E-Content Presentation
Subject
Class
Topic
: Chemistry
: +1
: Chemical Bonding & Molecular
Structure
Prepared By
: Mr. Rohit Kashmiri,
PGT Chemistry,
J.N.V. Pandoh,
Distt. Mandi, H.P.
Chandigarh Region
Technical Support : Mr. Rajeev Sharma
Provided By
F.C.S.A. J.N.V. Paprola,
Distt. Kangra, H.P.
INTRODUCTION
Linus Pauling won the
Nobel Prize in Chemistry in
1954 for his work on
chemical bonding.
A chemical bond is the physical
phenomenon of chemical substances being
held together by attraction of atoms to each
other through sharing, as well as
exchanging, of electrons or electrostatic
forces. In general, strong chemical bonds
are found in molecules, crystals or in solid
metal and they organize the atoms in ordered
structures. Weak chemical bonds are classically
explained to be effects of polarity between molecules
which contain strong polar bonds. Some very weak
bond-like interactions also result from induced polarity
London forces between the electron clouds of atoms, or
molecules. Such forces allow the liquification and
solidification of inert gases. At the very lowest strengths
of such interactions, there is no good operational
definition of what constitutes a proper definitional
"bond".
Learning Objectives
 Chemical Inertness of Noble Gases.
 Types of Chemical Bonds.
 Covalent Bonds.
 VSEPR Theory & Shapes of Molecules.
 Polarity in Covalent Bonds.
 Quantum theory of Covalent Bonds: Valence Bond
Theory.
 Molecular Orbital Theory.
 Hybridisation.
 Hydrogen Bonding
Entry Behaviour
1. Why do certain atoms combine to form
molecules whereas other do not?
2. What is the nature of the forces which
hold the atoms together in molecules?
3. Why do atoms have fixed combining
capacity?
Overview of Presentation
In theory, all bonds can be explained by quantum
theory, but in practice, chemical bonds are divided
in several categories. Simplifications of quantum
theory have been developed to describe and
predict the bonds and their properties. These
theories include octet theory, valence bond theory,
orbital hybridization theory, VSEPR theory, ligand
field theory and LCAO -method. Electrostatics
and other physical theories are used to describe
bond polarities and the effects they have on
chemical substances. However, in combination,
they constitute a powerful theory, which can be
applied in almost all of chemistry.
In quantum mechanics, in simplified terms, electrons
are located on an atomic orbital (AO), but in a
strong chemical bond, they form a molecular
orbital (MO). In many theories, these are divided
in bonding, anti-bonding, and non-bonding
orbitals. They are further divided according the
types of atomic orbitals hybridizing to form a
bond. These orbitals are results of electron-nucleus
interactions that are caused by the fundamental
force of electromagnetism. Chemical substances
will form a bond if their orbitals become lower in
energy when they interact with each other.
Different chemical bonds are distinguished that
differ by electron cloud shape and by energy
levels.
Vocabulary
Glossary of Terms
1.
2.
3.
Atomicity of a gas: The number of atoms present in the molecule of
a gas is called its atomicity.
Bond dipole moment ( m ).:A covalent bond between two atoms of
different elements is called a polar covalent bond . A polar bond is
partly covalent bond and partly ionic. The percentage of ionicity in a
covalent bond is called percentage ionic character in that bond .
The ionic character in a bond is expressed in terms of bond dipole
moment ( m ).
BORN-HABER CYCLE: This thermochemical cycle was devised
by Born and Haber in 1919. It relates the lattice energy of a
crystalline substance to other thermochemical data. The Born-Haber
cycle is the application of Hess's law to the enthalpy of formation of
an ionic solid at 298 K.
4. Chemical bond : The chemical force which keeps the atoms in any
molecule together is commonly described as a chemical bond.
5. Chemical compounds :Compounds are generally called chemical
compounds because they are formed due to the chemical combination
of the combining element.
6. Covalent bond : The Bond formed by Mutual sharing of electrons
between the combining atoms of the same or different elements is
called covalent bonds.
7. Double covalent bond :The bond formed between two atoms due to
the sharing of two electron-pairs is called a double covalent bond or
simply a double bond. It is denoted by two small horizontal lines (=)
drawn between the two atoms, e.g., O = O, O = C = O etc.
8. Electronegative or nonmetallic character :The tendency of an
element to accept electrons to form an anion is called its non metallic
or electronegative character.
9. ELECTRONEGATIVITY :The relative tendency of an atom in a
molecule to attract a shared pair of electrons towards itself is termed its
electronegativity.
10. Electronic configuration:The distribution of electrons amongst
various energy levels of a atom is termed its electronic configuration
11. HYBRIDISATION :The process of mixing of the atomic orbitals to
form new hybrid orbitals is called hybridisation.
12. Hybrid orbitals :According to the concept of hybridisation, certain
atomic orbitals of nearly the same energy undergo mixing to produce
equal number of new orbitals. The new orbitals so obtained are called
hybrid orbitals.
13. Hybridisation in carbon :Carbon shows sp 3 hybridisation in alkanes,
sp 2 hybridisation in alkenes and sp hybridisation in alkynes.
14. HYDROGEN BOND :The bond between the hydrogen atom of one
molecule and a more electronegative atom of the same or another
molecule is called hydrogen bond.
15. Ionic (or Electrovalent) bond : An ionic (or electrovalent) bond is
formed by a complete transfer of one or more electrons from the atom
of a metal to that of a non-metal.
16. LATTICE ENERGY (Lattice Enthalpy) :The strength of binding
forces in solids is described by the term lattice enthalpy ( D L H )
(earlier the term lattice energy was used). The molar enthalpy change
accompanying the complete separation of the constituent particles that
composed of the solid (such as ions for ionic solids and molecules for
molecular solids) under standard conditions is called lattice enthalpy (
D L H° ). The lattice enthalpy is a positive quantity.
17. Lewis Formula (or Electronic Formula) of a Compound :The formula
showing the mode of electron-sharing between different atoms in the molecule
of a compound is called its electronic formula or Lewis formula.
18. Metallic crystals : In metallic crystals, the valence electrons of all the atoms
form a pool of mobile electrons. The nuclei with their inner electrons (called
Kernels) are embedded into this pool of free electrons. Thus, the constituent
particles in a metallic crystal are the positive kernels in a pool of electrons.
19. NON-POLAR COVALENT BOND :When a covalent bond is formed
between two atoms of the same element, the electrons are shared equally
between the two atoms. In other words, the shared electron-pair will lie exactly
midway between the two atoms. The resulting molecule will be electrically
symmetrical, i.e ., centre of the negative charge coincides with the centre of
the positive charge. This type of covalent bond is described as a non-polar
covalent bond. The bonds in the molecules H 2 , O 2 , Cl 2 etc., are non-polar
covalent bonds.
