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Chemical Bonding
Concepts
(i) Formation of a chemical bond
(ii) Nature of a chemical bond
(iii) Lewis theory
(iv) Types of chemical bond
Introduction
Though the periodic table has a place for 118 elements, there are obviously more
substances in nature than 118 pure elements. This is because atoms of elements can react
with one another to form new substances called compounds. When two or more elements
combine, the resulting compound is unique both chemically and physically from its
parent atoms. For example, sodium is a silver coloured metal that reacts so violently with
water that flames are produced when sodium gets wet. The element chlorine is greenish
coloured gas that is so poisonous that it was used as a weapon in world war I. When
chemically bonded together, these two dangerous substances form the compound sodium
chloride,
a compound so safe that we eat it every day – common table salt.
Formation of a chemical bond
Free atoms of elements are in random motion and possess some energy. Farther
the atoms are, greater is their energy and lesser is the stability. Two or more atoms unite
to form a molecule because in doing so, the energy of the united atoms is lowered. Thus
the ‘molecule’ becomes stable in comparison to separate atoms. In other words, a stable
chemical union called ‘bond’ between two or more atoms comes into existence only if the
energy is lowered when the atoms come in close vicinity. The lower the energy of the
molecule, the stronger the bond and more is the stability to the bonded atoms.
Nature of chemical bond
A chemical bond is an attraction between atoms. It is the attraction caused by the
electromagnetic force between opposing charges either between electrons and nuclei or
as the result of a dipole attraction. Since opposite charges attract via a simple
electromagnetic force, the negatively charged electrons revolving round the nucleus and
the positively charged protons in the nucleus attract each other. Also an electron
positioned between two nuclei will be attracted to both of them. Thus, the most stable
configuration of nuclei and electrons is one in which the electrons spend more time
between nuclei than anywhere else in space. These electrons cause the nuclei to be
attracted to each other and this attraction results in the bond. Electrons occupy large
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volume compared to the nuclei and this volume keeps the atomic nuclei relatively far
apart as compared with the size of the nuclei themselves.
The force of attraction which holds the two atoms together in a molecule
is called a chemical bond.
Lewis theory
In 1916, an American chemist, Lewis proposed that chemical bonds are formed
between atoms because electrons from the atoms interact with each other. Lewis had
observed that many elements are most stable when they contain eight electrons in their
outermost or valence shell of the atom. He suggested that atoms with fewer than eight
electrons bond together to share electrons and complete their valence shell.
While some of Lewis predictions have since been proven incorrect ( he suggested
that electrons occupy cube – shaped orbitals ), his work established the basis of what is
known today about chemical bonding.
Essentials of Lewis theory
Between 1916 and 1919, Lewis, Kossel and Langmuir made several important
proposals on bonding which lead to the development of Lewis theory of bonding.
1) Valence electrons mainly play a fundamental role in bonding.
2) Ionic bonding involves the transfer of one or more electrons from one atom to another.
3) Covalent bonding involves sharing of electrons between atoms.
4) Electrons are transferred or shared between atoms such that each atom achieves the
electron configuration of a noble gas i.e. having eight electrons in the outermost shell
called octet.
5) This arrangement is called octet rule. ( Exception – He)
6) Exceptions to octet rule may occur.
Lewis proposed symbols which represent the resulting structures that follow the
octet rule. In a Lewis symbol, an element is surrounded by up to 8 dots where elemental
symbol represents the nucleus and the dots represents the valence electrons.
Activity 1 - Draw the Lewis dot formula for following molecules - BF3, KCl
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Types of chemical bonds
Following figure shows a road map of chemical bonding i.e. which elements will form
which type of bond
Figure 1 – Periodic table and elements forming different types of bonds.
