3. CHEMICAL BONDING - TopperLearning

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3. CHEMICAL BONDING

1. Chemical Bond

Most of the substances are found in nature in the form of clusters or aggregates of atoms. Any such aggregation, in which atoms are held together and which is electrically neutral is called a molecule. �e molecules are made of two or more atoms joined together by some force acting between them. �e force is termed as a chemical bond. �us, a chemical bond is de�ned as a force that acts between two or more atoms to hold them together as a stable molecule.

All is a game

of molecular interactions.

So let’s dive

into this lesson

to learn about

it.

3.2 Chemical Bonding

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Kossel and Lewis approach to chemical bonding: First of all in 1916, W. Kossel and G. N. Lewis independently proposed the electronic theory of valency. Kossel explained the bonding in ionic substances while Lewis explained bonding in other substances. �e basis of the theory was the stable electronic con�guration of noble gases. Later on, the idea of wave nature of electron was incorporated to explain the chemical bonding in the light of quantum theory proposed by other scientists like Linus Pauling, Heitler, London and R.S. Mullikan.

1.1 Electronic Theory of Chemical Bonding

On the basis of the electronic con�guration of the atoms of elements and their chemical behavior, the electronic theory of bonding may be stated as follows:

Atoms of various elements combine with one another by redistribution of valence shell electrons so as to acquire stable electronic con�guration in their outermost shell like their nearest noble gas.

Octet rule: �e electronic con�guration of noble gases like Ne, Ar, Kr, etc. show that the atom of each noble gas contains eight electrons in its outermost shell. Since each atom of noble gas is stable, G. N. Lewis formulated the rule that any atom with eight electrons in its outermost shell must be stable. �is is popularly called as octet rule, and is very useful to learn about the chemical bonding. According to the octet rule:

�e atoms of elements combine to have eight electrons in their outermost shell.

In this type of bond, valence electrons are transferred from one atom to another. One atom donates its excess electrons to another atom so that both the atoms may acquire a stable noble gas con�guration. �e atom which loses electron becomes positively charged and is called the cation. �e atom which takes up the electron lost by the �rst atom becomes negatively charged and is called the anion. �ese two oppositely charged ions are now held together by an electrostatic force of attraction. �is force of attraction binding the two atoms together is known as an electrovalent or ionic bond.

�us, the chemical bond formed between two atoms by the transfer of one or more valence electrons from one atom to the other is known as an electrovalent or ionic bond. It is also called a polar bond.

Example: Combination of sodium (Na) and chlorine (Cl) atoms to form sodium chloride (NaCl)

�e atomic number of sodium is 11. So its electronic con�guration is 2, 8, 1. It has only one electron in its outermost shell. �e Na atom transfer this electron to chlorine (Cl) atom and becomes positively charged sodium ion (Na+).

1e

2,8,1 2,8Na [Na]

�us, the electronic con�guration of the Na+ ion is the same as that of neon which is the noble gas nearest to sodium in the periodic table. Let us consider the chlorine atom (Cl). �e atomic number of chlorine is 17. So its electronic con�guration is 2, 8, 7. It has 7 electrons in the outermost shell. It lacks one

3.3 Foundation for Chemistry

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electron to acquire a stable noble gas con�guration. So a chlorine atom takes one electron transferred by the sodium atom and becomes negatively charged chlorine ion (Cl–).

2,8,7 (2,8,8)Cl 1e [Cl ]

�us, the chloride ion (Cl-) attains the con�guration of the nearest noble gas, argon. �e two ions (Na and Cl ), being oppositely charged, are now held together by an electrostatic force of attraction as Na Cl .

2,8,1 2,8,7 2,8 2,8,8Na Cl Na Cl

�e formation of sodium chloride can be shown diagrammatically as in given �gure.

�e force that holds Na+ and Cl- ions together is called an electrovalent bond. As this bond exists between ions, it is also called ionic bond. An electrovalent bond is polar, i.e., the positive and negative charges are separated.

Figure 3.1: Octate rule

PLANCESS CONCEPTS

Note: (a) In the formula of an ionic compound, the positive ion is written �rst.

(b) Charges on the ions of an ionic compound are usually not shown with the formula. So, sodium chloride is usually expressed as NaCl, not as Na+Cl–.

If you pour a handful of salt into a full glass of water, the water level will actually go down rather than over�owing the glass.

Neeraj Toshniwal

Gold Medalist, INChO

Figure 3.2: Solubility of NaCl


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Illustration 1: Chlorine atom and chloride ions

(a) have an equal number of protons (b) have an equal number of electrons

(c) form covalent bonds (d) react spontaneously with water

Sol: (a). Chlorine atom and chloride ions have an equal number of protons.

1.2 Electrovalency

�e number of electrovalent or ionic bonds an atom can form is called its electrovalency. �e electrovalency of an element is, therefore, equal to the number of electrons lost or gained by the atom to form an ion.

Elements which lose electrons show positive electrovalency and those which gain electrons show negative electrovalency. For example, in the formation of sodium chloride (Na+Cl–), the electrovalency of sodium (Na) is +1, while that of chlorine (Cl) is –1.

Element which lose or gain one, two, three ..., etc. ., electrons are said to be monovalent (or univalent), divalent (or bivalent), trivalent, ....etc., respectively.

Monovalent element: Na, Cl, F

Divalent element : Mg, Ca, Ba, O

Trivalent element : Al, B

Illustration 2: Why NaCl gives a white precipitate with AgNO3 solution but CCl4 does not?

Sol: NaCl is an ionic compound and hence gives Cl- ions in the solution which combine with Ag+

ions given by AgNO3 to form a white precipitate of AgCl but CCl4 is a covalent compound and does not give Cl- ions.

1.2.1 Characteristics of Electrovalent or Ionic compounds

1. Electrovalent compounds are made up of positively and negatively charged ions.

2. Electrovalent compounds have high melting and boiling points. �is is due to the presence of strong electrostatic forces of attraction between the positive and negative ions. A large amount of heat energy is required to break this force of attraction.

3. Electrovalent compounds are usually soluble in water but insoluble in organic solvents such as benzene, acetone, carbon disulphide and carbon tetra chloride.

4. Electrovalent compounds conduct electricity in molten state and in their aqueous solutions.

In solid electrovalent compounds, the ions are held together in �xed positions and cannot move. Hence, such compounds in the solid state do not conduct electricity.

When an electrovalent compound is dissolved in water or is melted, the crystal structure breaks


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down. �e ions now become free to move and can, therefore, conduct electricity. Due to presence of free ions.

PLANCESS CONCEPTS

Ionic bonding is mostly formed between metals located at the left side of the periodic table.

Vipul Singh

AIR 1, NSTSE 2009

Illustration 3: Noble gases exist an individual

(a) Atoms (b) Molecules (c) Ions (d) Compounds

Sol: (a). Atoms

1.2.2 Conditions for Forming Electrovalent or Ionic Bond

(i) An ionic bond is purely electrostatic in nature.

(ii) Its formation is favoured by:

(a) Low ionisation potential (I.P.): of the element that forms a cation on losing electrons (s). �e element should be metal, i.e., electropositive in nature.

(b) High electron a�nity (E.A.): of the element that forms an anion of gaining electron (s). �e element should be non- metal, i.e., electronegative in nature.

(c) High lattice energy (L.E.): �e energy released when isolated ions form a crystal. �e value of lattice energy depends on the charges present on the two ions and distance between them. It shall be high if charges are high and ionic radii are small.

(d) �e summation of three energies should be negative, i.e., energy is released.

I. P. + E. A. + L. E. = – ve

(iii) Highly electropositive elements of groups I and II combine with highly electronegative elements of VI and VII (or 16th and 17th) groups to form electrovalent or ionic compounds. Halides, oxides, sulphides, nitrides and hydrides of alkali and alkaline earth metals are generally ionic.

(iv) Greater the di�erence of electronegativity between two atoms, higher will be the possibility of ionic bond formation.


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PLANCESS CONCEPTS

Partial Ionic character of Covalent Bond: �e extent of partial ionic character is determined by the di�erence in electronegativities of the combining atoms. More is the di�erence in electronegativities, greater will be the ionic character.

(i) If the electronegativity di�erence between two atoms is 1.7, the bond is said to have 50% covalent character.

(ii) If the electronegativity di�erence between two atoms is more than 1.7, the partial ionic character of the bond is more than 50% and the bond is taken as ionic.

(iii) If the electronegativity di�erence between two atoms is less than 1.7, the bond is predominantly covalent.

Shivam Agarwal

Gold Medalist, INChO

�e following types of ions are encountered.

(i) Ions with inert gas con�guration: �e atoms of the representative elements of group I, II and III by complete loss of their valency electrons and the elements of group V, VI and VII by gaining 3, 2 and 1 electrons respectively form ions either with ns2 con�guration or ns2p6 con�guration.

(a) Ions with ns2 (He) con�guration: H–, Li+, Be2+, etc. �e formation of Li+ and Be2+ is di�cult due to their small size and high ionisation energy.

(b) Ions with ns2p6 con�guration: More than three electrons are hardly lost or gained in the ion formation.

Cations: Na+, Ca2+, Al3+, etc.

Anions: Cl–, O2–, N3–, etc.

(ii) Ions with pseudo inert gas con�guration: �e Zn2+ ion is formed when zinc atom loses its outer 4s electrons. �e outer shell con�guration of Zn2+ ion is 3s23p63d10. �e ns2np6nd10 outer shell con�guration is often called pseudo noble gas con�guration which is considered as stable one.

Examples: Zn2+, Cd2+, Hg2+, Cu+, Ag+, Au+, Ga3+, etc.

(iii) Exceptional con�guration: Many d- and f- block elements produce ions with con�gurations di�erent than the above two.

Cr = [Ar] 4s13d5

Cu = [Ar] 4s23d10

Cu+ = [Ar] 3d10 �ese atoms have such con�guration because of the extra stability of the half-�lled and full-�lled

sub shell in these case 3d sub shell.


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However, such ions are comparatively less stable.

(iv) Ions with ns2 con�guration: Heavier member of groups III, IV and V lose p-electrons only to form ions with ns2 con�guration. Ti+, Sn2+, Pb2+, Bi3+ are the examples of this type. �ese are stable ions.

(v) Polyatomic ions: �e ions which are composed of more than one atom are called polyatomic ions. �ese ions move as such in chemical reactions. Some common polyatomic ions are:

4NH (Ammonium); 3NO (Nitrate);

34PO (Phosphate); 2

4SO (Sulphate);

23CO (Carbonate); 2

3SO (Sulphite); etc.

�e atoms within the polyatomic ions are held to each other by covalent bonds. �e electrovalencies of an ion (any type) is equal to the number of charges present on it.

PLANCESS CONCEPTS

In metallic bonding, a large number of atoms lose their electrons.

Anand K

AIR 1, NSO 2011

Illustration 4: Which of the following can lose two electrons to attain the con�guration of argon?

(a) Mg (b) Br (c) S (d) Ca

Sol: (d). Ca

1.3 Lattice Energy

When one mole of a solid is formed from its constituent gaseous ions, the energy released is called the lattice energy.

Energetics of Formation of Ionic Substances: �e energy included in the formation of an ionic compound from its constituent elements may be considered as shown by the Born-Haber cycle for the formation of one mole of sodium chloride from sodium and chlorine.

sublimationS I

Na(s) Na(g) Na(g) e

addition of eDissociation

2 EA1/2D

1Cl (g) Cl(g) Cl(g)

2 ; Crystal formation

UNa(g) Cl (g) NaCl(s)


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Where S = heat of sublimation of sodium metal

I = ionization energy of sodium.