20. OCTET RULE : According to this theory, the atoms tend to adjust
the arrangement of their electrons in such a way that they ( except H
and He ) achieve eight electrons in their outermost shell. This is known
as the octet rule .
21. Pi ( p ) Bond : A covalent bond formed between the two atoms due to
the sideways overlap of their p -orbitals is called a pi ( p ) bond.
22. POLAR COVALENT BOND :When a covalent bond is formed between
two atoms of different elements, the bonding pair of electrons does not
lie exactly midway between the two atoms. In fact, it lies more towards
the atom which has more affinity for electrons. The atom with higher
affinity for electrons, thus, develops a slight negative charge, and the
atom with lesser affinity for electrons a slight positive charge. Such
molecules are called polar molecules. The covalent bond between two
unlike atoms which differ in their affinities for electrons is said to be a
polar covalent bond.
23. RESONANCE : When a molecule is represented by a number of
electronic structures such that none of them can exactly describe all the
properties of the molecule, but each structure has a contribution to it,
then the molecule is termed as a resonance hybrid of all these
structures. Such structures are called resonance structures and such a
phenomenon is called resonance.
24. Resonance hybrid :When a molecule is represented by a number of
electronic structures such that none of them can exactly describe all the
properties of the molecule, but each structure has a contribution to it,
then the molecule is termed as a resonance hybrid of all these
structures.
25. Sigma ( s ) Bond : A covalent bond formed due to the overlap of orbitals of
the two atoms along the line joining the two nuclei (orbital axis) is called
sigma ( s ) bond.
26. Single Covalent Bond : A covalent bond formed by mutual sharing of one
pair of electrons is called a single covalent bond, or simply a single bond. A
single covalent bond is represented by a small line (-) between the two atoms.
27. Triple covalent bond : Bond formed due to the sharing of three electronpairs is called a triple covalent bond or simply a triple bond.
28. Valence electrons : Valence is one of the most important chemical property of
the elements. The chemical behaviour of an element depends upon the number
of electrons in the outermost shell of its atom. The electrons present in the
outermost shell are called valence electrons. The electrons in the outermost
shell are called valence electrons because the electrons in the outermost shell
determine the valence of an element .
29. Valency : The combining capacity of an atom of an element is described in
terms of its valency. It may be defined as,
The number of hydrogen or chlorine or double the number of oxygen atoms
which combine with one atom of the element is termed its valency.
It may also be defined as,
The number of electrons which an atom loses or gains or shares with other
atoms to attain noble gas configuration is termed its valency.
Sub Topic 1:
Chemical Inertness of Noble Gases
Noble gases (He, Ne, Ar, Kr, Xe, Rn), do not form
compounds neither among themselves, nor with
other elements. Xenon however, forms fluoride and
oxyfluoride compounds under drastic conditions.
Now, the question is that why do these elements not
show chemical reactivity? To answer this, let us
consider the electronic configurations of these
elements.
The electronic configurations of these elements are
given below.
From the electronic configurations given above, we see that
All noble gases have their outermost shells completely filled. The atoms
of all other elements which show chemical reactivity have less than
eight electrons in their outermost shell.
Thus, it appears that the chemical reactivity of any element is related to
the number and distribution of electrons in its atom. Theoretical
physicists have shown that certain number of electrons in certain
definite energy levels give stable atoms. The atoms having a total of 2,
10, 18, 36, 54, and 86 electrons are found to be the most stable. These
electronic configurations incidently correspond to those of the noble
gases. So, it is because of their stable electronic configurations that
the noble gases show no chemical reactivity. Logically, it means that
the atoms of all other elements are not so stable, and they tend to gain
stability by acquiring an electronic configuration of the nearest noble
gas element.
The atoms tend to adjust the arrangement of their electrons in such a way
that they (except H and He) achieve eight electrons in their outermost
shell. This is known as the octet rule.
VALENCE ELECTRONS AND VALENCY :
The electrons present in the outermost shell (generally termed as
valence shell) are called valence electrons. The number of
valence electrons in the atoms of certain elements are given
below.
"The number of electrons which an atom loses or gains or shares
with other atoms to attain noble gas configuration is termed its
valency."
ELECTRONIC THEORY OF VALENCY :
The combining tendency of atoms was explained by Kossel and Lewis
(1916) through their theory called electronic theory of valency. The
main postulates of this theory are,
(i) The tendency of an atom to take part in chemical combination is
determined by the number of valence electrons. The valence electrons
are the electrons in the outermost shell of the atom.
(ii) The atoms combine by mutual sharing or by transfer of one or more
electrons. In doing so, each combining atom acquires stable noble gas
electronic configuration having 8 electrons in its outermost shell. This
is called octet rule.
(iii) The number of electrons which an atom loses, gains or mutually
shares to attain noble gas configuration is called its valency. For
example, Li, Be, B and C having respectively 1, 2, 3 and 4 electrons,
have valence of 1, 2, 3 and 4. The elements N, O, F and Ne having 5,
6, 7 and 8 electrons in their outermost shell show common valence of
3, 2, 1 and 0.
Thus, "the common valency of an element is either equal to the
number of valence electrons, or it is equal to 8 minus the number
of valence electrons."
LEWIS ELECTRON DOT SYMBOLS
An American chemist, G.N. Lewis introduced simple notation to denote
the valence electrons in an atom. These notations are called electron
dot symbols or Lewis symbols, (or Lewis structures).
According to this method,
(a) The symbol of the element represents the nucleus along with all the
inner electrons which do not take part in the bond formation.
(b) The dots on the symbol represent the valence electrons. Thus, the
number of dots represents the number of valence electrons.
For example,
(i) An atom of hydrogen contains one electron in its valence shell. So,
its Lewis symbol is
H.
Here, H represents the nucleus of hydrogen and the dot (×) represents one
valence electrons in an atom of hydrogen.
(ii) The electronic configuration of chlorine is 2,8,7. Thus, there are
seven valence electrons. The Lewis symbol of chlorine atom is
Recapitulation Exercise
1. Draw the Lewis symbols for the following
elements.
Na, Ca, B, Br, Xe, As, Ge
2. Why are Noble gases monoatomic?
3. Draw Lewis Symbols of O2- , Mg2+ ions.
4. Define Octet Rule.
Sub Topic 2:
Types of Chemical Bonds
The type of chemical bond developed between the two combining atoms
depends upon the way these atoms acquire a stable noble gas
configuration.
Elements may combine through any one of the following ways to form
stable compounds.
i.
By the transfer of electrons from the atom of an element to the atom
or atoms of another. This gives rise to an ionic (or electrovalent)
bond.
ii.
By mutually sharing the electrons. This gives rise to a covalent
bond.
iii.
By one-sided sharing of electrons. This gives rise to a coordinate
bond.
