• There are about 109 different kinds of elements which constitute all matter.
• Except noble gases, no other element exists as independent atoms under ordinary conditions, all of them exist in combined state.
• Only the atoms of noble gases have independent existence.
• According to Dalton’s Atomic theory, atoms combine together to form compound atoms or combined atoms.
• The term compound atom is proved to be incorrect, and is replaced by the term molecule. (Molecule term is proposed by Avogadro).
• The number of atoms present in a molecule is called as Atomicity.
• Molecule consisiting of atoms of one kind of elements is called as Homo atomic molecule.
Ex: H2, N2, O2, O3, F2, Cl2, Br2, I2, P4, As4, Sb4, S8, Se8, Te8, ……. etc.
• Molecules consisting of atoms of different element are called as Hetero atomic molecules.
Ex. H2O, CO2, CH4, H2SO2, C6H12O6 …..etc.
Valency : Combining capacity of an elements is called as its valency
• Valency of the elements can be explained in terms of Hydrogen, Chlorine and Oxygen.
• Valency of elements with respect to Hydrogen is the most common valency.
• Valency of elements with respect to ‘H’ remains the same in groups while in periods from left to right it increases from 1 to 4 and decreases to 0 and with respect to oxygen it is equal to group number.
• Valency with respect to Chlorine is same as that of Hydrogen in groups and periods.
• Valency is defined as no of the Hydrogen atoms or no.of chlorine atoms or double the no. of oxygen atoms that can combine with atoms of an element
• Those electrons of an atom, which decides the valence of that element are called as valency electrons.
• The orbit to which those valence electrons belongs is called as valency orbit.
• For representative elements, electrons of nth orbit (outer most orbit) are the valency electrons.
• For transition elements, electrons belonging to both ns and (n-1)d are valence electrons.
• For inner transition elements. ns, (n-1)d and (n-2)f electrons are valence electrons.
Chemical Bond: The forces of attractions which keep atoms or Ions together to form Molecules or Compounds is called Chemical Bond.
• Energy changes or the rearrangement of electrons are primarily responsible for the bond formation.
• The energy released during the formation of chemical bonds in any substance is sufficient to rearrange the electrons of valency orbits of bonded atoms.
• The process of bonding is accompanied by decrease in energy.
• The greater the decrease in Energy in the formation of a bond, greater is the strength of the bond.
• Energy of a molecule is always less than the total energy of all the atoms in it.
• Formation of chemical bond is always an exothermic process, i.e., when a molecule is formed from the isolated atoms energy is always released.
Ex. H + H → H2+104 K.Cal/Mole
[ΔH = –434.72 KJ/Mole]
• During formation of a molecule from the atoms, their Potential Energy is decreased.

• In 1916 Kossel and Lewis independently explain the formation of chemical bond interms of electrons.
• This is logically explained by “valence concept”. (which was based on the inertness of noble gases).
• Lewis pictured the atom in terms of a “positively charged kernel” and the valence shell that accommodate a maximum of eight electrons. (these 8 electrons assumed to be occupy the corners of a cube which surround the kernel)
Ex: i) one valence shell of sodium, occupy one corner of the cube
ii) In Noble gas, 8 electrons occupy eight corners of the cube.
• Lewis postulated that “atoms achieve the stable octet by transfer of electrons (or) sharing of a pair of electrons between the atoms.

LEWIS SYMBOLS: Kernel = Nucleus + inner electrons:
• In the formation of a molecule, only outer shell electrons take part in chemical reaction called as “valence electrons”.
• The inner shell electrons (core electrons) not involved in chemical reaction.
• G.N. Lewis introduced simple notations to represents valence electrons in an atom are called “Lewis symbols”.
Significance of Lewis symbols: The number of dots around the symbol represents number of valence electrons, useful to calculate common (or) group valance of elements.
• Valence electrons = Group number of an element (or) group valence of an element = number of dots in Lewis symbol (or) 8 – no. of valence electrons.
Ex: Lewis symbols for second period elements