D = heat of dissociation of molecular chlorine.

EA = electron a�nity of chlorine, and U = lattice energy of sodium chloride.

�e amount of heat liberated in the overall reaction is the heat of formation of sodium chloride. From the above

f A1

H S I D E U2

�e most important of these energy terms are I, E and U, since there are considerably greater than the remaining terms S and D.

More the negative value of the heat of formation, greater would be the stability of the ionic compound produced. �us on the basis of the above equation, formation of an ionic compound is favoured by

(a) Low ionization energy (I) of the metal.

(b) High electron a�nity (EA) of the other element.

(c) Higher lattice energy (U) of the resulting compound.

PLANCESS CONCEPTS

1. For dissolution of ionic solid in a solvent (water), hydration energy must be higher than the lattice energy.

2. A number of ionic compounds are almost insoluble in water as their hydration energy is less than lattice energy.

eg., BaSO4, AgCl, PbS, CaCO3, AgBr, AgI, CaF2, Ag2CrO4, PbSO4, etc.

Vaibhav Gupta

AIR 2, NSO

Illustration 5: Electrovalent compounds are usually

(a) Solids with low melting points (b) Solids with high melting points

(c) Volatile liquids (d) Organic compounds

Sol: (b). Solids with high melting points Formation of ions with higher charges: formation of a cation with unit positive charge is easy if

the �rst ionization energy is low as in the case of alkali metals. Alkaline earth metals ionizes in two successive steps.


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Mg Mg e ; 2Mg Mg e

But energy needed to ionise alkaline earth metals are higher than alkali metals. However,

bipositive ions like 2 2 2Mg , Ca , Sr and 2Ba are quite common. Formation of a tripositive

ion like Al3+ requires much more energy (= 5138 kJ) which is not available ordinarily successive ionization energies of aluminium are:

E11Al Al e E 577KJ

;

E 222Al Al e E 1816KJ ;

E 33

3Al Al e E 2745KJ

It is on this account that most of aluminium compounds are covalent. In solution, however, aluminium is known to give hydrated ions [Al.6H2O]3+. �is is possible because of the high heat of hydration of Al3+. �e energy liberated during hydration of ions is de�cient for ionization.

Similarly, anions with unit negative charge are very common. �is is because of electron a�nity of these atoms is positive and quite high. Formation of anion is carrying two units of negative charge (e.g S2-, O2-) is not as easy as their electron a�nities are negative i.e., energy is needed to add second electron. Formation of anions carrying three units of negative charge (e.g N3-,P3-) is almost rare.

�e scientist Fajan suggested that as covalent bonds have some ionic character. In the same way, the ionic bonds also have partial covalent character. �e extent of covalent character of an ionic bond depends on the nature of its constituent cations and anions.

�e percent covalent character of ionic bond is determined by the factors such as polarizing power of the cation, polarizability of the anion and degree of polarization of anion. �erefore, we shall give some idea of the polarizing power of cation and polarizability of anion.

(i) Polarizing power of a cation is its power to distort the electronic charge of the anion. In general, the polarizing power increases as cations become smaller and more highly charged.

(ii) Polarizability of an anion is the extent to which its electronic charge can be distorted. In general, anions with larger size and large negative charge are more polarizable than smaller ones.

PLANCESS CONCEPTS

�e anions such as 2 2I , Br , S , Se etc. are easily polarisable and have longer tendency to induce covalent character in ionic compounds.

Chen Reddy Sandeep Reddy

KVPY Fellow


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Illustration 6: Which of the following statement is wrong?

(a) An atom is electrically neutral

(b) �e size of a cation is smaller than that of the corresponding atom

(c) �e size of an anion is bigger than that of the corresponding atom

(d) An atom and its ion have an unequal number of protons

Sol: (d). An atom and its ion have an unequal number of protons

1.4 Fajan’s Rules

(i) When the size of the cation is small and the size of the anion is large, the covalent character of an ionic bond is large.

Reason: In small cation, the positive charge is concentrated over a small area. �is makes the cation highly polarizing and more e�ective in distorting the negative charge of the large anions.

(ii) When the charge on cation or on anion, or on both ions is large, the covalent character of an ionic bond is large.

Reason: A high charge increases the extent of polarization.

(iii) When the cation does not have a noble gas con�guration, the covalent character of an ionic bond is large.

Explanation: A noble gas con�guration has closed outermost shell and it shields the nuclear charge most e�ectively. �us, a cation with noble gas electronic con�guration has less power to polarize an anion.

A cation without the noble gas con�guration has high positive charge at its surfaces, and thus it has high power to polarize an anion.

Some cations which do not have a noble gas con�guration and favour covalent character:(a) Cations of transition metals such as Ti3+, V3+, Cr2+, Mn2+, and Cu+, Ti+, Pb2+, and Bi3+

(b) Cations of some lanthanide metals such as Ce3+ and Eu2+

Hydrogen Bonding

Nitrogen, oxygen and �uorine are the highly electronegative elements. When they are attached to a hydrogen atom to form a covalent bond, the electrons of the covalent bond are shifted towards the more electronegative atom. �is partially positively charged hydrogen atom forms a bond with the other more electronegative atom. �is bond is known as hydrogen bond and is weaker than the covalent bond. For example, in HF molecule, the hydrogen bond exists between hydrogen atom of one molecule and �uorine atom of another molecule as depicted below :


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Here, hydrogen bond acts as a bridge between two atoms which holds one atom by covalent bond and the other by hydrogen bond. Hydrogen bond is represented by a dotted line ( – – –) while a solid line represents the covalent bond. �us, hydrogen bond can be de�ned as the attractive force which binds hydrogen atom of one molecule with the electronegative atom (F, O or N) of another molecule.

PLANCESS CONCEPTS

Hydro�uoric acid is so corrosive that it will dissolve glass. Although it is corrosive, hydro�uoric acid is considered to be a ‘weak acid’.

T P Varun

KVPY Fellow

1.5 Cause of Formation of Hydrogen Bond

When hydrogen is bonded to strongly electronegative element ‘X’, the electron pair shared between the two atoms moves far away from hydrogen atom. As a result, the hydrogen atom becomes highly electropositive with respect to the other atom ‘X’. Since there is displacement of electrons towards X, the hydrogen acquires fractional positive charge (+) while ‘X’ attain fractional negative charge δ-. �is results in the formation of a polar molecule having electrostatic force of attraction which can be represented as :

�e magnitude of H-bonding depends on the physical state of the compound. It is maximum in the solid state and minimum in the gaseous in�uence on the structure and properties of the compounds.

1.5.1 Types of H-Bonds

There are two types of H-bonds

(i) Intermolecular hydrogen bond

(ii) Intramolecular hydrogen bond

(i) Intermolecular hydrogen bond: It is formed between two di�erent molecules of the same or di�erent compounds. For example, H-bond in case of HF molecule, alcohol or water molecules, etc.

(ii) Intramolecular hydrogen bond: It is formed when hydrogen atom is in between the two highly electronegative (F, O, N) atoms present within the same molecule. For example, in o-nitrophenol, the hydrogen is in between the two oxygen atoms.


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1.6. Covalent Bond

“A chemical bond formed by sharing one or more electron pairs between atoms when each atom contributes equally is called a covalent bond”.

Covalent bond may be single, double or a triple bond. Double and triple covalent bonds are called multiple covalent bonds. Single covalent bond is formed by sharing of only one electron pair. �is bond is represented by single dash. Double and triple covalent bonds are formed when atoms bonded together share two or three electron pairs, respectively. �ese bonds are represented by double dash ( = ) and triple dash ( ) respectively. Some examples of covalent bonding are given below:

(i) Formation of hydrogen molecule: In the formation of hydrogen molecule, each hydrogen atom contributes one electron and then the pair is shared between two atoms. Both the atoms acquire stable con�guration of helium. �us, the molecule consists of one single covalent bond.

(ii) Formation of chlorine molecule: Chlorine atom has seven electrons in the valence shell. In the formation of chlorine molecule, each Chlorine atom contributes one electron and then the pair of electrons is shared between two atoms. Both the atoms acquire the stable con�guration of argon.

(iii) Formation of O2 molecule: Each oxygen atom contributes two electrons and two pairs of electrons are then shared equally. Both the atoms acquire the con�guration of neon.

(iv) Formation of N2 molecule: Nitrogen atom have �ve valence electrons. Both nitrogen atoms achieve con�guration of neon by sharing three pairs of electrons i.e. each atom contributes 3 electrons.

Illustration 7: Which amongst the following contains a covalent bond?

(a) MgCl2 (b) CaF2 (c) Al2O3 (d) HCl

Sol: (d). HCl


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PLANCESS CONCEPTS

�e covalency of an atom is the number of its electrons taking part in the formation of shared pairs. �us, the covalency of hydrogen is one, oxygen two, nitrogen three and carbon four.

Vijay Senapathi

KVPY Fellow

Illustration 8: In each of the following compounds, there is at least one multiple bond. Draw a Lewis structure for each molecule. Use lines to indicate electrons in bonds and dot to indicate nonbonding electrons.

(a) H3CCN (b) H3CNO

Sol: (a) In methyl group (CH3), three covalent bonds are constructed by sharing three of carbon’s four available electrons with the single electrons of the three hydrogens.

�e other atoms, the carbon and the nitrogen, have four and �ve electrons respectively in their valence shell.

Formation of a triple bond between carbon and nitrogen complete the picture, and nitrogen is left with a non-bonding pair of electrons.

(b) H3CNO – Once again there is a methyl group. Nitrogen has �ve available electrons, and oxygen has 6 electrons. A carbon – nitrogen single bond and nitrogen – oxygen double bond can be formed. Nitrogen is left with one pair of nonbonding electrons, and oxygen with two pairs.

PLANCESS CONCEPTS

Since noble gases have a naturally full outer shell they seldom react.

Krishan Mittal

KVPY Fellow


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Illustration 9: Which of the following indicates that the bonding in hydrogen chloride is covalent?

(i) An aqueous solution of hydrogen chloride form hydrogen with metals

(ii) Hydrogen chloride is a gas

(iii) �e di�erence in the electronegativity between hydrogen and chlorine is less than 1.7

(iv) Dry hydrogen chloride is a non-conductor of electricity

(a) (i), (iii) and (iv) only (b) (i), (ii) and (iii) only

(c) (i) and (iv) only (d) (i), (ii), (iii) and (iv)

Sol: Correct answer is (a).

1.6.1 Co-ordinate Bond

A coordinate bond is de�ned as the chemical bond between two atoms in which a pair of electrons is contributed by one atom but shared by both the atoms.

Mechanism of formation of coordinate bond: A coordinate bond is formed when an atom in a molecule contains lone pair of electrons and the other atom or ion is de�cient of electrons. �e atom that contributes the pair of electrons is called a donor and the other one which merely shares it is called an acceptor. �erefore, the coordinate bond is also known as donor-acceptor bond or dative bond. A dative bond is represented by an arrow drawn between the two atoms of the molecule or ion. �e arrow head of the coordinate bond is directed towards the acceptor.

(A) Formation of coordinate bond in ammonium ion NH4+:

A molecule of ammonia has a lone pair of electrons on nitrogen atom: ( : NH3) while H+ ion does not have any electron. �erefore, when ( : NH3) and H+ interact, the lone pair of electrons on N atom of ( : NH3) is donated and then shared by both N and H as

Note: In ammonium ion, though one N—H bond is a coordinate bond, it is similar in length and strength to the other three N—H bonds. Why? Because once a bond is formed by sharing, it becomes di�cult to distinguish one electron pair from the others.