Ionic Bond
An IONIC BOND is an electrostatic interaction that holds together a
positively charged ion (cation) and a negatively charged ion
(anion). In an ionic bond, one atom loses an electron to another
atom, forming a cation and anion, respectively. And, as
everyone knows, opposites attract.
Formation of NaCl
Sodium chloride
results from ionic
bonding.
In table salt, for example, a valence electron from a sodium
atom is transferred to a chlorine atom, forming Na+ and Cl-.
Because the ions have opposite charges, they are attracted to
each other. The loss of a valence electron and the attraction to
the atom that took it happen simultaneously.
Properties of Ionic (or electrovalent) Bond An ionic or
electrovalent bond has the following characteristics. :
(i) An ionic bond is formed due to the coulombic attraction between the
positively and negatively charged ions.
(ii) An ionic bond is non-directional, i.e., the strength of interaction
between two ions depend upon distance, but not on the direction.
(iii) An ionic bond gets broken when the substance is dissolved in a polar
solvent such as water, or when the substance is melted.
Formation of Some Typical Ionic Compounds
The formation of some ionic compounds is explained below.
1. Formation of magnesium chloride, (MgCl2).
The electronic configuration of magnesium (At. No. 12) is 2,8,2. So, it
has two electrons in its valence shell. The electronic configuration of
chlorine (At. no. 17) is 2,8,7. So, it has seven valence electrons. In
terms of the Lewis (electron dot) structures, one can write,
2. Formation of aluminium fluoride (AlF3).
Factors Influencing the Formation of an Ionic Bond
Formation of an ionic bond is favoured by,
(i) Low ionisation enthalpy of the metallic element which forms the cation.
(ii) Large electron gain enthalpy (electron affinity) of the non-metallic
element which forms the anion.
(iii) Large lattice energy, i.e., the smaller size and higher charge of the ions.
Lattice Energy
The strength of binding forces in solids is described by the term lattice enthalpy
(DLH) (earlier the term lattice energy was used). The molar enthalpy change
accompanying the complete separation of the constituent particles that composed of
the solid (such as ions for ionic solids and molecules for molecular solids) under
standard conditions is called lattice enthalpy (DLH).
The lattice enthalpy is a positive quantity. For example, The enthalpy
change for the reaction, under standard conditions.
NaCl(s)
Na+(g)+ Cl-(g)
is the lattice enthalpy of NaCl(s), (DLH(NaCl(s)) and that for the reaction,
H2O(s)
H2O(g)
is the lattice enthalpy of the molecular solid ice.
The lattice enthalpy of a molecular solid is the same as its standard
enthalpy of sublimation. The lattice enthalpy of a metal is the same
as its enthalpy of atomisation.
The enthalpy change for the reaction,
Na+(g) + Cl-(g)  Na+Cl-(crystal) + lattice Enthalpy (U).
is equal to -DLH (NaCl,s), i.e., heat equal to the lattice enthalpy is released
during the formation of crystalline sodium chloride from gaseous
ions.
Lattice enthalpies (energies) are usually estimated from the
thermochemical data using the Born-Haber cycle or by theoretical
calculations. It depends upon the following factors:
1. Size of the ions. Smaller the size of the ions, lesser is the internuclear
distance. Consequently the inter-ionic attractions will be high & the
lattice enthalpy will also be large.
2. Charge on the ions.: Directly proportional to the magnitude of charge
on the ions.
BORN-HABER CYCLE
This thermochemical cycle was devised by Born and Haber in 1919. It
relates the lattice energy of a crystalline substance to other
thermochemical data. The Born-Haber cycle is the application of
Hess's law to the enthalpy of formation of an ionic solid at 298 K.
Formation of crystalline sodium chloride form sodium metal and chlorine
gas can be described by the reaction.
Na(s) + ½ Cl2(g)  NaCl (crystal) DrH = DfH = - 411 kJ mol-1 (energy evolved)
This overall reaction can be considered to proceed in a stepwise manner as follows
The signs of the energy involved in each step follow the rule that energy evolved is
negative and energy absorbed is positive. These steps are summarized in Fig. 6.3.
From the Born-Haber cycle the value for any one of the steps can be calculated if data
for all the other steps are known.
Recapitulation Exercise
1. Name the factors which favour the
formation of an ionic bond.
2. Out of MgO and NaCl, which has higher
lattice energy and why?
3. Draw the Born-Haber for a simple ionic
solid such as MX
Sub Topic 3:
Covalent Bonds
A covalent bond is formed between two atoms (similar or dissimilar) by a
mutual sharing of electrons. The shared pairs of electrons are counted
towards the stability of both the participating atoms. A covalent bond
is defined as the force of attraction arising due to mutual sharing
of electrons between the two atoms. The combining atoms may share
one, two or three pairs of electrons.
When the two atoms combine by mutual sharing of electrons, then each of
the atoms acquires stable configuration of the nearest noble gas. The
compounds formed due to covalent bonding are called covalent
compounds.
Covalency
The number of electrons which an atom contributes towards mutual
sharing during the formation of a chemical bond is called its covalency
in that compound. Thus, the covalency of hydrogen in H2 (H - H, H
H) is one; that of oxygen in O2 is two (O = O, O O), and that of
nitrogen in N2 is three (N  N, N N)
Characteristics of a Covalent Bond A covalent bond has the
following characteristics.
(i) Mode of formation. Covalent bonds are formed due to mutual
sharing of one or more pairs of electrons.
(ii) Directional character. Covalent bonds are directional in nature.
This is because in a covalent bond, the shared pair of electrons
remains localised in a definite space between the nuclei of the two
atoms. This gives a directional character to the covalent bond.
Single Covalent Bond
A covalent bond formed by mutual sharing of one pair of electrons is
called a single covalent bond, or simply a single bond. A single
covalent bond is represented by a small line (-) between the two
atoms.
Formation of ammonia (NH3).
The electronic configurations of nitrogen and hydrogen are
N
1s2 2s2 2p3
or
2,5
H
1s1
or
1
Thus, each nitrogen atom requires three more electrons to acquire a stable
noble gas configuration. On the other hand, each H-atom requires only
one electron to achieve the stable helium configuration. This is done
by mutually sharing three pairs of electrons between one nitrogen and
three hydrogen atoms, as shown below.
The unshared pair of electrons on the nitrogen atom (in ammonia molecule) is
not involved in bond formation and is called a lone pair of electrons. Since,
lone pair of electrons does not take part in bonding, hence it is also called
non-bonding pair of electrons.
MULTIPLE COVALENT BONDS:
The covalent bonds developed due to mutual sharing of more than one
pairs of electrons are termed multiple covalent bonds. These are,
Double covalent bond.
The bond formed between two atoms due to the sharing of two
electron-pairs is called a double covalent bond or simply a double
bond. It is denoted by two small horizontal lines (=) drawn between
the two atoms, e.g., O = O, O = C = O etc.