KOSSEL in relation to Chemical Bonding :
drew attention to the following falts :
• Periodic table contains highly electronegative.
halogens and highly electropositive alkalimetals.
• Halogens form a negative ion by gain of electron.
Ex : Cl + e → Cl
[Ne] 3s23p5      [Ne] 3s23p6 (or) [Ar]
• Alkalimetals form a positive ion by loss of electron.
Ex : Na → Na+ + e
[Ne]3s1         [Ne]
• The negative and positive ions attains noble gas configurations & gets more stability.
• Noble gases have a stable octet electronic configuration i.e., ns2np6 (or) 8 valency electrons [except :- He has only 2e i.e., 1s2]
• The negative and positive ions are brought together by strong electrostatic force of attraction and the bond formed between them was termed as “Ionic bond” (or) “Electrovalent bond”.
Ex :- Ca → Ca+2 + 2e                           O + 2e → O-2
[Ar]4s2     [Ar]                                 [he]2s22p4        [Ne]
:.Ca+2 +O-2 → CaO

Electrovalency :- The number of electrons gained (or) lost by an atom in the formation of Ionic bond is termed as “Electrovalency”.
It was given by Kossel
Ex (1) : In the formation of NaCl
Electrovalency of Na, Cl are +1, –1
Ex (2) : In the formation of CaF2
Electrovalency of Ca, F are +2, –1

Merits :
• Kossel postulations useful to understanding the
(i) formation of ions by transfer of electrons.
(ii) formation and systematisation of ionic compounds.

Demerits :
• Kossel recognise that a large number of compounds did not fit into these concept.


• Term covalent bond introduced by “Langmuir” based on the Lewis postulations by a bonding the ideal of the stationary cubical arrangement of the octet.
Ex :- Lewis Langmuir theory can be understood by the formation of Cl2 molecule.
• Electronic configuration of Cl is [Ne]3s23p5

Each ‘Cl’ atom attains octet configuration by sharing of one electron to the shared pair.

Lewis Representation of simple molecules (or) Lewis dot structures.
• Lewis dot structures can be written by adopting the following steps.
i) Total number of electrons required for writing the structure are obtained by adding valence electrons of the combining atoms.
Ex :- For CH4 molecule. [ 4 valence e– from C & 4 from H]
Eight electrons available for bonding.
ii) For Anions & Cations :-
Number of electrons available for bonding
= Total valence electrons of all atoms ± charge.

Lewis structure and covalent bond :
Dots represent electrons
Electron dot structures, also known as Lewis structures of covalent molecules, are written in accordance with octet rule.
i) Each bond is formed by sharing of an electron pair between the atoms.
ii) Each combining atom contributes one electron to the shared pair.
iii) Structure in which valence electrons are represented by dots are called lewis structures.
iv) All atoms in a formula will have a total of eight electrons by sharing in the the valence shell (except the H-atom which forms the largest number of bonds with other atoms placed in the centre of skeleton structure. Other atoms surrounds it to complete the octet).
v) Structure represented by line (–) or dashes are known as couper structure.
vi) Lewis dot formulae show only the number of valency electron, the number and kinds of bonds, but do not depict the three dimensional shapes molecules and polyatomic ions.
Ex :- (i) For NH4+ ion :-
Number of electrons for bonding = 5+4–1=8e

Lewis dot structue :-

• In general the least electronegative atom taken as central atom.
Ex :- In NF3 central atom is ‘N’ and ‘F’ is terminal atom.
• In the Total number of electrons,
a) First account the shared pairs of electrons for single bonds, than the remaining electron pairs are either utilised for multiple bonding (or) lone pairs.
b) The basic requirement being that each bonded atom gets an octet of electrons. Provide the skeletal structure of the compound.
Note :- Lewis dot structure i) Provide a picture of bonding in molecules & ions.
ii) Do not explain the bonding and behaviour of a molecule completely.
iii) It does not help in understanding the formation and properties of a molecule.