Illustration 10: �e bonding in ammonium chloride

(a) Is covalent only

(b) Is electrovalent only


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(c) Consists of three covalent nitrogen-hydrogen bonds and an electrovalent bond between the ammonia molecule and the chlorine atom.

(d) Consist of four covalent nitrogen-hydrogen bonds and one electrovalent bond between the ammonium ion and chloride ion.

Sol: (d). Consist of four covalent nitrogen-hydrogen bonds and one electrovalent bond between the ammonium ion and chloride ion.

(B) Formation of coordinate bond in ammonia boron tri�uoride complex:

In ammonia molecule (: NH3), the octet of N atom is complete and it contains a lone pair of electrons. However, in BF3 ,the octet of each F atom is complete but B atom is electron de�cient as there are only six electrons around it. �erefore, when NH3 and BF3 molecules come in closer contact, the lone pair of electrons on N atom of NH3 is donated to B atom of BF3, and a coordinate bond between N and B is formed.

Now, with the formation of a coordinate bond, the octet of boron atom is also complete.

PLANCESS CONCEPTS

A pair of electrons which is not shared with any other atom is known as the lone pair of electrons but it is provided to the other atom for the formation of coordinate bond.

Krishan Mittal

KVPY Fellow

Illustration 11: Sodium atom and sodium ion

(a) Are chemically the same (b) Have the same number of protons

(c) Have the same number of electrons (d) React spontaneously with water

Sol: (b). Have the same number of protons

1.7 Valence Bond Theory

�is theory was proposed by Heitler and London, in 1927, to explain how a covalent bond is formed. �is theory was extended by Pauling and Slater in 1931.�e main postulates of the theory are:

(a) A covalent bond is formed by overlapping of atomic orbitals of valency shell of the two atoms.


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(b) Only half �lled atomic orbitals. i.e,orbitals singly occupied can enter into overlapping process.

�e resulting bond acquires a pair of electrons with opposite spins.

(c) �e atoms with half-�lled orbitals must come closer to one another with their axes in proper directions for overlapping.

(d) As a result of overlapping, there is maximum electron density somewhere between the two atoms. A large part of bonding force come into existence from the electrostatic attraction between the nuclei and the accumulated electron cloud between them.

Figure 3.3: Electron distribution according to VBT

(e) Greater the overlapping, higher is the strength of the chemical bond. �e amount of energy released per mole during overlapping is termed bond energy. �is energy stabilises the system. Hence, the molecule formed has less energy and consequently more stability than the isolated atoms.

(f ) Electrons which are already paired in valency shell can enter into bond formation if they can be unpaired �rst and shifted to vacant orbitals of slightly higher energy of the same main energy shell (valency shell). �is point explains the trivalency of boron, tetravalency of carbon, pentavalency of phosphorus, hexavalency of sulphur and heptavalency of halogens (Cl, Br, I) inspite of the fact that these atoms have paired orbitals in the valency shell.

(g) Between two orbitals of the same stability (i.e., having same energy) one more directionally concentrated would form a stronger bond. Dumb-bell shaped p-orbitals will form stronger bond as compared to spherically symmetrical s-orbital. It is formed by head on or axial overlap.

Two types of bonds are formed on account of overlapping. �ese are (a) Sigma and (b) Pi

(a) Sigma (σ) Bond: A bond formed between two atoms by the overlap of singly occupied orbitals along their axes (end to end overlap) is called sigma bond. In such a bond formation, maximum overlap is possible between electron cloud of this bond is symmetrical about the line joining the two nuclei of the two atoms. Sigma bond can, thus, be de�ned as:


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“Bond orbital which is symmetrical about the line joining the two nuclei is known as sigma bond.” It is formed by head on or axial overlap.

Sigma bonds are formed by three types of overlapping:

(i) s-s overlapping (Formation of hydrogen molecule): Each hydrogen atom has one electron in 1s-orbital which is spherical. 1s-orbital of both the hydrogen atoms approach each other closely and when they reach a point of maximum attraction by the two nuclei, they overlap and form a sigma bond.

Figure 3.4: s-s overlapping in H2 molecule

�e bond has two electrons which have opposite spins. �e probability of �nding these electrons is maximum in the region between the two nuclei on the molecular axis. �e electron density of the bond is distributed symmetrically about the molecular axis.

(ii) s-p overlapping (Formation of HF, H2O, NH3 molecules):

(a) Formation of HF molecule: In the formation of HF molecule the 1s-orbital of hydrogen overlaps with the p-orbital of �uorine containing unpaired electron.

Figure 3.5: s-p overlapping in HF

(b) Formation of water molecule: Oxygen atom has the con�guration of valency shell 1s2pz

2py1px

1, i.e., it has two orbitals singly occupied. �ese two orbitals overlap with 1s-orbital of two hydrogen atoms forming sigma bonds. Since the two orbitals of oxygen are at right angles to each other, an angle of 90 is expected between two sigma bonds but actual bond angle observed is 104.5 .

Figure 3.6: s-p overlapping in water


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(c) Formation of ammonia molecule: Nitrogen atom has the con�guration of valency shell 2s2 2px

1 2py1 2pz

1, i.e. three singly occupied orbitals are present. �ese orbitals overlap with 1s-orbitals of three hydrogen atoms forming three sigma bonds.

Since the three orbitals of nitrogen are at right angles to each other, the expected angle between two sigma bonds should be 90° but actual bond angle observed is 107°.

Figure 3.7: s-p Overlapping in NH3

(iii) p-p overlapping (Formation of �uorine molecule):

�is is illustrated by the formation of �uorine molecule. �e electronic con�guration of �uorine atom is 1s2 2s2 2px

2 2py22pz

1, i.e., one orbital is singly occupied. When p-orbitals of two �uorine atoms approach each other with their heads directly towards one another, they overlap and form a sigma bond.

(b) Pi (π) Bond: Bonds are formed by the sidewise or lateral overlapping of p-orbitals. �e overlapping takes place at the side of two lobes and hence, the extent of overlapping is relatively smaller. �us, π-bond is a weaker bond in comparison to sigma bond. �e molecular orbital is oriented above and below the plane containing nuclear axis.

Formation of oxygen molecule: Oxygen atom has two p-orbitals singly occupied in the valency shell. When two oxygen atoms approach each other, one set of p-orbitals experiences head on overlaps axially to form σ bond and the other set of the orbitals overlaps sidewise to form a π

Figure 3.8: p-p overlapping in F2

Figure 3.9: p-p overlapping


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-bond. �us, oxygen molecule has one σ-and one π-bond.

Similarly, the formation of nitrogen molecule can be explained. It has one sigma bond and two π bonds.

Figure 3.10: p-p overlapping in oxygen molecule

Illustration 12: An element with low ionization energy combines with an element having high electron a�nity to form

(a) Ionic bond (b) Covalent bond

(c) Dative bond (d) None of these

Sol: Ionic bond

1.8 Hybridisation

In order to explain the characteristic geometrical shapes of polyatomic molecules like CH4, NH3 and H2O etc., Pauling introduced the concept of hybridisation. According to him the atomic orbitals combine to form new set of equivalent orbitals known as hybrid orbitals. Unlike pure orbitals, the hybrid orbitals are used in bond formation. �e phenomenon is known as hybridisation which can be de�ned as the process of intermixing of the orbitals of slightly di�erent energies so as to redistribute their energies, resulting in


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the formation of new set of orbitals of equivalent energies and shape. For example when one 2s and three 2p-orbitals of carbon hybridise, there is the formation of four new sp3 hybrid orbitals. NH3 and H2O etc., Pauling introduced the concept of hybridisation. According to him the atomic orbitals combine to form new set of equivalent orbitals known as hybrid orbitals. Unlike pure orbitals, the hybrid orbitals are used in bond formation. �e phenomenon is known as hybridisation which can be de�ned as the process of intermixing of the orbitals of slightly di�erent energies so as to redistribute their energies, resulting in the formation of new set of orbitals of equivalent energies and shape. For example when one 2s and three 2p-orbitals of carbon hybridise, there is the formation of four new sp3 hybrid orbitals.”

Salient features of hybridisation: �e main features of hybridisation are as under –

1. �e number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridised.

2. �e hybridised orbitals are always equivalent in energy and shape.

3. �e hybrid orbitals are more e�ective in forming stable bonds than the pure atomic orbitals.

4. �ese hybrid orbitals are directed in space in some preferred direction to have minimum repulsion between electron pairs and thus a stable arrangement. �erefore, the type of hybridisation indicates the geometry of the molecules.

Important conditions for hybridisation :

1. �e orbitals present in the valence shell of the atom are hybridised.

2. �e orbitals undergoing hybridisation should have almost equal energy.

3. Promotion of electron is not essential condition prior to hybridisation.

4. It is not necessary that only half �lled orbitals participate in hybridisation. In some cases, even �lled orbitals of valence shell take part in hybridisation.

Types of Hybridisation:

�ere are various types of hybridisation involving s, p and d orbitals. �e di�erent types of hybridisation are under –

(I) sp hybridisation: �is type of hybridisation involves the mixing of one s and one p orbital resulting in the formation of two equivalent sp hybrid orbitals. �e suitable orbitals for sp hybridisation are s and pz, if the hybrid orbitals are to lie along the z-axis. Each sp hybrid orbitals has 50% s-character and 50% p-character. Such a molecule in which the central atom is sp-hybridised and linked directly to two other central atoms possesses linear geometry. �is type of hybridisation is also known as diagonal hybridisation.

�e two sp hybrids point in the opposite direction along the z-axis with projecting positive lobes and very small negative lobes, which provides more e�ective overlapping resulting in the formation of stronger bonds.

Examples of molecule having sp hybridisation:

BeCl2: �e ground state electronic con�guration of Be is 1s22s2. In the excited state one of the 2s -electronic is promoted to vacant 2p orbital to account for its divalency. One 2s and


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one 2p-orbitals get hybridised to form two sp hybridised orbitals. �ese two sp hybrid orbitals are oriented in opposite direction forming an angle of 180°. Each of the sp hybridised orbital overlaps with the 2p-orbital of chlorine axially and form two Be-Cl sigma bonds.

Figure 3.11: Formation of BeCl2

(II) sp2 hybridisation: In this hybridisation there is involvement of one s and two p-orbitals in order to form three equivalent sp2 hybridised orbitals. For example, in BeCl3 molecule, the ground state electronic con�guration of central boron atom is 1s22s22p1. In the excited state, one of the 2s electrons is promoted to vacant 2p orbital as a result boron has three unpaired electrons. �ese three orbitals (one 2s and two 2p) hybridise to form three sp2 hybrid orbitals. �e three hybrid planar arrangement and overlap with 2p orbitals of chlorine to form three Be-Cl bonds. �erefore, in BeCl3 , the geometry is trigonal planar with ClBeCl bond angle of 120°.

In BeCl3 , the geometry is trigonal planar with ClBeCl bond angle of 120°.”

(III) sp3 hybridisation: �is type of hybridisation can be explained by taking the example of CH4 molecule in which there is mixing of one s-orbital and three p-orbitals of the valence shell to form four sp3 hybrid orbital of equivalent energies and shape. �ere is 25% s-character and 75% p-character in each sp3 hybrid orbital. �e four sp3 hybrid orbitals so formed are directed towards the four corners of the tetrahedron. �e angle between sp3 hybrids orbital is 109.5°.