Triple covalent bond.
Bond formed due to the sharing of three electron-pairs is called a
triple covalent bond or simply a triple bond. Three small horizontal
lines between the two atoms denote a triple bond, e.g., N  N, and H C  C H (acetylene).
Formation of Molecules Having Double & Triple Bonds
Some typical molecules having double & triple bonds are described
below.
Formation of oxygen (O2) molecule.
Oxygen molecules
shares two electrons to
make a double covalent
bond.
Formation of ethyne (C2H2) molecule.
In a molecule
of carbon, two
atoms share
three
electrons -- a
triple covalent
bond.
Comparison Between Single, Double and Triple Covalent Bonds
Single, double and triple covalent bonds differ from each other in the
following ways.
(i) Single bond is formed by the sharing of one electron pair, (two electrons),
double bond is formed by the sharing of two electron pairs, (four
electrons), whereas a triple bond involves sharing of three electron pairs,
(six electrons).
(ii) In a triple bond, six electrons attract the nuclei with greater force. This
decreases the distance of separation between the two nuclei. In a double
bond, four electrons attract the nuclei with a relatively lesser force, and in
a single bond, two electrons hold the nuclei with a still lesser force.
Therefore, the bond lengths follow the order,
Triple bond length < Double bond length < Single bond length
Since, a shorter bond means greater bond strength hence, the energy required
to separate the bonded atoms (called bond energy) follows the order,
Triple bond > Double bond > Single bond
Formation of a covalent bond is favoured by,
(i) High ionisation enthalpy (or energy) of the combining elements.
(ii) Nearly equal electron gain enthalpies (or electron affinities) and equal
electronegativities of the combining elements.
(iii) High nuclear charge and small atomic size of the combining elements.
COORDINATE COVALENT BOND
Coordinate bond is formed when the shared electron-pair is provided by one of the
combining atoms. The atom which provides the electron-pair is termed as the
donor atom, while the other atom which accepts it, is termed as the acceptor
atom.
The bond formed when one-sided sharing of electrons take place is called a
coordinate bond. Such a bond is also known as dative bond. A coordinate
bond is represented by an arrow () pointing towards the acceptor atom.
Formation of Coordinate Bond During the Formation of a Molecule or
Molecular ion
The formation of such a bond is illustrated through some examples given
below.
(i) Formation of ammonium (NH4+) ion.
During the formation of ammonium ion, nitrogen is the donor atom, while H+
is the acceptor ion as shown below.
EXCEPTIONS TO THE OCTET RULE
The octet rule is very useful for describing bonding in a large
number of compounds. However, there are many
exceptions to this rule.
(i) Where duplet is formed.
A hydrogen atom has only one electron in its valence shell.
It needs one more electron to fill its valence shell. The
completed shell has the electronic arrangement of the
noble gas helium. In this case, therefore, an octet is not
completed, but we still get a stable molecule. The electron
dot structures of a few molecules containing hydrogen
atom are shown below:
(ii) Where the octet remains incomplete.
The elements of group 1,2 and 13 contain less than four electrons in
their valence shell. These elements therefore, cannot achieve an octet
by electron sharing. As a result, therefore these elements should not
form covalent compounds. But, elements of these groups form some
covalent compounds. Boron halides (BF3 and BCl3) are covalent
compounds where the octet is incomplete (only six electrons surround
the boron atom).
These compounds are thus electron-deficient compounds.
(iii) Where the octet is expanded.
The elements belonging to groups 15, 16 and 17 have more than four
electrons in their outermost shell. The elements of these groups form
stable compounds in which there are more than eight electrons around
the central atom. For example, PF5, PCl5 and SF6 are some typical
compounds of this type.
The Cl and F atoms have 7 electrons in their valence shell. Therefore, they
need one more electron each to attain the noble gas configuration.
Phosphorus atom has five, and sulphur atom has six electrons in their
valence shell. Then, the Lewis structures of PF5 PCl5, and SF6 are
written as follows.
Thus, P and S atoms have expanded their shells to accommodate
more than eight electrons. Here again, octet rule is violated, but the
compounds formed are stable.
Recapitulation Exercise
1.
2.
3.
4.
5.
6.
7.
Give one example of a compound containing double bond &
one containing a triple bond.
Describe a coordinate bond, giving one example. How does it
differ from the Covalent bond?
How does bond multiplicity affect the bond length?
Explain the term Electrovalency & Covalency.
Write one main difference between an ionic and a covalent
bond.
Where are the bonding electrons most likely to be found in a
diatomic covalent molecule?
What happens to the valence electrons when a covalent bond
is formed between two atoms?
Sub Topic 4:
VSEPR Theory & Shapes of Molecules
The VSEPR theory was proposed by R.J. Gillespie and R.S. Nyholmm in
1957. This theory was developed to predict the shapes of the
molecules in which the atoms are bonded together with single bonds
only.
This theory is based on the repulsions between the electron-pairs in the
valence-shell of the atoms in the molecule. The main postulates of
the VSEPR theory are:
(i) The geometry of a molecule is determined by the total number of
electron pairs (bonding and non-bonding) around the central atom of
the molecule. The shape of the molecule depends upon the
orientation of these electron pairs in the space around the central
atom.
(ii) The electron pairs (shared, or lone pairs) around the central atom in a
molecule tend to stay as far away from each other as possible so as to
minimize the repulsion forces between them.
(iii) The strength of repulsions between different electron pairs follows
the order:
Lone pair - Lone pair > Lone pair - Shared pair > Shared pair
Shared pair
The shared pairs of electrons are also called bond pairs of electrons.
The presence of lone pair(s) of electrons on the central atom causes some
distortions in the expected regular shape of the molecule.
Predicting the Shape of Molecules on the Basis of VSEPR Theory
According to the VSEPR theory, the geometry of a molecule is determined by
the number of electron-pairs around the central atom. So, to use this theory
for predicting the shapes of molecules just count the number of electron
pairs (both, shared and lone pairs). The use of this theory in predicting the
shapes of molecules is illustrated below by taking a typical molecule of the
type ABn, where A is the central atom, B atoms are bonded to A by single
electron pair bonds (single covalent bonds), and n is the number of B atoms
bonded to one atom of A.
For the sake of easier understanding we have divided molecules into various
categories
Shapes of the molecules having only the bond (shared) pairs of electrons
(i) Molecules with two bond pairs.
In a molecule having two bond pairs of electrons around its central atom, the
bond pairs are located on the opposite sides (at an angle of 180), of the central
atom so that the repulsion between them is minimum. Such molecules are
therefore linear. For example, in a molecule of the type AB2, in which the
central atom A has two electron pairs, the two electron pairs are located on
either side of A. Thus, the molecule AB2 takes a linear geometry. Some
molecules which show linear geometry are, BeF2 (beryllium fluoride), BeCl2
(beryllium chloride), BeH2 (beryllium hydride), ZnCl2 (zinc chloride), and
HgCl2 (mercuric chloride).