Definition for Covalent Bond :
The bond formed by the pairing and mutual sharing of electron pairs of bonded atoms is
called as Covalent Bond.
• Formation of Covalent Bond was initially explained by Lewis.
• No. of Covalent Bonds formed between two atoms is equal to the No. of electron pairs shared between them.
• Single Bond is formed when one electron pair is shared between two bonded’ atoms (H2,F2,CI2, Br2, I2)
• A double covalent Bond is formed, when two electron pairs are shared between two bonded atoms Ex: O2
• When more than one bond is formed (double, Triple) between bonded, atoms, it is called as “Multiple bond”.
• Maximum number of bonds formed between any two atoms never exceeds three.
• Pure convalent bond is formed, when electron pair(s) are shared between atoms of same element with similar E.N.values.
Ex : H2, F2, Cl2, O2, N2, P4, S8 etc.,
Covalency:- The number of electrons an atom of the element contribute in the formation of a covalent compound is known as covalency of the element
a) The covalency of Oxygen in water = 2
b) The covalency of Hydrogen in water =1
c) The covalency of Nitrogen in NH3 = 3
d) The covalency of Nitrogen in NH4+ = 4
e) The covalency of Phosphorous in PCl3 = 3
f) The covalency of Phosphorous in PCl5 = 5
• BeCI2,BF3,BCI3 .. Molecules whose central atoms have contracted octet configuation (less than octet) are called as Electron Deficient Molecules
• PCl5, SF6, IF7, …Molecules whose central atoms have expanded octet configuation (more than octet)
BCI3,BF3 B has 6-electrons in its valency orbit PCl5 P has 10 electrons in its valency orbit
BeCI2 Be has 4 electrons in its valency orbit
SF6 S has 12 electrons in its valency orbit
IF7 I has 14 electrons in its valency orbit

Formal Charge :
• Formal charge is a factor based on a pure covalent bond formed by the sharing of electron pairs equally by neighboring atoms.
• Formal charge may be regarded as the charge that an atom in a molecule would have if all the atoms had the same electronegativity.
• It may or may not approximate the real ionic charge. In case of a polyatomic ions, the net charge is possessed.
• the real ion as a whole and not by an particular atom. It is, however, feasible to assign a formal charge on an atom in a polyatomic molecule or ion.
Formal charge [F.C.] on an atom in a Lewis structure= [no. of valence electrons in the free atom] – [no. of unshared electrons on the atom] – [no. of bonds around the atom]
Qf = [NA–NM]=[NA–NLP – 1/2NBP]
NA = number of electrons in the valence shell in the free atom
NB = number of electrons belonging to the atom in the molecule
NLP = number of electrons in unshared pairs, i.e. number of electrons in lone pairs
NBP = number of electrons in bond pairs, respectively.
Qf = formal charge
Generally the lowest energy structure is the one with smallest formal charge an atoms.
Ex-1 : PH3 molecule : The Lewis dot formula of PH3 is

Formal charge of P :
Qf= [NA–NM] = [NA–NLP –1/2NBP]
= {5 – 2 – 1/2(6)} = (5 – 5)=0
Ex-2: O3 molecule:

Formal charge of oxygen atoms marked for :
1) = 6 – 2 – 3 = +1
2) = 6 – 4 – 2 = 0
3) = 6 – 6 – 1 = –1
Therefore O3 is represented as


Octet Rule :-
Definition : Atoms prefer to get eight electrons in their outer most shell by losing (or) gaining (or) sharing electrons and attains stability is called “Octer Rule”.
Based on octer rule molecules are classified in to two types.
(i) molecules obey octet rule.
(ii) molecules not obey octet rule.

(i) Molecules obey octet rule :
The molecules (or) ions in which central atom contains 8e in its outer most shell obey octet rule & must be stable.
Ex :

ii) Molecules not obey octet rule:
• NO, NO2, CIO2 molecules are odd electron species, they are paramagnetic in nature.
• NO, NO2, CIO2, O2 molecules contains three electron bond.
• H21+ ion contains single electron bond.
• O2-1 superoxide also have three electron bond.

Limitations of the octet rule :
• Molecules not obey octet rule are placed in this theory.
• These are again classified in to two types.
(i) Central atom has incomplete octet (<8e)
• In these compounds, the no. of electrons surrounding the central atom is less than eight.
• These compounds are also called as “electron deficient compounds”.
• These compounds central atom may surrounded by 2, 4, 6 valence electrons.
Ex :- LiCl, BeH2, BeCl2, BF3, BCl3, BBr3, BI3, AlCl3,………
• These compounds generally acts as “Lewis acids”. (by accepting e– pair)
(ii) The expanded octet :
• In these compouns, the no. of electrons surrounding the central atom is more than eight are called as expanded octet configuration.
Ex :-

Drawbacks of octet theory :
• Octet Rule is based upon the chemical inertness of noble gases.
i) This theory does not considered the shape of the molecules.
ii) It does not explain the relative stability (or) energy of the molecules.