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Chemical bonds can be divided into three major types : ionic bonds which occur between
a metal and a non-metal; covalent bonds which occur between two non-metals; and
metallic bonds which occur within metals. Some people consider hydrogen bond as a
separate type of bond. In an ionic bond, one or more electrons are transferred from metal
to non-metal and the resultant ions are attracted to each other by coulombic forces. In a
covalent bond, non-metals share electrons that interact with the nuclei of both atoms via
coulombic forces, holding the atoms together. In a metallic bond, the atoms form a lattice
in which each metal atom loses electrons to an ‘ electron sea’. The attraction of the
positively charged metal ions to the electron - sea holds the metal atoms together.
Hydrogen bond occurs in some restricted hydrides. In addition, there are dipole – dipole
interactions and van – der – Waals forces which are small in magnitude and play a role
in bonding limited substances.
.
Activity 2 – Select one element from left hand side, one element from the right hand side
and one element from the middle of the periodic table. Predict how many types of bonds
each element can form with its own atoms as well as other atoms.
Check your understanding
(i) Why do atoms tend to combine and form a bond ?
(ii) When atoms come close, which forces come into existence ?
(iii) What is Lewis theory of bond formation ?
(iv) How many main types of bonds are known ?
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Concepts
(i) Formation of ionic bond
(iii) Formation of cation and anion
(ii) Characteristic properties of ionic compounds
(iv) Difference between atoms and ions
Formation of ionic bond
An ionic bond ( also called as electrovalent bond ) is a type of chemical bond that
involves a metal ion and a non-metal ion ( or polyatomic ions such as ammonium )
through electrostatic attraction. In short, it is a bond formed by the attraction between two
oppositely charged ions.
The metal donates one or more electrons, forming a positively charged ion or
cation with a stable electron configuration. These electrons then enter the non-metal,
causing it to form a negatively charged ion or anion which also has a stable electron
configuration. The electrostatic attraction between the oppositely charged ions causes
them to come together and form a bond.
For example, when sodium ( Na ) and chlorine ( Cl) are combined, the sodium
atoms each lose an electron, forming a cation (Na+) and the chlorine atoms each gain an
electron to form an anion (Cl-). These ions then are attracted to each other in 1:1
proportion to form sodium chloride NaCl.
Na + Cl →
Na+ + Cl-
→
NaCl
→
Figure 2 - Combination of Na and Cl to form Na+ and Cl-
The electrostatic force of attraction between two oppositely charged ions formed by
transfer of electrons from one atom to another is called an ionic or electrovalent bond.
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The figure given below shows the variation of potential energy as a function of distance
of separation between sodium or chlorine atoms. An atom of sodium has one electron
extra outside the closed shell and it takes 5.14 electron volts of energy to remove that
electron (its ionization potential is 5.14 eV).
( diagram is not to the scale)
Figure 3 – P.E. diagram for NaCl molecule
The chlorine atom is short of one electron to fill a shell and it releases 3.62 electron volts
when it acquires that electron ( its electron affinity is 3.62 eV). This means that it takes
only 1.52 eV( 5.14 – 3.62 ) of energy to donate one of the sodium electrons to chlorine
when they are far apart. When the resultant ions are brought close together, their electric
potential becomes more and more negative, reaching – 1.52 eV at about 0.94 nm
separation. This means that if neutral sodium and chlorine atoms found themselves closer
than 0.94 nm, it would be energetically favourable to transfer electron from Na to Cl and
form the ionic bond.
The potential energy curve shows that there is a minimum at 0.236 nm separation
and then a steep rise in potential which represents a repulsive force. This repulsive force
is more than just an electrostatic repulsion between the electron clouds of the two atoms.
The removal of electron from the atom is endothermic and causes the ions to have a
higher energy. There may also be energy changes associated with breaking of existing
bonds or the addition of more than one electron to form anions. However, the attraction
of the ions to each other lowers their energy.
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The energy balance cycle for NaCl is shown below.
(i) Gaseous sodium atom is formed from solid sodium metal
Na (s) + 108 kJ mol-1 → Na(g)
(ii) Sodium ion is formed from gaseous sodium atom.
Na (g) + 496 kJ mol-1 → Na+ (g) + e(iii) Chlorine molecule dissociates into gaseous chlorine atoms.