(ii) Molecules with three bond pairs.
In a molecule having three bond pairs of electrons around its central
atom, the electron pairs form an equilateral triangular arrangement
around the central atom. Thus, the three bond pairs are at 120C with
respect of each other. Therefore, the molecules having three bond pairs
around its central atom have trigonal planar (or triangular planar)
shape.
For example, in a molecule of the type AB3, the three bond pairs of
electrons are located around A in a triangular arrangement. Thus, the
molecule AB3 has a triangular planar geometry. Some molecules
which show triangular planar geometry are; BCl3, BF3 etc.
(iii) Molecules with four bond pairs (AB4 TYPE).
(iv) Molecules with five bond pairs.
A molecule having five bond pairs around its central atom has a
triangular bipyramidal shape.
(v) Molecules with six bond pairs.
The molecules of the type AB6 are octahedral. The molecule SF6 has
an octahedral geometry.
Shapes of the molecules having bond pairs and lone pairs of electrons
The pair of electrons in the valence shell of an atom which is not
involved in bonding is called lone pair of electrons. For example, the
nitrogen atom in ammonia molecule has one lone pair of electrons;
the oxygen atom in water molecule has two lone pairs of electrons. We
describe a few examples of the molecules having one or more lone
pairs of electrons.
(i) Molecules having three bond pairs and one lone pair.
A molecule of ÄB3 type having three bond pairs and one lone pair
has a triangular pyramidal shape. Typical molecules of this type are
NH3, NF3, PCl3, H3O+ etc.
(ii) Molecules with two bond pairs and two lone pairs.
The four electron pairs (two bond pairs + two lone pairs) are distributed
tetrahedrally around the central atom as shown in Fig. 6.4, for a molecule .
The two lone pairs on the central atom repel the bond pairs slightly inwards
due to greater lone pair - bond pair repulsion. As a result, the bond angle in
such a molecule is less than the tetrahedral value of 10928.
These two representations of the H2O molecule show the electron density
as a phantom shading (left) and contour lines (right.)
Note how most of the negative charge is concentrated around the oxygen
atom.
(iii) Molecules with four bond pairs and two lone pairs.
The four bond pairs are distributed in a square planar distribution. The
two lone pairs are in a direction at right angles to this plane. Thus,
giving a square planar shape to such molecules. Examples include
ICl4-, XeF4 and [Ni(CN)4]2-.
(iv) Molecules with Five Bond Pairs & one lone pair:
(iv) Molecules with Four Bond Pairs & one lone pair, three
bond pairs & two lone pairs and two bond pairs & three
lone pairs respectively:
Recapitulation Exercise
1.
2.
3.
4.
Arrange the following according to the increasing bond angles
in them: NH3, H2O and CH4.
The HPH angle in PH3 is smaller than the HNH angle in NH3.
Why?
Predict the shapes of the following molecules following the
VSEPR theory.
(i) Ammonia (NH3) molecule (ii) Water (H2O) molecule.
Predict the shapes of the following molecules using the
valence shell electron pair repulsion (VSEPR) theory. BeCl2,
SiCl4, AsF5, H2S, HgBr2, PH3, GeF2
Sub Topic 5:
Polarity in Covalent Bonds
Depending upon the chemical nature of the combining elements, the following
two types of covalent bonds are formed.
Non-polar Covalent Bond
When a covalent bond is formed between two atoms of the same element, the
electrons are shared equally between the two atoms. In other words, the shared
electron-pair will lie exactly midway between the two atoms. The resulting
molecule will be electrically symmetrical, i.e., centre of the negative charge
coincides with the centre of the positive charge. This type of covalent bond is
described as a non-polar covalent bond. The bonds in the molecules H2, O2,
Cl2 etc., are non-polar covalent bonds.
Polar Covalent Bond
When a covalent bond is formed between two atoms of different elements, the
bonding pair of electrons does not lie exactly midway between the two atoms.
In fact, it lies more towards the atom which has more affinity for electrons.
The atom with higher affinity for electrons, thus, develops a slight negative
charge, and the atom with lesser affinity for electrons a slight positive charge.
Such molecules are called polar molecules. The covalent bond between two
unlike atoms which differ in their affinities for electrons is said to be a polar
covalent bond.
The compounds having polar bonds are termed
polar compounds. Polar substance do not
conduct electricity when in their pure forms,
but give conducting solutions when dissolved
in polar solvents.
The polarity in a bond arises due to the difference in the electronegativities of the combining atoms. Therefore, greater is the
difference in the electronegativity values of the combining atoms,
greater is the polar character in the bond so formed. For example,
in the series H - X (X = F, Cl, Br, I), the electronegativity difference
between H and X atom follows the order,
H - F > H - Cl
> H - Br
> H
-I
Electronegativity difference: (4 - 2.1) (3.0 - 2.1)
(2.8 2.1) (2.5 - 2.1)
1.9
0.9
0.7
0.4
Therefore, the polarity in the H - X bond follows the order,
H - F > H - Cl > H - Br
> HI
i.e., H - F bond is the most polar and H - I bond is the least polar in this
series of compounds
IONIC CHARACTER OF BONDS IN POLAR MOLECULES
An ionic bond is formed due to complete transfer of electrons from one atom to
another. A covalent bond is formed due to mutual sharing of electrons. A
covalent bond between two atoms of different elements is called a polar
covalent bond. A polar bond is partly covalent bond and partly ionic. The
percentage of ionicity in a covalent bond is called percentage ionic character
in that bond. The ionic character in a bond is expressed in terms of bond
dipole moment (m).
The dipole moment of a bond depends upon the difference in the electronegativity
of the two atoms held together by the chemical bond.
The dipole moment of two equal and opposite charges is given by the product of
the charge and the distance separating them. Thus,
Dipole moment, m = Charge (q) × Distance of separation, (r) = q × r
The dipole moment is a vector quantity, and is represented by an arrow showing
the direction from the positive to the negative end of the dipole. The length of
the arrow represents the magnitude of the dipole moment
For non-polar bonds, there is no charge separation. Therefore, q = 0. As a result, therefore
the dipole moment of a non-polar bond is zero.
Unit of dipole moment. The dipole moments are expressed in the units of Debye (D).
1 Debye = 3.338 × 10-30 m C
The ionic character in a bond is related to the dipole moment of the bond by the relationship
given below.
Ionic character in a bond increases with the increase in the electronegativity difference
between the combining atoms as shown below.