1.8.1 Bond Characteristics

Covalent bonds are characterised by the following four parameters:

(1) Bond length (2) Bond energy

(3) Bond angle (4) Bond order

(1) Bond Length:

�e distance between the nuclei of two atoms bonded together is termed bond length or bond distance. It is expressed in angstrom (Å) or picometre (pm) units:

[1Å = 10-8 cm; 1 pm = 10-12 m]


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Bond LengthsTable 3.1: Bond Lengths

Bond Bond Length 0

(A)

H H (in 2H ) 0.74

Cl Cl (in 2Cl ) 1.99

Br Br (in 2Br ) 2.28

I I (in 2I ) 2.67

F F (in 2F ) 1.42

C – C 1.54

C C (in 2 4C H ) 1.34

C C (in 2 2C H ) 1.21

C N (in amines) 1.47

C H (in organic molecules) 1.08N O 1.36

H F (in HF) 0.92

H Cl (in HCl) 1.27

H – Br (in HBr) 1.41H – I (in HI) 1.61O – H (in H2O) 0.96O = O (in O2) 1.21C – O (in alcohols) 1.43C = O (in aldehydes) 1.22N –H 1.03S – H 1.35N = O 1.22

Bond lengths are measured spectroscopically or by X-ray di�raction of solids and electron di�raction of gases. In ionic compounds, bond length is obtained by adding up of the ionic radii of two bonded ions.


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Bond length in ionic compound = rc+ + rA–

Similarly, in a covalent compound, bond length is obtained by adding up the covalent (atomic) radii of two bonded atoms.

Bond length in covalent compound (AB) = rA + rB

�e factors such as resonance, electronegativity, hybridization, steric e�ects, etc., which a�ect the radii of atoms, also apply to bond lengths, i.e., the values of bond lengths do not hold good for all the compounds. Bond lengths of some common bonds are tabulated in the table given above.

Important features

(1) Bond length

(i) �e bond length of the homonuclear diatomic molecules are twice the covalent radii. Bond length increases with the increase in the size of atoms.

H2 C—H C—Cl

0.74 Å 1.09 Å 1.77 Å

(ii) �e lengths of double bonds are less than the lengths of single bonds between the same two atoms and triple bonds are even shorter than double bonds.

Single bond > Double bond > Triple bond

(iii) Bond length decreases with increase in s-character since s-orbital is smaller than a p-orbital.

sp3 C— H = 1.112 Å; sp2 C— H = 1.103 Å;

(25% s-character as in alkanes) (33.3% s-character as in alkenes)

sp C— H = 1.08 Å

(50% s- character as in alkynes)

(iv) Bond length of polar bond is smaller than the theoretical non-polar bond length.

(2) Bond Energy or Bond Strength

Bond energy or bond strength is de�ned as the amount of energy required to break a bond in a molecule. Each bond has a characteristic value of this energy and is measure of the strength of the bond. It is generally observed that shorter the bond length, greater is the bond strength or bond energy of the bond.

1H H(g) 2H(g); H 436 KJmol (Bond length = 0.74

0A )

1Cl Cl (g) 2Cl (g); H 242 KJmol (Bond length = 1.99

0A )

1H Cl (g) H (g) Cl(g); H 431 KJmol (Bond length = 1.27

0A )

Consider the dissociation of water molecule which consists of two O – H bonds.


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H O H(g) 2H(g) O(g) ; 1H 926kJmol

�e average bond energy of O – H bond 926463

2 kJ

Similarly the average bond energy of C – H bond in CH4 is equal to one fourth of the energy of dissociation of CH4 into (C + 4H). However, the bond energy of each C – H bond is CH4 is di�erent.

14 3CH (g) CH (g) H(g); H 426kJmol

1

3 2CH (g) CH (g) H(g); H 439kJmol

1

2CH (g) CH (g) H(g); H 451kJmol

1CH (g) C(g) H(g); H 374kJmol

1

4CH (g) C(g) 4H(g); H 1663kJmol

Hence, average bond energy 11663C H 416 kJ mol

4

�us, the bond energy of C - H bond depends on the order in which the particular hydrogen atom is lost from the molecule. A similar situation exists for all molecules with more than two atoms. �e strengths of bonds depends on the order in which they are broken. Average bond energies of the some common bonds are listed below:

Bond energies in kJ mol-1

Table 3.2: Bond energies for diatomic molecules

Bond Bond energy Bond Bond energy

H –H 436 Cl –Cl 242

O =O 497 Br – Br 193

N N 945 I – I 151

F – F 158 H –F 563

H - Cl 431 N – H (in NH3) 389

H – Br 366 O –H (in water) 464

H – I 299 C – C 346

C – H (Methane) 416 C = C (in C2H4) 598


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1. �e magnitude of the bond energy depends on the type of bonding. Most of the covalent bonds have energy between 50 to 100 Kcal mol–1 (200-400 kJ mol–1). Strength of sigma bond is more than that of a π-bond.

2. A double bond in a diatomic molecules has a higher bond energy than a single bond and a triple bond has a higher bond energy than a double bond between the same atoms.

C C C C C — C

3. �e magnitude of the bond energy depends on the size of the atoms forming the bond, i.e., bond length. Shorter the bond length, higher is the bond energy.

4. �e bond energy decreases with increase in number of lone pairs on the bonded atom. �is is due to electrostatic repulsion of lone pairs of electrons on the two bonded atoms.

5. Homolytic and heterolytic �ssion involve di�erent amounts of energies. Generally the values are low for homolytic �ssion of the bond in comparison to heterolytic �ssion.

6. Bond energy decreases down the group in case of similar molecules.

7. Bond energy increases in the following order:

s < p < sp < sp2 < sp3

C—C > N—N > O—O

346 163 146.4 kJ mol–1

(No lone pair) (One lone pair) (Two lone pairs)

(3) Bond Angles

Angle between two adjacent bonds at an atom in a molecule made up of three or more atoms is known as the bond angle. Bond angles mainly depend on the following three factors:

(i) Hybridization: Bond angle depends on the state of hybridization of the central atom.

Hybridization sp3 sp2 sp

Bond angle 109°28’ 120° 180°

Example CH4 BCl3 BeCl2

It is observed that as s-character increases in the hybrid bond, the bond angle increases.

(ii) Lone pair repulsion: Bond angle is a�ected by the presence of lone pair of electrons at the central atom. A lone pair of electrons at the central atom always tries to repel the shared pair (bonded pair) of electrons. Due to this, the bond are displaced slightly inside resulting in a decrease of bond angle. �e bond angle in NH3 is 107° and bond angle in H2O is 104.5° inspite of the fact that N-atom and O-atom undergo sp3 hybridization.


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(iii) Electronegativity: If the electronegativity of the central atom decreases, bond angle decreases.

In case the central atom remains the same, bond angle increases with decrease in the electronegativity of the surrounding atoms.

(4) Bond order: Bond order is de�ned as the number of bonds or number of shared electron pairs between two atoms in a molecule.

Isoelectronic molecules or ions have same bond order.

For example, F2 and O22- (18 electrons) have bond order 1. N2, CO and NO+ (14 electrons) have

bond order three. It is observed that with increase in bond order, bond enthalpy increases while bond length decreases.

Illustration 13: Which amongst the following has a polar covalent bond?

(a) NaCl (b) NH3 (c) O2 (d) CaO

Sol: (b). NH3

1.9 Valence Shell Electron Pair Repulsion (VSEPR) Theory

Lewis-Langmuir concept of covalent bonding is unable to explain the shapes of the covalent molecules. In case of covalent molecules having three or more atoms, one of the atoms acts as a central atom and rest of the atoms are linked to it de�nite directions in space. �is de�nite arrangement of the bonded atoms with the central atom in a molecule is known as shape or geometry of the molecule.

�e �rst simple theory that was put forward to predict the geometry or shape of a covalent molecule is known as valence shell election pair repulsion theory (or VSEPR theory). �is theory was further developed by Nyholm and Gillespie in 1957.

�e theory, known as VSEPR theory, is primarily based on the fact that in a polyatomic molecule, the direction of bond around the central atom depends upon the total number of electron pairs (bonding as well as non-bonding) in its valence shell. �ese electron pairs place themselves as far apart as possible in space in order to have minimum forces of repulsion between them. �e minimum repulsions correspond to the state of minimum energy and maximum stability of the molecule.


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�e main postulates of VSEPR theory are as follows:

i. �e shape of a molecule depends upon the number of valence shell electron pairs (bonded or non-bonded) around the central atom.

ii. Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.

iii. �e positions of the electron pairs in space around central atom are such that minimise repulsion and thus maximise distance between them.

iv. �e magnitudes of the di�erent types of electronic repulsions follows the order given below:

lone pair — lone pair > lone pair— bond pair > bond pair — bond pair

�ese repulsive forces alter the bond angles of the molecule or ion.

v. �e electronic repulsion between two pairs of electrons will be minimum if they are as far apart as possible. On this basis, the following geometrical arrangements are most suited.

Table 3.3: Geometry of molecule according to the number of lone pair

No .of electron pairs Geometrical �gure Examples

2 Linear2 2 2BaCl ,CO ,ZnCl

3 Triangular planar 23 3 3BCl ,SO ,CO

4 Tetrahedral4 2 3CH ,H O,NH

5 Trigonal bipyramidal5 3 4PF ,CIF ,SF

6 Octahedral 26 5 6SF ,IF ,SiF

If the central atom is linked to similar atom and is surrounded by bond pairs of electrons, the repulsions between them are similar. As a result the shape of the molecule is symmetrical and the molecule is said to have a regular geometry. On the other hand, the molecules in which the repulsive interactions between the electron pairs around the central atom are unequal have irregular or distorted geometry. In fact, the presence of lone pairs in addition to bond pairs in the molecule causes distortion in the geometry of the molecule. For example, in the NH3 and H2O molecules, the bond angles are not 109°28’ but 107° and 104.5° respectively due to presence of one lone pair in NH3 and two lone pairs in H2O.

[In the VSEPR notation used to describe molecular geometries, the central atom is denoted as A, terminal atoms as X and lone pair as E. AX2E2 describes a structure with two terminal atoms and two lone pair around a central atom, for example water molecule. Similarly, AX3E2 describes a structure with three terminal atoms and two lone pairs around a central atom. For example ClF3 molecule.]


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A multiple bond is treated as if it is a single electron pair and two or three electron pairs of a multiple bond are treated as a single pair.

A Strategy for Applying VSEPR Method:

To predict the shape of a molecule or polyatomic ion, the following four steps can be followed:

Step 1: Draw a Lewis structure of the molecule or polyatomic ion. �e structure must be plausible but it does not need to be the best one.

Step 2: Determine the number of electron groups around the central atom (bonding and non-bonding both). A double and a triple bond each is counted as one electron group.

Step 3: Identify the electron group geometry. �is may be linear, trigonal planar, tetrahedral, trigonal bipyramidal or octahedral corresponding to two, three, four, �ve and six electron groups respectively.

Step 4: Identify the molecular geometry. �is is based on the positions around the central atom occupied by other atoms (not by lone pairs).

Illustration of the �eory by Considering a Few Examples:

Molecules containing bond pairs only

Shape of BeF2 molecule: Be atom is surrounded by two bond pairs of electrons. �ese should be localised in such a way that there is minimum repulsion between them, i.e., bond angle should be 180°, hence BeF2 molecule is linear.