½ Cl2 (g) + 121 kJ mol-1 → Cl (g)
(iv) Chloride ion is formed from gaseous chlorine atom.
Cl (g) + e- → Cl- (g) + 349 kJ mol-1
(v) Sodium ions and chloride ions interact to form solid sodium chloride.
Na+ (g) + Cl- (g) → Na+ Cl- (s) + 787 kJ mol-1
Energy evolved = 349 + 787
= 1136 kJ
Energy absorbed = 108 + 496 + 121 = 725 kJ
----------------------------------------------------------Energy evolved = 411 kJ mol-1
Ionic bonding will occur only if the overall energy change for the reaction is
favourable – when the bonded atoms have a lower energy than the free ones. The larger
the resulting energy change, the stronger the bond. The low electronegativity of the
metals and high electronegativity of non-metals means that the energy change of the
reaction is most favourable when metals lose electrons and non-metals gain electrons.
Notice that when sodium loses its one valence electron, it gets smaller in size, while
chlorine grows larger when it gains an additional valence electron. This is typical of the
relative sizes of the ions to atoms. Positive ions tend to be smaller than the parent atoms
while negative ions tend to be larger than their parent. After the reaction takes place, the
charged Na+ and Cl- ions are held together by electrostatic forces, thus forming an ionic
bond.
Activity 3 - You are given following data – All energy values are in kJ mol-1
(i) Sublimation energy of K =
(iii) Dissociation energy of F2 =
(v) Lattice energy of KF
=
(vi) Electronegativity of K
=
89.2
158.8
- 821
0.82
(ii) Ionisation energy of K = 418.8
(iv) Electron affinity of F = - 328
( vii) Electronegativity of F = 4
What type of bond K and F will form and energetically will KF be stable ?
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Activity 4 – From the crystal structure of sodium chloride given in the book, find out the
coordination number of Na+ ion and Cl- ion and try to draw yourself the structure of
NaCl.
Other examples of ionic bonding
As stated earlier, more the difference in electronegativity of the two atoms, more are the
chances of forming ionic bonds. For example, two potassium atoms can lose one electron
each to oxygen atom and potassium and oxygen may combine to form ionic bond.
Similarly, rubidium and fluorine atom, magnesium and chlorine atom, calcium and
oxygen atom can form ionic bond.
Characteristic Properties of Ionic Compounds
Ionic compounds have following characteristic properties.
1) Ionic compounds involve ionic bonds which are formed between metals and
non-metals.
2) In naming simple ionic compounds, the metal is always first, the non-metal second
( e.g. sodium chloride )
3) Ionic compounds dissolve easily in water and other polar solvents.
4) In solution and in molten state ionic compounds easily conduct electricity.
5) Ionic compounds tend to form crystalline solids with high melting temperatures.
Pure ionic bonding is not known to exist. All ionic compounds have a degree of covalent
bonding. The larger the difference in electronegativity between two atoms, the more ionic
the bond.
Formation of cation and anion
(i) When an atom loses electron, it gets an overall positive charge because the number of
protons now exceed the number of electrons. The positively charged ion is called a
cation. The process of formation of a cation from its atom is called oxidation.
(ii) When an atom gains electron, it gets an overall negative charge because the number
of electrons now exceed the number of protons. The negatively charged ion is called
an anion. The process of formation of an anion from its atom is called reduction.
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Difference between atoms and ions
Atoms
Ions
(i) Atoms are electrically neutral because
(i) Ions are charged particles because of
protons and electrons are equal in number.
imbalance of protons and electrons.
(ii) The outermost shell may or may not
(ii) The outermost shell has a completed
have a completed duplet or octet e.g
duplet or octet e.g.
Ne = 2,8 ( completed octet )
Cl- = 2,8 ( completed octet )
Na = 2,8,1 ( incomplete octet )
Li+ = 2 ( completed duplet )
(iii) Atoms may be or may not be capable of (iii) Ions are capable of independent
free existence e.g.
existence in solution or gaseous state
He atom exists in uncombined state
e.g. NaCl → Na+ + Cl- ( in solution)
Hydrogen ( H2) exists in combined
Na → Na+ + e(Gaseous)
state
Check your understanding
(i) Which elements in the periodic table tend to form ionic bond ?