Ionic character in H - Cl bond. The HCl molecule has a dipole moment of 1.07 D. The H
- Cl bond length is 127.5 pm (= 1.275 × 10-10 m). Now, if a complete transfer of
electrons from H to Cl is assumed, then the
Dipole moment of H+ - Cl- = (Charge on an electron) × (Bond length)
mionic = = 6.12 D
Thus, had the bond between H and Cl been 100% ionic, the dipole moment of HCl molecule
should have been 6.12 D. But the actual dipole moment (mobs) of HCl molecule is 1.07
D*. So,
Ionic character in HCl bond =
× 100 =
× 100 = 17%
Thus, the estimated ionic character in H - Cl bond is about 17%.
A bond with an ionic character of 50% or more is termed as an ionic
bond. As a simple estimate, the bond between the two atoms with
electronegativity difference of 2 is an ionic bond. A bond between the
two atoms with electronegativity difference of less than 1.5 is usually
covalent.
Dipole Moment of Polyatomic Molecules
So the net dipole is
Zero.
Recapitulation Exercise
1.
2.
3.
4.
The molecule SO2 has a dipole moment. Is the molecule linear or
bent? Explain your reasoning.
Predict the dipole moment of
(i) a molecule of the type AX4 having square planar geometry.
(ii) a molecule of the type AX5 having a trigonal bipyramidal
geometry.
(iii) a molecule of the type AX6 having an octahedral geometry.
Which of the bonds given below is the most and which is the
least polar: S - O, Cl - Cl or
Cl - O? Justify your answer.
Should a polar molecule contain a polar bond?
Sub Topic 5:
Quantum theory of Covalent
Bonds: Valence Bond Theory
The first quantum mechanical model to explain the nature and stability of a
covalent bond was formulated by Heitler and London 1927. This theory
was then modified by Pauling and Slater in 1931. This theory is commonly
known as the Valence-bond theory.
The main postulates of the valence bond theory are:
i.
A covalent bond is formed due to the overlap of the outermost halffilled orbitals of the combining atoms. The strength of the bond is
determined by the extent of overlap.
ii.
The two half-filled orbitals involved in the covalent bond formation
should contain electrons with opposite spins. The two electrons then
move under the influence of both the nuclei.
iii.
The completely-filled orbitals (orbitals containing two paired
electrons) do not take part in the bond formation.
iv.
An s-orbital does not show any preference for direction. The nonspherical orbitals such as, p- and d-orbitals tend to form bonds in
the direction of the maximum overlap, i.e., along the orbital axis.
v.
Between the two orbitals of the same energy, the orbital which is
non-spherical (e.g., p- and d- orbitals forms stronger bonds than the
orbital which is spherically symmetrical, e.g., s-orbital.
vi.
The valence of an element is equal to the number of half-filled
orbitals present in it.
In the valence bond model, the stability of a molecule is explained in
terms of the following types of interactions.
a.
electron - nuclei attractive interactions, i.e., the electrons of one
atom are attracted by the nucleus of the other atom also.
b.
electron - electron repulsive interactions, i.e., electrons of one
atom are repelled by the electrons of the other atom.
c.
nucleus - nucleus repulsive interactions, i.e., nucleus of one atom
is repelled by the nucleus of the other atom.
Various interactions which act between the two atoms are shown in Fig.
The attractive and the repulsive interactions oppose each other. When
the attractive interactions are stronger than the repulsive
interactions, certain amount of energy is released. Due to the
lowering of energy the molecule becomes stable.
Valence Bond Description of Hydrogen Molecule
Chemical Bond and the Valence Bond Theory: The Orbital Overlap
In the beginning, atoms in a molecule were thought to be held by bonds. These
bonds are represented by drawing a small line (-) between the combining
atoms. Lewis described a chemical bond in terms of pair of electrons shared by
the combining atoms. So, the pair of electrons shared by the two nuclei may be
considered to be a chemical bond. Thus, the line drawn between the two atoms
to represent a chemical bond in the older concept, may be seen as a shared pair
of electrons in the Lewis concept.
Sharing of an electron pair is possible only when the atom are close enough to
overlap their orbitals. Thus, the orbital overlap is necessary for the sharing
of electrons and hence for the bond formation.
TYPES OF OVERLAPPING
Various types of atomic orbital overlap leading to the formation of covalent bond
are:
1. s - s overlap .
In this type of overlap, half-filled
s -orbitals of the two combining atoms overlap each other. This is shown in
Fig.
2. s - p overlap .
Here a half-filled s -orbital of one atom overlaps with one of the p orbitals having only one electron in it. This is shown in Fig.
3. p-p overlap along the orbital axis. This is called head on, end-on or endto-end linear overlap. Here, the overlap of the two half-filled p-orbitals takes
place along the line joining the two nuclei. This is shown in Fig.
4. p-p sideways overlap. This is also called lateral overlap. In this types
of overlap, two p-orbitals overlap each other along a line perpendicular
to the internuclear axis, i.e., the two overlapping p-orbitals are parallel
to each other. This is shown in Fig.
TYPES OF COVALENT BONDS: SIGMA (s) AND PI (p) BONDS
The overlapping of orbitals is possible in two ways.
(i) along their orbital axis so that the electron density along the axis is
maximum.
(ii) along a direction perpendicular to the bond axis due to sideways
overlap of the orbitals.
Depending upon the manner in which the two atomic orbitals overlap with
each other, two types of bonds are formed. These are called, sigma (s)
bond, and pi (p) bond.
A covalent bond formed due to the overlap of orbitals of the two atoms
along the line joining the two nuclei (orbital axis) is called sigma ( s )
bond. For example, the bond formed due to s-s and s-p, and p-p
overlap along the orbital axis are sigma bonds, (by convention Z-axis
is taken as inter-nuclear axis.
A covalent bond formed between the two atoms due to the sideways
overlap of their p -orbitals is called a pi ( p ) bond
Comparison of Sigma (s) and Pi (p) Bonds
Sigma (s) bond
1. It is formed due to axial overlap of the
twoorbitals. The overlap may be of s-s, s-p, pp orbitals.
Pi (p) bond
1. This bond is formed by the lateral (sideways)
overlap of two p-orbitals.
2. There can be only one sigma bond between
atoms.
3. The electron density is maximum and
cylindrically symmetrical about the bond
axis.
2. There can be more than one p-bond between
the two atoms.
3. The electron density is high along a direcion
tion at right angle to the bond axis.
4. The bonding is relatively strong.
4. The bonding due to a p-bond is weak.
5. Free rotation of atoms about sigma (s)
bond is possible.
5. Free rotation about a p bond is not possible.
6. It can be formed independently, i.e., there
can be a sigma (s) bond without having a p
bond in any molecule.
6. The p bond is formed only after s bond has
been formed.
Recapitulation Exercise
1. What is the basis of valence bond theory?
2. Identify two repulsive forces and two
attractive forces that influence the formation
of a chemical bond between two atoms.
