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

3

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.

4

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 (ii) Characteristic properties of ionic compounds

(iii) Formation of cation and anion (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 = 89.2 (ii) Ionisation energy of K = 418.8

(iii) Dissociation energy of F2 = 158.8 (iv) Electron affinity of F = - 328

(v) Lattice energy of KF = - 821

(vi) Electronegativity of K = 0.82 ( 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 (ii) Multiple bonds

(iii) Polar and non-polar covalent bonds

(iv) Coordinate bond (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 ( 19.5

0C) and ammonia ( - 33

0C ) 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 Cl

-

http://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/chemical-

bonding/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