(ii) In terms of electronegativity, what is the condition for formation of an ionic bond?
(iii) What is the criterion to know whether the ionic compound will be stable or not?
(iv) Which pair of elements in the periodic table will form the strongest ionic bond ?
(v) Why is it that the process of formation of cation is called oxidation and formation of
anion is called reduction ?
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Concepts
(i) Formation of covalent bond
(iii) Polar and non-polar covalent bonds
(iv) Coordinate bond
(ii) Multiple bonds
(v) Characteristics of covalent compounds
Formation of covalent bond
The second major type of chemical bond occurs when atoms share electrons. As
opposed to ionic bonding in which a complete transfer of electrons occurs, covalent
bonding occurs when two ( or more ) elements share electrons. Covalent bonding occurs
because the atoms in the molecule have a similar tendency for electrons ( generally to
gain electrons.) This most commonly occurs when two non-metals bond together.
Because both of the non-metals want to gain electrons , the elements involved will share
electrons in an effort to fill their valence shells. A good example of a covalent bond is
that which occurs between two hydrogen atoms. Atoms of hydrogen (H) have one
valence electron in their electron shell. Since the capacity of this shell is two electrons,
each hydrogen atom will ‘want’ to pick up a second electron. In an effort to pick up a
second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the
molecule H2. Since the hydrogen molecule is a combination of equally matched atoms,
the atoms will share each other’s single electron, forming one covalent bond. In this way,
both atoms share the stability of a full valence shell.
A chemical bond formed by sharing of electrons between atoms is
called a covalent bond.
As the two hydrogen atoms approach one another, in addition to nucleus – electron
attraction, nuclear-nuclear repulsion and electron – electron repulsion also come into
existence. When the two hydrogen atoms are at a distance of 0.074 nm, the potential
energy of the two hydrogen atoms together is at its minimum and releases 4.52 eV. At
this stage, a chemical bond is formed. If the hydrogen atoms come still closer, the
potential energy rises steeply making the molecule unstable. Thus, the sharing of
electrons is energetically favourable to both the hydrogen atoms with the formation of
stable single covalent bond.
The figure given below shows the variation of potential energy as a function of distance
of separation of hydrogen atoms.
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( diagram is not to the scale)
Figure 4 – P.E. diagram for H2 molecule
Following figure shows the formation of single covalent bond between two hydrogen
atoms and two chlorine atoms.
Figure 5 - Bonding in H2 and Cl2 molecule
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Two or more atoms of different elements can also share electrons to form a single bond
between them and complete the octet ( or duplet ) of each atom. For example, in methane,
one carbon and four hydrogen atoms share one electron pair each to form four C - H
bonds, in ammonia, one nitrogen and three hydrogen atoms share one electron pair each
to form three N – H bonds and in water, one oxygen and two hydrogen atoms share one
electron pair each to form two O – H bonds. This is shown in the following diagram.
Figure 6 – Bonding in H2O, NH3 and CH4 molecules
Multiple bonds
For every pair of electrons shared between two atoms, a single covalent bond is
formed. Some atoms can share two or three pairs of electrons forming multiple bonds i. e.
a double or triple bonds. For example, oxygen atom has six electrons in its outermost
shell. It needs two electrons to complete its octet and attains the configuration of neon.
Hence two oxygen atoms combine by sharing two pairs of electrons between them and
form a double bond. Similarly, nitrogen atom has five electrons in its outermost shell. It
needs three electrons to complete its octet and attain the configuration of the inert gas
neon. Hence, two nitrogen atoms combine by sharing three pairs of electrons between
them and form a triple bond. In HCN molecule, H and C atoms share one pair of electron
to form a single bond while C and N atoms share three pairs of electrons to form a triple
bond.