3. What is meant by orbital overlap.
4. Show the formation of a bond by the
overlap ofs-s overlap. Mark the area where
the electron density is higher.
1.
Sub Topic 6:
Molecular Orbital Theory
Molecular orbital theory was put forward by R.S. Mulliken to explain the nature of
bonding in
the molecules of covalent compounds. Mulliken was awarded Nobel Prize for
Chemistry in 1966.
Major postulates of the theory are:
(i) The wavefunction of an electron in a molecule is called molecular orbital
(MO). The
molecular orbital surrounds all the nuclei in the molecule, i.e., MO’s are
polycentric.
(ii) The atomic orbitals (AO’s) of nearly equal energy, and appropriate symmetry
combine to
give equal number of MO’s. The MO’s are constructed by the linear combination
of the atomic
orbitals (LCAO method).
(iii) MO of lower energy is called bonding molecular orbital (  b ), while that of
higher energy
as antibonding molecular orbital ( a ),
(v) The electrons of the constituent atoms of a molecule are distributed over all the
available
MO’s in accordance with the Aufbau principle, the Pauli's exclusion principle and
Hund’s rule.
(vi) Like atomic orbitals (AO’s), the molecular orbitals can also be arranged
according to their energies. The internuclear axis is taken to be in the z-direction.
For the molecule or molecular ions formed from Li, Be, B, C, and N, the energies
of 2s and 2p orbitals are quite close to each other. Because of the repulsion
between the electrons that occupy 2s and 2p orbitals, the energy of the s2p
molecular orbital gets raised. Relative to p 2p orbitals.
Splitting patterns for the second row
Diatomic
If we combine the splitting
schemes for the 2s and 2p orbitals, we
can predict bond order in all of the
diatomic molecules and ions
composed of elements in the first
complete row of the periodic table.
Remember that only the valence
orbitals of the atoms need be
considered.
One minor complication that you should be aware of is that
the relative energies of the s and p bonding molecular
orbitals are reversed in some of the second-row diatomics.
The presence of one or more unpaired electrons accounts for the paramagnetic
nature of the molecule. The electronic configuration in which all the
electrons are paired indicate the diamagnetic nature of the species.
The strength of a chemical bond is described in terms of a parameter called bond
order.
As per definition, the bond order is expressed as,
Bond order = (No. of electrons in BMO-No. of electrons in ABMO)/2=(Nb-Na)/2
where, N b is the total number of electrons in bonding MOs.
N a is the total number of electrons in antibonding MOs.
(a)
When, N b > N a : Bond order > 0 (+ ve). Then, a stable bond formation is
indicated.
(b) When, N b =N a : Bond order =0. Then, the bond is unstable. In fact. such a
bond is not formed.
Conditions For the Formation of MOs From the Atomic Orbitals
Formation of MOs by the combination of atomic orbitals takes place only if the
following conditions are satisfied:
(i) The combining atomic orbitals should have nearly equal energies. Only the
atomic orbitals of nearly the same energy combine to form MOs. For
example, 1s atomic orbitals of two atoms can combine to form one bonding
(s 1s ) and one antibonding (s * 1s ) orbitals. The 1s atomic orbital of one
atom cannot combine with 2s or 2p atomic orbital of the other atom.
(ii) The combining atomic orbitals should have the same symmetry.
The atomic orbitals are oriented in space. Only those atomic orbitals
can combine to form molecular orbitals which have the same
symmetry about the molecular axis. For example, a p x orbital of an
atom can combine with a p x orbital of another atom. A p x orbital
cannot combine with a p z orbital.
(iii) The combining atomic orbitals should overlap effectively. MOs are
formed only if the combining atomic orbitals overlap to a reasonable
extent.
In-phase and out-of-phase wave combinations
“Matter waves” corresponding to the two separate hydrogen 1s orbitals
interact; both in-phase and out-of-phase combinations are possible, and
both occur. One of the resultants is the bonding orbital that we just
considered. The other, corresponding to out-of-phase combination of
the two orbitals, gives rise to a molecular orbital that has its greatest
electron probability in what is clearly the antibonding region of space.
This second orbital is therefore called an antibonding orbital.
Dicarbon
Dioxygen
Recapitulation Exercise
1. Write the molecular orbital configurations for O22–
and O22+ , and stating a reason, predict which of
them will be more stable.
2. Calculate the bond order of O2– molecular ion.
3. What is the bond order in molecular ion O2–
4. Calculate the bond orders for He2 and O2+ .
5. Arrange the following molecular species in
increasing order of stability (giving bond orders):
O2 , O2+ , O2 – , O 22 – .
Sub Topic 7:
Hybridisation
The concept of hydridisation is used to explain the nature of bonds, and
shape of the polyatomic molecules. For an isolated atom hybridisation
has no meaning.
According to the concept of hybridisation, certain atomic orbitals of
nearly the same energy undergo mixing to produce equal number of
new orbitals. The new orbitals so obtained are called hybrid orbitals.
The process of mixing of the atomic orbitals to form new hybrid orbitals is
called hybridisation.
All hybrid orbitals of a particular kind have equal energy, identical shapes
and are symmetrically oriented in space.
The types of atomic orbitals involved in hybridisation, and the nature of
hybridisation depends upon the requirements of the reaction. For
example, carbon in methane (CH4) shows sp3 hybridisation. In ethene
(ethylene, C2H4), it exhibits sp2 hybridisation. In ethyne (acetylene,
C2H2) it shows sp hybridisation.
The hybrid orbitals are designated according to the type and the number of
atomic orbitals merging together. For example,
Mixing orbitals
one s and three pone s and two pone s and one p-
Hybrid orbital
four sp3 orbitals
three sp2 orbitals
two sp orbitals
Hybridisation
sp3 hybridisation
sp2 hybridisation
sp hybridisation
Conditions Necessary For Hybridisation
Atomic orbitals undergo hybridisation only if the following conditions are
satisfied.
(i) Atomic orbitals of the same atom participate in hybridisation. Electrons
present in these atomic orbitals do not participate in the hybridisation, and
occupy the hybrid orbitals as usual.
(ii) The atomic orbitals participating in hybridisation should have nearly
equal energy.
(ii) Characteristics of Hybrid Orbitals
The characteristics of hybrid orbitals are:
i. The number of hybrid orbitals formed is equal to the number of the
atomic orbitals participating in hybridisation.
ii.
All hybrid orbitals are equivalent in shape, and energy, but different
from the participating atomic orbitals.
iii. A hybrid orbital which takes part in the bond formation must contain
only one electron in it.
iv.
A hybrid orbital, like atomic orbitals, cannot have more than two
electrons. The two electrons should have their spins paired.
v.