Molecules such as BeCl2, ZnCl2, HgCl2, etc., have a linear shape.

Shape of BF3 molecule: �e Lewis structure of BF3 is

B-atom is surrounded by three bond pairs. According to VSEPR theory they must be situated at an angle of 120° as they will experience minimum repulsion, i.e., shape of BF3 is triangular planar.

Molecules such as BCl3, AlCl3, etc., have a triangular planar shape.

1.10 Molecular Orbital Theory

�e main features of the molecular orbital theory can be summed up as follows:

Like valence bond theory, this theory starts with atomic orbitals but the atomic orbitals of the atoms approaching for bonding overlap to undergo constructive interference as well as destructive interference to form molecular orbitals. As a result, the atomic orbitals lose their individual identity and all the electrons in the molecule are associated with molecular orbitals.

1. When two atomic orbitals overlap, they form two new orbitals called molecular orbitals. One of which is called bonding molecular orbitals and other is called antibonding molecular orbital. �ese are formed by addition and subtraction of wave functions respectively.


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2. Molecular orbitals are the energy states of a molecule in which the electrons of the molecule are �lled.

3. Bonding molecular orbital has energy lower than the combining atomic orbitals while antibonding orbital has higher energy than the combining atomic orbitals.

4. Only those atomic orbitals can overlap to form molecular orbitals which have comparable energies and proper orientation.

5. Electrons present in the bonding molecular orbital contribute towards the stability of molecule while electrons present in antibonding molecular orbital contribute to the repulsions between the nuclei of the atoms.

6. �e bonding molecular orbitals are denoted as etc., while antibonding molecular orbitals are denoted as etc.

7. A molecular orbital (bonding or antibonding cannot accommodate more than two electrons. Both the electrons must have opposite spins (Pauli’s exclusion principle).

8. Molecular orbitals are �lled in order of increasing energies starting with the orbitals of minimum energy (Aufbau principle).

8. In molecular orbitals of same energy (degenerate orbitals), the electron pairing occurs only when all of them are singly �lled (Hund’s rule).

9. �e shapes of the molecular orbitals formed depend upon the type of the combining orbitals.


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SUMMARY

• Electrovalent, covalent and co-ordinate bonding, structures of various compounds – orbit structure and electron dot structure.

• Structure of electrovalent compounds NaCl, MgCl2, CaO

Characteristic properties of electrovalent compounds – state of existence, melting and boiling points, conductivity (heat and electricity), ionization in solution and in molten state to be linked with electrolysis.

• Covalent Bond: de�nition and examples, structure of covalent molecules on the basis of duplet and octet of electrons (example: hydrogen, chlorine, nitrogen, water, ammonia, carbon tetrachloride, methane.)

• Characteristic properties of Covalent compounds: state of existence, melting and boiling points, conductivity (heat and electricity), ionization in solution.

Comparison of Electrovalent and covalent compounds.

• De�nition of Coordinate Bond: �e loan pair e�ect of the oxygen atom of the water molecule and the nitrogen atom of the ammonia molecule to explain the formation of 3H O and OH- ions in water and 4NH ion. �e meaning of lone pair; the formation of hydronium ion and ammonium ion must be explained with the help of electron dot diagrams.

• �eories of Bonding-VSEPR, molecular orbital theory and valence Band theory.


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SOLVED EXAMPLES

Multiple Choice Questions

Example 1. Number of non-bonding electrons in N2 is

(a) 4 (b) 10 (c) 12 (d) 14

Sol: 4

Example 2. �e nature of bonding in CCl4 and CaH2 is that

(a) Covalent in CCl4 and electrovalent in CaH2

(b) Electrovalent in both CCl4 and CaH2

(c) Covalent in both CCl4 and CaH2

(d) Electrovalent in CCl4 and covalent in CaH2

Sol: Covalent in CCl4 and electrovalent in CaH2

Example 3. Bond formed between elements whose electronegativities are 1.2 and 3.0, would be

(a) Ionic (b) Covalent

(c) Co-ordinate (d) Metallic

Sol: Ionic

Example 4. KCl easily dissolves in water because

(a) Heat of hydration is more than lattice energy

(b) It reacts with water

(c) It is an electrovalent compound

(d) Its ions are easily solvated

Sol: Both (a) and (d)


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Match the Following Columns

Example 5. Match the elements whose atomic numbers are given in Column I with the type of bond they are likely to form Column II.

Column I Column II

(A) Z = 1 and Z = 17 (p) Pure covalent

(B) Z = 7 and Z = 7 (q) Polar covalent

(C) Z = 11 and Z = 17 (r) No Bond

(D) Z = 2 and Z = 17 (s) Ionic Bond

Sol: A-q; B-p; C-s; D-r

Example 6. Match the molecules mentioned in Column I with the bond description in Column II.

Column I Column II

(A) Water molecules (p) One lone pair, three bond pairs

(B) Ammonia molecule (q) Four bond pairs

(C) Hydrogen chloride (r) Two lone pairs, two bond pairs

(D) Carbon dioxide (s) One bond pair

Sol: A-r; B-p; C-s; D-q

Assertion and Reason

(a) Both assertion and reason are correct and reason is the correct explanation of assertion

(b) Both assertion and reason are correct, but Reason is not the correct explanation of assertion.

(c) Assertion is correct but reason is incorrect.

(d) Assertion is incorrect but reason is correct.

Example 7. Assertion: Ionic bond is non–directional.

Reason: Each ion is surrounded by a uniformly distributed electric �eld.

Sol: Assertion and Reason both are true and Reason is correct explanation for Assertion.

Example 8. Assertion: As lattice energy increases, melting and boiling points of ionic compound increases.

Reason: As lattice energy increases, stability of ionic compound increases.



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Example 9. Assertion: Bond energy has order like C–C < C = C < C C.

Reason: Bond energy increases with increase in bond order.


Example 10. Assertion: In NH3, N is sp3 hybridized, but angle is found to be 107o.

Reason : �e decrease in bond angle is due to repulsion between the lone pair on nitrogen and bond pair between N and H.


Example 11. Assertion: Be (among alkaline earth metals) predominantly forms covalent bond.

Reason: Be is smaller in size and hence has greater polarising power.


Example 12. Assertion: pi bonds are weaker than σ - bonds.

Reason : pi bonds are formed by the overlapping of p-p orbitals along their axis.



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EXERCISE 1 – For School Examinations

Fill in the Blanks

Directions: Complete the following statements with an appropriate word/term to be �lled in the blank space(s).

Q.1 are formed by the sharing of electrons between two atoms.

Q.2 �e bond between sodium and chloride ions is bond.

Q.3 Solid NaCl is a conductor of electricity.

Q.4 Formation of bond the distance between the two concerned atoms.

Q.5 �ere are bonds in a nitrogen molecule.

Q.6 Two atoms of similar electronegativity are expected to form compounds.

Q.7 Multiple bonds exists in molecule of

Q.8 �e electronegativity next to F is

Q.9 Conversion of sodium atom into sodium ion is an process.

Q.10 Conversion of chlorine atom into chloride ion is an process.

True / False

Directions: Read the following statements and write your answer as true or false.

Q.11 Linear overlap of two atomic orbits leads to a bond.

True False

Q.12 In water molecule, the bond angle in H – O – H is 104.5 .

True False

Q.13 �e dipole moment of 3CH F is greater than that of 3CH Cl .

True False

Q.14 �e energy of a carbon – carbon bond is less than the energy of carbon – carbon bond.

True False

Q.15 All molecules with polar bonds have dipole moment.

True False


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Q.16 Covalent bonds is formed between two atoms are directional whereas ionic bonds are non-directional.

True False

Q.17 A polar bond is formed between two atoms of same electronegativity.

True False

Q.18 Element X is strongly electropositive and element Y is strongly electronegative. Both are univalent. �e compound formed is X+Y-.

True False

Q.19 In its covalent compound, the hydrogen atom has the electronic con�guration analogous to that of neon.

True False

Q.20 34PO ion exhibits phenomenon of resonance.

True False


Directions: Each question contains statements given in two columns which have to be matched. Statements (A, B, C, D) in column I have to be matched with statements (p, q, r, s) in Column II

Q.21

Column I Column II

(A) A B 0

(p) bond A – B is more ionic and less covalent.

(B) A B 1.7

(q) bond A – B is less ionic and more covalent.

(C) A B 1.7

(r) bond A – B is 100% covalent or non-polar

(D) A B 1.7

(s) bond A – B is 50% covalent and 50% ionic or or polar covalent bond.

Very Short Answer Questions

Directions: Give answer in one word or one sentence.

Q.22 What is the electric charge on the calcium ion in calcium chloride, 2CaCl ?

Q.23 How many electrons in an oxygen atom able to attract from other atoms?


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Q.24 How many covalent bonds in an oxygen atom able to form?

Q.25 What is dipole?

Q.26 How many electrons does the calcium atom tend to lose?

Q.27 Why covalent solids having the giant molecules are insoluble in all the solvents?

Q.28 Why the reaction of covalent compound are slow?

Q.29 If the electro negativities of H, N, P, As, Sb are 2. 1, 3. 0, 2. 1, 2.0 and 1.9 respectively, then arrange in order of decreasing polarity of bonds in 3 3 3 3SbH ,AsH ,PH ,NH .

Q.30 What would be the electron dot structure of carbon dioxide which has formula 2CO ?

Q.31 What would be the electron dot structure of a molecule of sulphur which is made up of eight atoms of sulphur?

Q.32 Which electrons take part in bond formation?

Q.33 What type of orbitals can overlap to form a covalent bond?

Q.34 What orbitals can overlap to form a σ -bond and which orbitals can do so to form a π - bond?

Q.35 Why free rotation about a π- bond is not possible?

Q.36 What changes in energy takes place when a molecule is formed from its atoms?

Q.37 What type of forces holds the atoms together in an ionic compound?

Short Answer Questions

Directions: Give answer in two to three sentences.

Q.38 Where valence electrons are located, and why are they important?

Q.39 What is the chemical formula for the ionic compound magnesium oxide?

Q.40 How many non-bonding pairs are there in the valence shell of an oxygen atom?

Q.41 Why does the �uorine atom tend to gain only one electron?

Q.42 What force holds two atoms together in a covalent bond?

Q.43 Which is more polar; a carbon – oxygen bond or a carbon – nitrogen bond?

Q.44 Classify the following bond as ionic, polar covalent or nonpolar covalent (O, atomic number 8; F, atomic number 9; Na, atomic number 11; Cl atomic number 17; Ca, atomic number 20; U, atomic number 92):

O with F

Ca with Cl

U with Cl


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Q.45 List the following bonds in order to increasing polarity:

N – N, N – F, N – O, H – F,

(least polar) (most polar)______ , _______ , _______ , _______

Q.46 �e dipole moment of KCl is 293.336 10 cm. �e interatomic distance K and Cl ion in KCl is 260 pm. Calculate the dipole moments of KCl, if there were opposite charges of the fundamental unit located at each nucleus. Calculate the percentage ionic character of KCl.

Q.47 Arrange the bonds in order of increasing ionic character in the molecules: 2 2 2 3LiF, K O, N , SO and ClF

Q.48 Explain why 2BeH molecule has a zero dipole moment although the Be – H bonds are polar.

Q.49 Draw the lewis structures for the ionic compounds

(i) 2CaF (ii) 2K O

(iii) 2BaCl

Q.50 Dipole moment of HBr is 0.779 D and the inter atomic spacing is 1.41 Å. What is the percent ionic character of HBr?