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Following figure shows the multiple bonds in O2, N2 and HCN molecules.
Figure 7 - Bonding in O2, N2 and HCN molecules.
Activity 5 - Carbon atom has four electrons in its outermost shell. Oxygen atom has six
electrons in its outermost shell. Arrange the valence electrons around these two atoms
and draw the Lewis dot formula in such a way that each atom completes its octet. Name
and count the types of bonds in the molecule.
Polar and Non-polar covalent bonds
There are two subtypes of covalent bonds – non-polar and polar. The H2 molecule
is a good example of the first subtype of covalent bond. Since both atoms in H 2 molecule
have an equal attraction ( or affinity ) for electrons, the bonding electrons are equally
shared between the two atoms i.e. the shared pair lies exactly in the middle of two atoms
and a non-polar covalent bond is formed. There is no charge separation and the molecule
is non-polar. Whenever two atoms of the same element bond together, a non-polar
covalent bond is formed. Following figure shows the non-polar covalent bond between
H2 and O2 molecules.
Figure 8 - Non-polar covalent bonds in H2 and O2 molecules
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A polar covalent bond is formed when electrons are unequally shared between
two atoms. Polar covalent bonding occurs because one atom has stronger affinity for
electrons than the other ( yet not enough to pull the electrons away completely and form
an ion).In a polar covalent bond, the bonding electrons spend more time around the atom
that has the stronger affinity for electrons. Due to this uneven distribution of charge, one
end of the molecule acquires a slightly positive charge while the other end acquires a
slightly negative charge. These slight charges are represented by the symbols ∂ + and ∂ (called delta). Good examples of polar covalent bond are HCl and H2O. The figure given
below shows the polar covalent bond in HCl and H2O molecule.
Figure 9 - Polar covalent bond in HCl and H2O molecules.
The polar or non-polar nature of the covalent bond can be predicted from the
electronegativity values of the two atoms. There is a correlation between the
electronegativity difference and the percentage ionic character of the molecule. In case of
HCl, the electronegativity difference between H and Cl is 0.9 and the ionic character is
20%. In case of NaCl molecule, the elctronegativity difference between Na and Cl is 2.1
and the ionic character is 65%. In order to have 50% ionic character in a molecule, the
atoms should have 1.7 as the difference in electronegativity values.
It is also possible that the multi-bond molecule is non-polar but the individual bonds in
the molecule are polar. This is the case in carbon tetrachloride molecule. Each C – Cl
bond is slightly polar but the overall molecule is non-polar. When the directions of the
bonds are taken into account, the net effect of the polarity of four C-Cl bonds is zero.
Following figure shows the individual polarities of bond in carbon tetrachloride
molecule.
Figure 10 - Non-polar carbon tetrachloride molecule
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Activity 6 - Draw the Lewis dot formula and show the polar covalent bond formation in
HBr molecule.
Coordinate bond
A coordinate bond, also known as dative or semi polar bond, is a special type of
covalent bond in which the shared pair of electrons comes from one of the bonding atoms
only. This bond is formed when an electron pair donor ( Lewis base ) donates a pair of
electrons to an electron pair acceptor ( Lewis acid ) to give a so called adduct. The
process of forming a coordinate bond is called coordination. In this process, the electron
donor acquires a formal positive charge while the electron acceptor acquires a formal
negative charge. Since a dipole is created, this bond is, sometimes, called as a dipolar
bond. The distinction between a normal covalent bond and a coordinate bond is artificial.
Once the coordinate bond is formed, its strength and description is no different from that
of other polar covalent bond. Any atom, ion or molecule which has a lone pair of
electrons is capable of forming a coordinate bond. For example, ammonia molecule has a
lone pair of electrons. It can act as electron donor ( Lewis base). Hydrogen ion is electron
deficient and can act as an electron acceptor ( Lewis acid ). When they come together,
they form a coordinate bond. In this process, nitrogen of the ammonia molecule acquires
a formal positive charge while hydrogen ion acquires a formal negative charge. Once the
coordinate bond is formed all four N – H bonds in ammonium ion become identical in all
respects. The figure given below shows the formation of a coordinate bond between
ammonia molecule and H+ ion.