Due to the electronic repulsions between the hybrid orbitals, they tend
to remain at the maximum distance from each other
Types of Hybridisation
Depending upon the nature of the orbitals involved in hybridisation,
different types of hybridisation become possible. The type of
hybridisation shown by an atom depends upon the requirements of the
reaction.
•
sp3 (es pee three) hybridisation. In any atom, corresponding to energy
levels (or shells) for which n  2, there is one s orbital and three p
orbitals. For example, for n = 2, we have one 2s and three 2p orbitals;
for n = 3, we have one 3s, and three 3p orbitals. These four orbitals
undergo mixing to provide four new hybrid orbitals.
s
+
(px + py + pz)

one s-orbital
three p-orbitals
sp3
four hybrid orbitals
Hybridisation sp3 , geometry
tetrahedral, bond angle 109.5’
• sp2 (ess pee two) hybridisation. In certain reactions, one s and two p
(say px and py) orbitals of an atom undergo mixing to produce three
equivalent sp2 hybridised orbitals. The three sp2 hybrid orbitals are
oriented in a plane along the three corners of an equilateral triangle,
i.e., they are inclined to each other at an angle of 120. The third porbital (say pz here) remains unchanged. Each hybrid orbital has 33.3%
s-character and 66.7% p-character. Formation of sp2-hybrid orbitals
from one s and two p-orbitals is shown in Fig. 6.43.
s + (px + py)  sp2
one s-orbital two p-orbitals three hybrid orbitals
Boron trifluoride has a plane trigonal shape
in which all three bonds are identical.
Hybridisation sp2 , geometry trigonal planar, bond angle 120’
• sp (ess-pee) hybridisation. In this type of hybridisation, one s and one
p (say pz) orbitals belonging to the same main energy level hybridise to
give two sp hybrid orbitals. These sp hybrid orbitals are oriented at an
angle of 180 to each other. Each hybrid orbital has 50% s- and 50% pcharacter. The other two p-orbitals (say 2px and 2py) remain
unhybridised and are oriented at right angles to each other and to the
internuclear axis.
Hybridisation sp , geometry Linear, bond angle 180’
Formation of ethene molecule.
Formation of ethyne (acetylene) molecule.
Recapitulation Exercise
1. Give one example each of sp, sp2 & sp3
Hybridisation.
2. Arrange the following in the decreasing order of
their bond angles:
NH3, CO2, CH4, C2H6, , C2H2
3. Describe all the conditions necessary for
Hybridisation.
Sub Topic 8:
Hydrogen Bonding
If a hydrogen atom is bonded to a highly electronegative element such as
fluorine, oxygen, nitrogen, then the shared pair of electrons lies more
towards the electronegative element. This leads to a polarity in the
bond in such a way that a slight positive charge gets developed on Hatom, viz.,
H+ d : O- d H+ d : F- d H+ d : N- d
This positive charge on hydrogen can exert electrostatic attraction on the
negatively charged electronegative atom of the same or the other
molecule forming a bridge-like structure such as
Xd - - Hd+ × × × × × × Yd- - Hd+
where X and Y are the atoms of strongly electronegative elements.
The bond between the hydrogen atom of one molecule and a more
electronegative atom of the same or another molecule is called
hydrogen bond.
A hydrogen bond is shown by a dotted line (.....).
Hydrogen bond is not a covalent bond as the 1s orbital of hydrogen is
already completed, and the 2s level is high up in its energy.
Conditions Necessary For the Formation of Hydrogen Bond
Hydrogen bond is formed only when the following conditions are
satisfied. (i) Only the molecules in which hydrogen atoms is linked to
an atom of highly electro-negative element, are capable of forming
hydrogen bonds.
(ii) The atom of the highly electronegative element should be small.
These conditions are met by fluorine, oxygen and nitrogen atoms. As a
results, all compounds containing hydrogen atom linked to an atom
of either N, O, or F exhibit hydrogen bonding.
Some Typical Compounds Showing Hydrogen Bonding
Hydrogen fluoride (HF).
Water (H2O).
Ice (H2O(s)).
Each Oxygen is "linked" in by a combination of a covalent bond and a
hydrogen bond to 4 other Oxygens.
Notice that each Oxygen can be linked to Hydrogen in one of two ways.
Or
Types of Hydrogen Bonding
There are two types of hydrogen bonding, viz.,
(a) Intermolecular hydrogen bonding (b) Intermolecular hydrogen
bonding
When the hydrogen bonding is between the H-atom of one molecule
and an atom of the electronegative element of another molecule, it is
termed as intermolecular hydrogen bonding. For example, hydrogen
bonding in water, ammonia etc., is intermolecular hydrogen bonding.
The intramolecular hydrogen bonding is between the hydrogen of one
functional group, and the electronegative atom of the adjacent
functional group in the same molecule. For example, the molecule of
o-nitrophenol, shows intramolecular hydrogen bonding. The pnitrophenol shows intermolecular hydrogen bonding.
Recapitulation Exercise
1. What is a hydrogen bond?
2. Which type of compounds exhibit hydrogen
bonding?
3. By giving suitable examples differentiate
between intermolecular & intramolecular
hydrogen bonding.
SUMMARY
• Here's a quick summary of how to predict the type of
bond likely to be found in any given substance:
1. Two nonmetal atoms usually form covalent bonds.
A metal and a nonmetal atom usually form ionic
bonds.
Why? Great question! It happens by the numbers —
in this case, electronegativity values. The covalent
bonds between nonmetal atoms have relatively low
differences in electronegativity values. Looking at the
periodic table (again!) reveals that electronegativity
values for the nonmetals generally range from 2.0 to
4.0 (with a few exceptions). Because the
electronegativity values are fairly similar for all
nonmetal atoms, they should have low differences in
electronegativity values. For example, a carbonoxygen bond has an electronegativity difference of
1.0. This value is considered rather low.
• Guess what happens with ionic bonds because they usually
occur between a metal and a nonmetal atom? The periodic table
shows that metal atoms generally have electronegativity values
that range from 0.7 to 1.5. The difference in electronegativity
between metal and nonmetal atoms in an ionic bond is therefore
relatively high compared to the difference in covalent bonds.
• For example, a common ionic solid is table salt (NaCl). The
difference in electronegativity values for sodium and chlorine is
2.1. This value for the ionic bond between sodium and chlorine
is two times greater than the value of 1.0 for the covalent
carbon-oxygen bond.
Try It Out!
• Identify the following bonds as covalent,
ionic, or polar covalent:
H-O
K-Cl
C-C
H-H
References
1.
2.
3.
4.
5.
6.
Physical Chemistry by Puri, Sharma & Pathania.
Inorganic Chemistry by J.D.Lee.
Modern abc of Chemistry by S.P. Johar.
New Course Chemistry by Pradeep Publications.
Organic Chemistry by Morrison & Void.
Internet