Q.51 What will be the formula and electron dot structure of cyclopentane?

Q.52 Explain the nature of the covalent bond using the bond formation in 3CH Cl .

Q.53 Draw reasonating structures of

(i) 3N H (ii) 2HNO

(iii) CO

Q.54 Calculate the formal charge on

(i) S in 4HSO ion (ii) Cl in 4HClO

Long Answer Questions

Directions: Give answer in four to �ve sentences.

Q.55 (i) Draw Lewis structures for the following molecules: 3 2 4 2 3PH ,H S,SiCl ,BeF ,AlI ,HCOOH.

(ii) Draw the Lewis electron dot diagram for each of the following ion.2

3 3 3(a) SO (b) CO (c) CN (d) BrF

Q.56 (i) �e skeletal structure of 3CH COOH as shown below is correct, but some of the bonds are


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wrongly shown. Write the correct Lewis structure for acetic acid.

(ii) Write Lewis structure of the following compounds –

(a) BaO (b) MgCl2

(c) K2S

Q.57 (i) Which of the molecule, in each of the following pairs have zero dipole moment?

2 2 3 3 3 4 2(a) HF, F (b) N ,NF (c) BF ,CH Cl (d) CCl ,H O

(ii) Arrange the following in increasing order of their dipole moment.

3 3 3 3CH Cl, CH Br, CH F and CH I

Q.58 (i) Calculate the percentage ionic character of HCl. Given that the observed dipole moment is 1.03 D and bond length of HCl is 1.275 Å.

(ii) �e dipole moment of LiH is 291.964 10 cm and the interatomic distance between Li and H in the molecule is 1.596 Å. Calculate the percent ionic character of the molecule.

Q.59 You are given the electronic con�guration of �ve neutral atoms – A, B, C, D and E.

A - 2 2 6 21s 2s 2p 3s , B - 2 2 6 11s 2s 2p 3s , C - 2 2 11s 2s 2p , D - 2 2 51s 2s 2p , E - 2 2 61s 2s 2p

Write the empirical formula for the substance containing

(i) A and D (ii) B and D

(iii) Only D (iv) Only E


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EXERCISE 2 – For Competitive Examinations


Directions: �is section contains multiple choice questions. Each question has 4 choices (a), (b), (c) and (d) out of which ONLY ONE is correct.

Q.1 When two atoms combine to form a molecule

(a) Energy is released

(b) Energy is absorbed

(c) Energy is neither released nor absorbed.

(d) Energy may either released or absorbed.

Q.2 Which of the following compounds have highest melting point?

(a) BeCl2 (b) MgCl2 (c) Cal2 (d) SrCl2

Q.3 Which of the following is most covalent?

(a) AlF3 (b) AlCl3 (c) AlBr3 (d) AlI3

Q.4 �e molecule having a high dipole moment among the following is –

(a) H2S (b) BF3 (c) CO2 (d) CCl4

Q.5 Bond length of HCl is 1.275 Å 10(q 4.8 10 e.s.u.) if 1.02D , then HCl is

(a) 100 % ionic (b) 83% covalent

(c) 50% covalent (d) 40% ionic

Q.6 Which of the following bonds is most polar?

(a) O – H (b) P – H (c) C – F (d) S – Cl

Q.7 Element X is strongly electropositive and element Y is strongly electronegative. Both are univalent. �e compound formed would be

(a) X Y (b) X X

(c) X Y (d) X Y

Q.8 �e total number of electrons that take part in forming the bond in N2 is

(a) 2 (b) 4 (c) 6 (d) 10

Q.9 Number of paired electrons in O2 molecule is

(a) 7 (b) 8 (c) 16 (d) 14


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Q.10 Bonding in ferric chloride is

(a) Covalent (b) Ionic

(c) Co-ordinate (d) None of these

Q.11 �e correct order of decreasing polarizable ions is

(a) Cl Br I F (b) F I Br Cl

(c) F Cl Br I (d) I Br Cl F

Q.12 Consider the following statements:

(i) A sigma (σ ) bond is formed when two s-orbitals overlap.

(ii) A pi ( π ) bond is formed when two p-orbitals axially overlap.

(iii) A σ -bond is weaker than π -bond.

Which of the above statements is/are correct?

(a) i and ii (b) ii and iii

(c) i alone (d) ii alone

Q.13 �e values of electronegativity of atoms A and B are 1.20 and 4.0 respectively. �e percentage of ionic character of A – B bond is

(a) 50% (b) 72.24 %

(c) 55.3 % (d) 43%

Q.14 �e electronegativities of F, Cl, Br and I are 4.0, 3.0, 2.8, 2.5 respectively. �e hydrogen halide with a high percentage of ionic character is

(a) HF (b) HCl

(c) HBr (d) HI

Q.15 A lone pair of electrons in an atom implies

(a) A pair of valence electrons

(b) A pair of electrons

(c) A pair of electrons involved in bonding

(d) A pair of valence electrons not involved in bonding

Q.16 Among LiCl, BeCl2, BCl3 and CCl4 the covalent bond character follows the order

(a) 2 3 4LiCl BeCl BCl CCl (b) 2 3 4BeCl BCl CCl LiCl

(c) 2 3 4LiCl BeCl BCl CCl (d) 2 3 4LiCl BeCl BCl CCl


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Q.17 �e electronegativity di�erence between N and F is greater than that between N and H yet the dipole moment of NH3 (1.5 D) is larger than that of NF3 (0.2 D). �is is because

(a) In NH3 the atomic dipole and bond dipole are in same direction whereas in NF3 these are in opposite directions.

(b) In NH3 as well as in NF3 the atomic dipole and bond dipole are in opposite directions.

(c) In NH3 the atomic dipole and bond dipole are in the opposite directions whereas in NF3 these are in the same direction.

(d) In NH3 as well as in NF3 the atomic dipole and bond dipole are in the same direction.

Q.18 Which of the following compounds has electrovalent linkage?

(a) CH3Cl (b) NaCl (c) CH4 (d) Cl2

Q.19 Elements whose electronegativities are 1.2 and 3.0. Bond formed between them would be

(a) Ionic (b) Covalent

(c) Co-ordinate (d) Metallic

Q.20 In a double bond connecting two atoms there is a sharing of

(a) 2 electrons (b) 4 electrons

(c) 1 electrons (d) All electrons

Q.21 Outermost shells of two elements X and Y have two and six electrons respectively. If they combine, expected formula of compound will be

(a) X Y (b) X2Y (c) X2Y3 (d) XY2

Q.22 An atom of an element A has three electrons in its outermost shell and that of B has six electrons in the outermost shell. �e formula of the compound between these two will be:

(a) A3.B4 (b) A2B3 (c) A3B2 (D) A2B

Q.23 Hydrogen chloride molecule contains

(a) Polar covalent bond (b) Metallic bond

(c) Co-ordinate bond (d) Electrovalent bond

Q.24 �e molecular formula of chloride of a metal M is MCl3 the formula of its carbonate would be

(a) MCO3 (b) M2(CO3)3

(c) M2CO3 (d) M(CO3)2

Q.25 �e rule which states that the atoms tend to form bonds until each atom has eight electrons in its outermost shell is called

(a) Pauli’s exclusion principle (b) Newland’s rule

(c) Octet rule (d) Dipole rule


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Q.26 �e dipole moment can be expressed as the

(a) Distance between the two atoms

(b) Electrostatic charges between the atoms

(c) Product of charge and distance between the atoms

(d) None of the above

Q.27 For non-polar diatomic molecules, the dipole moment is always

(a) Zero (b) One

(c) Two (d) Negative one

Q.28 �e most favourable conditions for the ionic bonding to occur are

(a) High charge, large anions and small cation

(b) High charge, small anion and large cation

(c) Low charge, small anion and large cation

(d) Low charge, large anion and small cation

Q.29 �e sodium chloride in the molten state conducts electricity because of the presence of

(a) Free molecules (b) Free electrons

(c) Free ions of sodium and chloride (d) Free atoms of sodium and chloride

Q.30 What causes the atom of two di�erent elements or same elements to go for bonding?

(a) �e opposite charges they carry

(b) �ere tendency to acquire a stable inert gas structure

(c) �e di�erent in their ionization energies

(d) None of the above

Q.31 Which of the following structure best describe the bond formation between carbon and oxygen to form a molecule of carbon dioxide?

Q.32 �e tendency of an atom of an element to attract the shared pair of electrons to form a covalent compound is called

(a) Covalency (b) Electronegativity

(c) Electron a�nity (d) Electron attractive force


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More than One Correct

Directions: �is section contains multiple choice questions. Each question has 4 choices (a), (b), (c) and (d) out of which ONE OR MORE may be correct.

Q.33 Valency expresses

(a) Total electrons in valence shell of an atom

(b) Atomicity of an element

(c) Oxidation number of an element

(d) Combining capacity of an element

Q.34 Which of the following compounds is/are covalent?

(a) H2 (b) CaO (c) CHCl3 (d) BI3

Q.35 Polarization power of a cation increases when

(a) Charge on the cation increases

(b) Size of the cation increases

(c) Cation acquire pseudo inert gas con�guration

(d) Has no relation to its size or charge

Q.36 Which statement is/are correct?

(a) A sigma bond is weaker than a π-bond.

(b) A sigma bond is stronger than a pi-bond.

(c) A double bond is stronger than a single bond.

(d) A double bond is shorter than a single bond.

Q.37 Which of the following statements is/are correct for sigma and pi-bonds formed between two carbon atoms?

(a) Sigma-bond determines the direction between carbon atoms but a pi-bond has no primary e�ect in this regard

(b) Sigma-bond is stronger than a pi-bond

(c) Bond energies of sigma- and pi-bonds are of the order of 264 kJ/mol and 347 kJ/mol, respectively

(d) Free rotation of atoms about a sigma-bond is allowed but not in case of a pi-bond

Q.38 Covalent compounds are soluble in

(a) Polar solvents (b) Non-polar solvents

(c) Slightly polar solvents (d) All types of solvents


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Q.39 As compared to covalent compounds, electrovalent compounds generally have

(a) Low melting points (b) Low boiling points

(c) High boiling points (d) High melting points

Q.40 �e factors which favours the formation of electrovalent bond is

(a) High electron a�nity (b) High lattice energy

(c) Low ionization potential (d) High ionization potential

Q.41 Which of the following is/are an electrovalent (ionic) compound(s)?

(a) Calcium chloride (CaCl2) (b) Sodium sulphate (Na2S)

(c) Magnesium �uoride (MgF2) (d) Carbon tetrachloride (CCl4)

Q.42 Which gas(es) has a triple bond in between its atoms?

(a) Hydrogen (b) Nitrogen (c) Oxygen (d) Acetylene

Q.43 Which of the statement(s) below is/are incorrect?

(a) �e covalent compound exist as ions.

(b) �e electrovalent compounds exist as molecules

(c) �e covalent compound exist as molecule not ions

(d) �e covalent compounds exist both as ions and molecules.

Q.44 Which of the following is/are exists as monoatomic?

(a) Helium (b) Fluorine (c) Radon (d) Oxygen

Fill in the Blanks

Directions: Complete the following statements with an appropriate word/term to be �lled in the passage.

Q.45 Majorly there are two types of bonds (1) and (2) . �e bond formed by transfer of electrons from one atom to other is (3) whereas bonds formed by mutual sharing of pair of electrons between two atoms is called covalent bond. Covalent bonds are (4) and compounds containing these bonds are (5)

conductors of electricity. Ionic bonds are (6) and compounds containing these bonds are (7) conductors of electricity in (8)

state.