Figure 11 - Formation of coordinate bond.
Formation of H3O+ ion and NH3 → BF3 adduct are some more examples of coordinate
bonding.
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Characteristic Properties of Covalent Compounds
Covalent compounds have following characteristic properties
1) Covalent compounds do not exist as ions but exist as molecules. They may occur in
solid, liquid or gaseous state.
2) They are generally soft and have low melting and boiling points.
3) Covalent compounds are generally insoluble or less soluble in water and in other polar
solvents.
4) Covalent compounds are poor conductors of electricity in fused or dissolved state.
Check your understanding
(i) Draw a potential energy curve for H2 molecule and show the bond length and potential
energy at which H2 molecule is formed.
(ii) What is the difference between covalent bond and coordinate bond?
(iii) Choose the pairs of atoms which will form (i) non-polar (ii) polar covalent bond.
Be, B, C, N, O, F, N, O, F
(iv) Identify the types of bonds in NH4Cl molecule.
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Concepts
(i) Metallic bonding
(ii) Characteristic properties of metals
Metallic bonding
The elements which are placed on the extreme left, the middle and a few on the
right of the periodic table are metals. Alkali metals like sodium, potassium, alkaline earth
metals like magnesium, calcium, transition metals like iron, cobalt, nickel, copper and
others like lead, tin represent the family of metals. They have low electronegativity.
They tend to lose their valence electrons easily. When we have a macroscopic collection
of metal atoms, the valence electrons are detached from the atoms but not held by any of
the other atoms. In other words, these valence electrons are free from any particular atom
and are held only collectively by the entire assembly of atoms. When atoms lose their
outer-shell electrons they become positive ions. The outer electrons become a ‘sea’ of
mobile electrons surrounding a lattice of positive ions. The positive ion cores are held
more or less at fixed places in an ordered or crystal lattice. The valence electrons are free
to move about under applied stimulation like electrical field or heat. This is called
‘electron sea model’ of metals
.
The force of attraction which holds the delocalized (or mobile) electrons and the metallic
nuclei together in a metal is called a metallic bond.
Following figure shows electron sea model of metals.
Figure 12 - Electron sea model of a metal
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Although the term ‘metallic bond’ is often used in contrast to the term ‘covalent
bond’ it is preferable to use the term ‘metallic bonding’ because this type of bonding is
collective in nature and a single ‘metallic bond’ does not exist.
Characteristic Properties of Metals
Metals show following characteristic physical properties:
1) At room temperature, they are solids (except mercury)
2) They are opaque to light.
3) They, generally, have high density.
4) They show metallic luster.
5) They are malleable and ductile in their solid state.
6) They are good conductors of heat and electricity.
7) They have crystal structure in which each atom is surrounded by eight to twelve near
neighbours.
Activity 7 – Draw the picture of a metal lattice and show the position of metal nuclei and
valence electrons in the lattice.
Activity 8 – ‘Metals generally have high densities.’ Support this statement by giving
densities of some metals.
Check your understanding
(i) Why the crystal structure of metal is described as a sea of electrons?
(ii) Give any one property of metals which can be explained by its crystal structure.
Justify your answer.
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Concepts
(i) Hydrogen bond
(ii) Effects of hydrogen bonding
Hydrogen bond
This is a different type of bond. It is restricted to only some molecules containing
hydrogen atoms.
The force of attraction between the hydrogen atom attached to an electronegative atom of
one molecule and an electronegative atom of another molecule is called hydrogen bond.
Usually, the electronegative atom is O, N or F. In a molecule, the O, N or F atom has a
partial negative charge and then the hydrogen atom which has a very small size has a
partial positive charge. This type of bond always involves hydrogen atom and hence the
name hydrogen bond.
