEPR theory
states that, “The electron pairs (both
lone pairs and shared pairs) surrounding the central atom will be arranged in
space as far apart as possible to minimize the electrostatic repulsion between
them”.
Postulates
of VSEPR theory
1. There is
spatial arrangement of electrons pairs (lone pairs and shared pairs) around the
central atom.
2. Atoms of
molecules are arranged in such a way that repulsion between electron pairs is
minimum as much as possible.
3. The order of
repulsion between electron pairs are L.P.-L.P.>L.P.-B.P.>B.P.-B.P.
4. If lone pair
of electrons is/are present in the central atom, then the structure of molecule
gets deviated from ideal geometry.
5. The shape of
molecule is determined by the repulsion between all types of electron pairs
present around the central atom.
6. The molecules or ions containing 2,3,4,5 and 6 bonded electron pairs in central atom have linear, trigonal planar, tetrahedral, trigonal bi-pyramidal and octahedral geometry respectively.
Geometry
of some molecules in the light of VSEPR theory
i. Shape of BeF2:
Lewis structure of BeF2
Fig.: Linear
geometry of BeF2
ii. Shape of BF3
The central atom boron contains three bond pairs of electrons. To minimize repulsion, three bonded pairs of electrons get stretched to assume trigonal planar geometry with bond angle 120o between each B-F bonds.
Fig. Trigonal planar geometry
iii. Shape of CH4
Fig.:Lewis
structure of CH4
Central atom carbon has four bonded pairs of electrons which cause repulsion equally to each other. Therefore, to have a minimum repulsion between them, four bonded pairs electrons are stretched symmetrically in the space and form an angle of 109.5o with each other and hence molecule have regular tetrahedral geometry.
Limitation
of VSEPR theory
It cannot
explain the shape of molecules having very polar bonds.
Example: Li2O
should have the same structure as H2O but actually it is linear.
It fails to
explain the shape of molecules having extensive De-localized pi electron system.
This theory
is unable to predict the shape of certain transition metal complex.
It does not help in determining the exact bond angle.