Q.46 Electronegativity is highest for elements at the (1) of the periodic table and lowest for the elements at the (2) . On the basis of (3) di�erence between bonded atoms the covalent bond is classi�ed as polar and non-polar bond. �us bond formed between lithium and �uorine is (4) and between aluminium and chlorine is (5) .


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Passage Based Questions

Directions: Study the given passage(s) and answer the following questions.

When anions and cations approach each other, the valence shell of anions are pulled towards a cation, is known as polarisation and ability of the cation to polarize the anion is called as polarising power of cation. Due to polarisation, sharing of electrons occurs between two ions to some extent and bond shows some covalent character. Fajan’s suggested following factors on which polarisation depend.

(i) As the charge on cation or anion increases polarisation increases.

(ii) Size of cation decreases or size of anion increases, polarisation increases.

(iii) Cation with pseudo noble gas con�guration shows highest polarisation power.

Q.47 Which is most covalent in nature?

(a) NaCl (b) MgCl2 (c) AlCl3 (d) CaCl2

Q.48 Which is having highest melting point?

(a) LiF (b) LiCl (c) LiBr (d) Lil

Q.49 Highest polarisation is shown in which of the following compounds?

(a) MgCl2 (b) BaCl2 (c) AgCl (d) AgI


Directions: Each of these questions contains an assertion followed by reason. Read them carefully and answer the questions on the basis of following options. You have to select from following the one that best describes the two statements.

(a) If both assertion and reason are correct and reason is the correct explanation of assertion.

(b) If both assertion and reason are correct, but reason is not the correct explanation of assertion.

(c) If assertion is correct but reason is incorrect.

(d) If assertion is incorrect but reason is correct.

Q.50 Assertion: Boron always forms covalent bond.

Reason: �e small size of B3+ favours formation of covalent bond.

Q.51 Assertion: LiCl is covalent whereas NaCl is ionic.

Reason: Greater size of the cation, greater is its polarising power.

Q.52 Assertion: �e dipole moment helps to predict whether a molecule is polar or non-polar.

Reason: �e dipole moment helps to predict the geometry of molecules.


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Q.53 Assertion: According to Fajan’s rule, covalent character is favoured by small cation and small anion.

Reason: �e magnitude of covalent character in the ionic bond depends upon the extent of polarisation.

Q.54 Assertion: Dipole moment of BeF2 is not de�nite.

Reason: Structure of BeF2 is certain.

Q.55 Assertion: �e electronic structure of O3 is

Reason: Structure is not allowed because octet around O cannot be expanded.

Q.56 Assertion: LiCl is predominantly a covalent compound.

Reason: Electronegativity di�erence between Li and Cl is too small.

Multiple Matching Questions

Directions: Each question contains statements given in two columns which have to be matched. Statements (A, B, C, D) in column I have to be matched with statements (p, q, r, s) in Column II

Q.57

Column I Column II

(A) HCl (p) Covalent compounds with directional bond

(B) CO2 (q) Ionic compound with nondirectional bonds.

(C) NaCl (r) Polar molecule

(D) CCl4 (s) Non-polar molecule

Subjective Questions

Directions: Answer the following questions.

Q.58 List these bonds in order of increasing polarity: P – F, S – F, Ga — F, Ge—F (F, �uorine, atomic number 9; P, phosphorus, atomic number 15; 5, sulfur, atomic number 16; Ga, gallium, atomic number 31; Ge, germanium, atomic number 32):

(least polar) ________, __________, ________, _________ (most polar)

Q.59 Is an ionic compound an example of a chemical compound, or is a chemical compound an example of an ionic compound?

because octet around O cannot be expanded.


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Q. 60 What is an ionic crystal?

Q.61 Which element in the periodic table has the greatest electronegativity? Which has the least electronegativity?

Q.62 How is a polar covalent bond similar to an ionic bond?

Q.63 How can a molecule be nonpolar when it consists of atoms that have di�erent electronegativities?

Q.64 Atoms of non-metallic elements form covalent bonds, but they can also form ionic bonds. How is this possible?

Q.65 Which bond is most polar? H—N, N—C, C-O, C—C, O—H, C—H.

Q.66 Which molecule is most polar? S = C = S, O = C = O, O = C = S

Q.67 �e ionic bond is called a non-directional but covalent bond is called a directional bond. Why?

Q.68 Give the lewis structures and empirical formula for the ionic compounds formed from the following pair of elements.

(i) Al, Cl

(ii) Mg, P

(iv) Na, S

Q.69 Why an ionic bond is formed between two elements having large di�erence in their electronegativity?

Q.70 Why is HCl predominantly covalent in the gaseous state but ionic in the aqueous solution?


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SOLUTIONS

Exercise 1 – For School Examinations

Fill in the Blanks

1. Covalent bonds 2. Ionic 3. Bad

4. Shortens

5. 2; 2N N(N ) has 1σ and 2π bonds. (A triple bond consists of 1σ and 2π bonds)

6. Non-polar or covalent 7. 2N or 2O 8. Oxygen

9. Endothermic 10. Exothermic

True / False

11. True 12. True 13. True 14. False

15. False : Symmetrical molecules with polar bonds have zero dipole moment.

16. True 17. False 18. True 19. False

20. True


21. A – (r), B – (s), C – (p) , D – (q)

Very Short Answer Questions

22. In 2CaCl , the charge on calcium ion is 2+ (i.e., 2Ca )

23. An atom of oxygen can attract two valence electrons from other atoms and acquire near by stable gas con�guration.

24. An atom of oxygen can form two covalent bonds.

25. A dipole refers to separation of charge that takes place in a chemical bond due to di�erence in electro negativities of the bonded atoms.

26. Calcium atom (atomic number = 20) has two valence electrons (2, 8, 8, 2) and it has a tendency to lose electrons to form 2Ca .


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27. Due to their big size they are not able to interact with solvent molecule. �ey do not easily

undergo solvation as extent of solvation ChargeSize

28. Because reaction between covalent compounds involves breaking of covalent bonds which requires energy.

29. 3 3 3 3PH AsH SbH NH

30.

31.

32. Valence electrons present in the outermost shell.

33. Half-�lled atomic orbitals containing electrons with opposite spin.

34. s-s, s-p, p-p for σ - bond and only p-p for π - bond.

35. �e overlapping vanishes and the bond breaks.

36. Lowering of energy takes place.

37. Electrostatic forces of attraction.

Short Answer Questions

38. Valence electrons are located in the outermost occupied shell of an atom. �ey are important because they play a leading role in determining the chemical properties of the atom.

39. As magnesium has con�guration of (2, 8, 2), you know a magnesium atom must lose two electrons to form a 2Mg ion which has stable neon con�guration (2, 8). As oxygen has con�guration of (2, 8), an oxygen atom gains two electrons to form an 2O ion which has stable neon con�guration (2, 8). �ese charges balance in a one-to-one ratio, and so the formula for magnesium oxide is MgO.


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40. In the valence shell of oxygen there are six electrons. Atomic number of oxygen is 8 and thus an atom of oxygen has eight electrons. Out of these eight electrons two are present in shell one and six in shell two. In this case electrons in shell two (which is outermost shell) represent valence electrons. Oxygen atom has 2 non-bonding pairs (i.e. four electrons) and 2 unpaired electrons (i.e. 2 electrons).

41. Fluorine atom (atomic number = 9) has seven valence electrons (2, 7). In this case second shell is outermost shell and it can have 8 electrons. By gaining one electron the outermost shell gets 8 electrons and become stable. Due to this �uorine atom has a tendency to gain only one electron and forms F- which has a stable arrangement of electrons (2, 8).

42. �e force that holds two atoms together in a covalent bond is the electrical force of attraction that two atoms have for the shared pair of electrons.

43. More the di�erence in electronegativity values of the atoms forming a covalent bond more is polarity of bond. In case of carbon – oxygen bond the electronegativity di�erence is 0.89(3.44 – 2.55= 0.89) and in case of carbon – nitrogen bond the electronegativity di�erence is 0.49 (3.04 – 2.55 = 0.49). Since electronegativity di�erence in case of carbon - oxygen is more than that in case of carbon – nitrogen so carbon – oxygen bond is more polar.

44. �e type of bond is shown against each. Bond of O with F is polar covalent. Bond of Ca with Cl is ionic. Bond of U with Cl is polar covalent.

45. Greater the di�erence between electronegativities of elements higher polar will be the molecule. �e following is the order of increasing polarity. Electronegativity di�erence

0.940(least polar) 0.40 1.68(most polar)N N N O N F H F

46. From the given data, q = 1.602 x 10-19C,

r = 260 pm 12 10 260 x 10 m = 2.6x10 m .

Magnitude of dipole moment for 100% ionic character 19 10 29| | qr (1.602 10 )(2.6 10 ) 4.165 10 Cm

Actual dipole moment = 3.336 x 10-29 Cm29

29

3.336 10% of ionic bond 100 80.1%

4.165 10.

�e bond is 80.1% ionic.

47.

Molecule LiF K2O N2 SO2 CIF3

Di�. in the electronegativities of the two atoms 3.0 2.7 0 1.0 1.0

So, the ionic character in the bonds follows the order, 2 3 2 2N ClF SO K O LiF .


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48. �is is because BeH2 is a linear molecule and the two bond dipoles oriented at an angle of 180° neutralize each other, As a result, the net dipole moment of the molecule is zero.

49.

50. Dipole moment of a 100% ionic molecule at the given internuclear distance would be e x d

i.e. Dipole moment 10 84.8 x10 esu x 1.41 x10 cm. 186.768 x 10 esu. cm or 6.768 D

obs

cal

0.779 D% ionic character 100 100 11.50%

6.786 D

51.


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52. Chloromethane is made up of one carbon atom, three hydrogen atoms and one chlorine atom. Carbon atom has,4 outermost electrons, each hydrogen atoms has 1 outermost electron and chlorine has 7 outermost electrons. Carbon shares its four outermost electrons with three hydrogen atoms and chlorine atom to form CH3Cl as follows:

Each atom in CH3Cl has a noble gas electronic con�guration. Carbon attain as the nearest noble gas con�guration of neon, hydrogen acquires the con�guration of helium while chlorine achieves the con�guration of argon. Chloromethane contains three C - H and one C - Cl covalent bond.

53. (i) Hydrazoic acid (N3H)

(ii) Nitrous acid (HNO2)

(iii) Carbon monoxide(CO)

54. (i) Lewis structure of 4HSO ion is

Applying the formula for calculation of formal charge. Formal charge on

1S 6 0 (8) 6 4 2

2

(ii) Lewis structure of HClO4 is

Formal charge on 1Cl 7 0 (8) 3

2


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Long Answer Questions

55. (i) �e Lewis structures of the given molecules or ions are shown below:

(ii)

56. (i) �e correct Lewis structure for acetic acid is

(ii) (a) Ba = [Xe] 6s2 : two electrons excess to inert gas con�guration; can lose two electrons.

O = 1s2 2s2 2p4; two electrons less than octet; can gain two electrons


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(b) Mg = [Ne] 3s2 : two electrons excess to inert gas con�guration; can lose two electrons.

CI = [Ne] 3s2 3p5 : one electron less than octet; two chlorine atoms gain two electrons which are lost by magnesium.