In order to form a hydrogen bond, it is necessary that the electronegative atom
should have one or more lone pairs of electrons and a partial negative charge so that there
is a force of attraction termed as dipole-dipole interaction. The hydrogen atom which has
a partial positive charge tries to find another atom of O,N or F with excess of electrons to
share and is attracted to partial negative charge. This forms the basis of hydrogen bond.
The hydrogen bond can occur between molecules ( intermolecular ) like HF or
within different parts of a single molecule ( intramolecular ) like o-nitro phenol. The
hydrogen bond is stronger than van-der-Waals’ bond but weaker than covalent or ionic
bond. The hydrogen bond has the bond energy in the range 5 to 30 kJ per mole.
Following figure shows hydrogen bonding in HF molecules.
Figure 13 - Hydrogen bonding in HF
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Activity 1 – Draw the structure of water molecules with hydrogen bonding.
Effects of hydrogen bonding
Hydrogen bonding has effects on the properties of certain substances.
(i) Hydrogen bonding leads to association of molecules which affects the physical state
of a substance. For example, HF which should be a gas at room temperature,
becomes a liquid due to association of molecules.
(ii) Covalent compounds are normally insoluble in water. But compounds like ethanol,
lower aldehydes, ketones, though covalent, are soluble in water due to formation of
hydrogen bonds with water molecules.
(iii) The boiling points of water ( 1000C), HF ( 19.50C) and ammonia ( - 330C ) are
exceptionally high as compared to other Group 16 hydrides which have no hydrogen
bonds.
(iv) Intramolecular hydrogen bonding is partly responsible for secondary, tertiary and
quaternary structure of proteins and nucleic acids. It also plays an important role in
the structure of polymers.
Activity 2 - Draw the structure of o – nitro phenol and show the intramolecular bonding
in it.
Check your understanding
(i) Hydrogen bonding is known only in the hydrides of O, N and F. Why?
(ii)Water molecules are joined by hydrogen bonds. Is hydrogen bonding present in ice
also?
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References / Figures / Diagrams etc
1) Fig. 1 – Periodic table and elements forming different types of bonds
www. Smallscalechemistry.colostate.edu/…/chemicalbonding.pdf
2) Fig. 2 – Combination of Na and Cl to form Na+ and Clhttp://www.visionlearning.com/library/module_viewer.php?mid=55
3) Fig. 3 – Potential energy diagram for NaCl molecule
http://hyperphysics.phy-astr.gsu.edu/ hbase/chemical/bond.html
4) Fig. 4 – Potential energy diagram for H2 molecule
http://hyperphysics.phy-astr.gsu.edu/ hbase/molecule/hmol.html
5) Fig. 5 – Bonding in H2 and Cl2 molecule
Http://www.tutorvista.com/content/chemistry/chemistry-ii/chemicalbonding/covalent-bonding.php
6) Fig. 6 – Bonding in H2O, NH3 and CH4 molecules
http://hyperphysics.phy-astr.gsu.edu/hbase/chemical/lewis
7) Fig.7 - Bonding in O2, N2 and HCN
www.tutornext.com/covalent.bond-continued/2255
8) Fig. 8 – Non-polar covalent bond in H2 and O2 molecules
For H2 – http://www.tutorvista.com/topic/non-polar-covalentbond
For O2 - www.tutornext.com/covalent.bond-continued/2255
9) Fig. 9 – Polar covalent bond in HCl and H2O molecules
For H2O - www.ausetute.com.au/molpolar.html
For HCl – http://users.stlcc.edu/gkrishnan/polar.html
10) Fig. 10 – Non-polar carbon tetrachloride molecule
http://www.chemguide.co.uk/atoms/bonding/electroneg.html
11) Fig. 11 - Formation of coordinate bond
www.tutorvista.com/content/chemistry-iii/chemical-bonding/dative-bond.php
12) Fig. 12 – Electron sea model of a metal
www.chemguide.co.uk/atoms/bonding/metallic.html
13) Fig. 13 – Hydrogen bonding in HF
http://en.wikipedia.org/wiki/hydrogen fluoride