(c)

57. (i) (a) HF is unsymmetrical molecule and thus has some value of dipole moment whereas F2 is a diatomic molecule having both F atoms. �erefore, its dipole moment is zero.

(b) N2 is Linear molecule have zero dipole moment whereas NF3 is having trigonal pyramidal geometry having some dipole moment.

(c) BF3 is planar molecule and thus has zero dipole moment. CH3Cl is tetrahedral in shape.

(d) CCl4 is tetrahedral in geometry but cancel out each others e�ect. �us, dipole moment is zero. On the other hand H2O molecule has bent shape with some value of dipole moment.

(ii) Greater the electronegativity of halogen more is its dipole moment. �erefore, increasing order of dipole moment is CH3I < CH3Br < CH3C1 < CH3F

58. (i) If HCl were 100 % ionic, each end would carry charge equal to one unit viz, 104.8 10 esu. As bond length of HCl is

01.275A , its dipole moment for 100% ionic character

would be10 8

ionic q d 4.8 10 esu 1.275 10 cm. 186.12 10 esu 6.12D

observed 1.03D(Given)

observed

ionic

%ionic character 100 16.83%

(ii) If the molecule were 100 % ionic,

19 10 29ionic q d (1.602 10 C) (1.596 10 m) 2.557 10 Cm


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29observed

29ionic

1.964 10 Cm%ionic character 100 100 76.81%

2.557 10 Cm

59. (i) Empirical formula of the compound formed by A and D is AD2 as A has two valence electrons and D has seven. Atom A transfers its two electrons to two D atoms to complete their octets.

(ii) Empirical formula of the compound formed by 13 and D is BD as B transfers its one electron to D.

(iii) D2 as both the atoms of D share one electron each to form a covalent bond.

(iv) Since it is a noble gas, no compound is formed.

Exercise 2 – For Competitive Examinations


1. (a). To attain stability energy is released.

2. (d). As the size of cation increases, polarizing power decreases hence ionic character increases.

3. (d). As the electronegativity di�erence decreases, covalent character increases.

4. (a). It is an unsymmetrical molecule with the bond angle H – S – H slightly greater than 90 .

5. (b).

observed 1.02% ionic character 100 100 17%

experimental 1.275 4.8

ionic = 83% covalent

6. (c). Due to maximum electronegativity di�erence.

7. (a). X Y

8. (c). N N

9. (d). O2 – Oxygen (Z=8) has following electronic con�guration 1s22s22p63s23p4. i.e., 2 unpaired and 14 paired electrons.

10. (b). Covalent bond on the basis of Fajan’s rule.

11. (d). �e larger the size of anion the more is its polarizability.

12. (c). Statement 1 is correct.


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13. (b). Electronegativity di�erence is 4.0 - 1.20 = 2.8 percentage ionic character is 72.24% when the electronegativity di�erence is 1.7, the % ionic character is approx. 51%.

14. (a). Ionic character follows the order

HF > HCl > HBr > HI

15. (d). Lone pair of valence electrons are not involved in bonding.

16. (c). As we move from Li Be B C , the electronegativity (EN) increases and hence the EN di�erence between the element and Cl decreases and accordingly the covalent character increases �us option (c) i.e., LiCl < BeCl2 < BCl3 < CCl4 is correct.

17. (a). In NH3 the atomic dipole and bond dipole are in the same direction whereas in NF3 these are in opposite direction so in the former case they are added up whereas in the latter case net result is reduction of dipole moment. It has been shown in the following �gure :

18. (b). Out of given four compounds, only NaCl has electrovalent linkage and others have covalent linkages Na has 1 e- in its outermost shell thus tends to loose it and Cl has 7e- in its outermost shell thus tends to gain an e-, hence the bond that is formed between Na and Cl is formed by transfer of electron i.e, electrovalent bond.

19. (a). When the electronegativity di�erence is 1.7, the bond existing between two atoms is 50% ionic and 50% covalent. If the electronegativity di�erence is more than 1.7, the chemical bond is predominately ionic whereas if the di�erence is less than 1.7, the bond is mainly covalent.

Two pairs of electrons. e.g. in C2H4.

20. (b). Two pairs of electrons. Eg. In C2H4 .

21. (a).

Hence, the formula is XY

22. (b). Clearly, valency of element A = 3. valency of element B = 8 - 6 = 2 Hence, the compound

2A3+ + 3B2- A2B3


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23. (a). A gaseous HCl molecule has hydrogen and chlorine linked by a covalent bond. Here electronegativity of chlorine is greater than that of hydrogen. Due to this the shared pair of electron is more attracted towards chlorine. �us, chlorine end of molecule has higher electron density and becomes slightly negative and the hydrogen becomes slightly positive. Hence the covalent bond in HCl has a polar character as shown below

H Clδ δ

24. (b). 33MCl M 3Cl ; 3 2

3 2 3 32M 3CO M (CO )

25. (c). According to octet rule, all atoms in the formula of compound will have a total of eight electrons by sharing in the valence shell.

26. (c). Dipole moment ( ) = charge (q) distance between two atoms (d)

27. (a). For example, dipole moment of O2 = zero

28. (b).

29. (c). Due to presence of Na+ and Cl ions, NaCl conducts electricity in molten and aqueous state.

30. (b). Self-explanatory



More than One Correct

33. (a, d). It is basic de�nition.

34. (a, c, d). All these compounds are covalent in nature.

35. (a,c) Polarizing power (P = charge/radius.

36. (b, c, d). A σ - bond is stronger than a π -bond hence option (a) is not correct. Sigma (σ ) bonds are formed by head on overlap of s-s, p-p or s-p orbitals hence a bonds are strong bonds. Whereas pi ( π )-bonds are formed by side ways overlap of p and d orbitals hence π bonds are weak bonds.

37. (a, b, d). As sigma bond is stronger than the π (pi) bond, so it must be having higher bond energy than π (pi) bond.

38. (a, b, c). Covalent compounds are both polar and non polar.

39. (c, d). Ionic compounds generally have high melting and boiling points because of the strong electrostatic force of attraction between oppositely charged ions. Consequently, a considerable amount of energy is required to overcome strong attractive inter-ionic forces and to break down the crystal lattice.

40. (a,b,c).


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41. (a, b, c). Only CCl4 is covalent in nature as electronegativity di�erence between carbon and chlorine is not high.

42. (b, d) N N and

H C C H

both has triple bond in between its atoms.

43. (c). Example 34PO is a covalent and is an ion also.

44. (a,c). He and Xe exists as a monoatomic gas.

Fill in the Blanks

45. (1) Ionic (2) Covalent (3) Ionic bond (4) Directional

(5) Bad (6) Non-directional (7) Good (8) Molten

46. (1) Upper right (2) Lower left (3) Electronegativity (4) Ionic

(5) Covalent

Passage Based Questions

47. (c). Higher charge density on Al leads to higher polarising power of Al3+ cation. So AlCl3 is most covalent in nature.

48. (a). LiF is most ionic (F–) is small in size therefore it has least polarisability). �erefore, it has higher melting point.

49. (d). AgI shows highest polarisation due to following two factors (i)Ag+ has pseudo noble gas con�guration (ii) I- is biggest anion of halogen family.


50. (a). Smaller the size of cation more will be polarisation.

51. (c). Smaller the size of cation, greater is its polarising power.

52. (a).

53. (d). Covalent character is favoured by small cation and larger anion.

54. (d). Dipole moment of BeF2 is zero. Structure of BeF2 is linear.

55. (a). Both assertion and reason are correct. �e reason explains the assertion as the central O-atom cannot have more than 8 electrons (octet).

56. (c). LiCl is a covalent compound. Due to the large size of the anion (Cl) its e�ective nuclear charge lessens and its valence shell is held less tightly towards its nucleus. Here, assertion is correct but reason is incorrect.


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Multiple Matching Questions

57. A – (p,r), B – (p,s) C- (q), D – (p,s)

Subjective Questions

58. �e greater the di�erence in electronegativites between two bonded atoms, the greater the polarity of the bond, and so the order of increasing polarity is S-F < P-F < Ge-F < Ga-F.

Note that this answer can be obtained by looking only at the relative positions of these elements in the periodic table rather than by calculating the di�erences in their electronegativities.

59. An ionic compound is an example of a chemical compound. A chemical compound may be an ionic compound, covalent compound or a coordinate compound and not necessarily an ionic compound.

60. �e orderly arrangement of ions in case of ionic compounds is called an ionic crystal. For example, In sodium chloride crystal each sodium ion is surrounded by six chloride and each chloride ion is surrounded by six sodium ion.

61. Electronegativity is greatest for elements at the upper right of the periodic table.(except inert gases) and lowest for the elements at the lower left, �erefore, in the periodic table the electronegativity is maximum for �uorine, F (3.90) and is minimum for francium, Fr (0.7).

62. In a polar covalent bond, the centers of partial positive charge (δ+) and partial negative charge (δ-) are separated whereas in an ionic bond the electron is completely transferred from one atom to other. In both type of bonds the atoms are held together by the electrostatic force of attraction.

63. A molecule consisting of atoms of di�erent elements having di�erent values of electronegativity can still be non-polar, if there is no net dipole in a molecule. For example, CO2 is a non-polar molecule because of its linear structure. In CO2 two C = O bonds are oriented in the opposite directions at an angle of 1800. �us, due to linear geometry dipole moment of one C = O bond cancels that of another.

64. �e atoms of non-metallic elements are short of a few electrons in their valence shell and they can complete their outermost shell by sharing electrons (one or more) amongst themselves and so they can form covalent bond e.g. F – F , O = O, N N

�e atoms of non-metallic elements can also complete their outermost shell by gain of electrons and in such cases, they form ionic bond e.g. NaCl.

65. Of the following bonds, the most polar is one that has maximum ‘di�erence in electro negativities of participating atoms.

�e di�erence is maximum in case of O - H (3.44 – 2.2 = 1.24) because O and H have maximum separation in periodic table as compared to H – N (3.04 - 2.2 - 0.84), N – C (3.04 - 2.55 = 0.49), C – O (3.44 – 2.55 = 0.89) or C – C (0).


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66. Out of the given molecules, the dipole cancels out in case of S = C = S and O = C - O (due to symmetry) and so the most polar molecule is O = C = S which has some resultant polarity.

67. Because in ionic bond, an ion can attract other opposite charged ions from any direction and extends equally in all directions. So the nature of ionic bond is non-directional, but on the other hand a covalent bond is formed by the proper overlapping of orbitals so it have directional character.

68. (i) 3

2,8,3 2,82,8,7 2,8,8Al , Cl Al , Cl (ions)

3

3Lewis formula Al 3Cl or AlCl

(ii) 2 3

2,8,5 2,8,82,8,2 2,8Mg, P Mg ,P (ions)

2 3

3 2Lewis formula 3Mg 2P Mg P

(iii) 2

2,8,1 2,8,6 2,8 2,8,8Na , S Na ,S (ions)

22Lewis formula 2Na S Na S

69. �e more electronegative element will attract the shared pair of electrons to such a large extent than the other that it will amount to transfer of electron resulting in the formation of ions.

70. �e electronegativity di�erence between Cl and H atoms is 3.0 – 2.1 = 0.9. Hence, HCl is

predominantly covalent in the gaseous state. However, being a polar molecule (H Cl),δ δ

when dissolved in water, the polar H2O molecules interact with HCl molecule as follows :

As a result, the bond between H and Cl is broken and we get hydrated H+ and Cl- ions in the solution.

3. CHEMICAL BONDING - TopperLearning

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