IONIC BOND or electrovalent bond : (Electrostatic Attraction between ions )
• The electrostatic forces of attraction, which keep oppositely charged ions together, that are formed by the transfer of electrons from one atom to the other is called as electrovalent bond or ionic bond.
• Formation of ionic bond is explained in three steps.
a) Transfer of electrons from valence shell of one atom to other.
b) Formation of oppositely charged ions.
c) Combination of oppositely charged ions.
d) For the formation ionic bond DEN is more than 1.7

• A metal atom loses electrons
• A nonmetal atom gains electrons
• Atom with low IE, high E.P. loses electrons.
• Atom with high EA, high EN. gains electrons.
• Metal atom is converted to cation.
• Non metal atom is converted to Anion.
• Metal atom gets Oxidises.
• Non metal atom gets reduces.
• Oxidation state of metal atom increases.
• Oxidation state of Non Metal atom decreases.

Conditions for forming ionic bonds :
• Formation of ionic bond depends upon following factors
a. Ionisation energy :
• Lesser ionisation energy → Greater tendency to form cation which further form ionic bond more
Ex: Na+ > Mg+2 > Al+3
Cs+ > Rb+ > K+ > Na+ > Li+
• Group 1, 2 elements have lower ionisation energy and form cation easily.

b. Electron affinity :
• Higher electron affinity Greater tendency to form anion which further form ionic bond more
Ex: Cl > F > Br > I
F > O–2 > N–3
• Group 16, 17 elements have high electron affinity and form anion easily.

c. Electronegativity :
• Ionic bond electro negativity difference between bonded atoms
• Ex: In generally, between metals of group 1, 2 and non-metals of groups 16, 17 favoured for the formation of ionic bond.
d. Lattice energy :
higher lattice energy Greater will be the stability or strength of ionic compound.
· Ionic compounds in the crystalline state consist of orderly three – dimensional arrangements of cations and anions held together by coloumbic attractions.
· Different crystal structures can be determined by
(i) Size of ions
(ii) Packing arrangement of ions & other factors
· Ionic crystal becomes stabilised when,
Electron gain enthalpy + Ionization enthalpy = +ve
but, crystal structure gets stabilized because, energy is liberated in the formation of the crystal lattice.
Ex :- Na(g) Na+(g) + e– ; I.P = 495.8 KJ.mole–1
Cl(g) + e– Cl–(g) ; EA =–348.7KJ mole–1
Na(g)+ Cl(g)Na+(g)+Cl–(g) ; DH = 147.1KJ.
· This value is much more than the enthalpy of lattice formation of NaCl(s) [ L.E = –788KJ mole–1]
· Lattice energy indirectly calculated by using Born-Haber cycle. [Depends on Hess law]

· All ionic compounds are solids at room temparature.
· They are hard and brittle in nature.
· Their Mpt and Bpt are very high due to strong electro static forces of attractions between ions.
· They are soluble in polar solvents like, H2O, Liq. NH3 …etc. (due to high dielectric constant values of these solvents).
· The energy required to break ionic solid in water is derived from hydration energy of ions.
· Aqueous solutions of ionic compounds have ionic – dipole interations between ions and water molecules.
· In aqueous solutions of ionic compounds ions get solvated.
· In molten and aqueous state they are good conductors of electricity due to the mobilities of free ions.
· Due to the non directional nature of ionic bond, ionic compounds does not exibit SPACE ISOMERISM.
· Electrovalent compounds undergo ionic reactions in solution. phase, which have very high reaction rate.
· Due to the non availability of discrete molecules, ionic compounds does not have molecular formula, but have only empirical formula.
· If anion is common in the ionic compounds, then a compound with high percentage of ionic nature posses high Mpt.
Ex: BaCI2, BeCI2
Ionic nature of BaCI2 > BeCI2.
· If cation is common then compound with high LE will have high Mpt.
Ex: LE of NaF > Nal
· Equation to calculate percentage of ionic character
% of I.C = x 100
· % of Ionic character in binary compound AB can be calculated with the help of Hanny- smith formula
% of I.C = 16 (XB – XA) + 3.5 (XB – XA)
XB & XA are the elelctronegativity of atoms of ‘B’ & ‘A’
1) Bond Length :
· The equillibrium distance between the nucleii of two bonded atoms is known as bond length.
· The bond length is expressed in angstrom units (or) nanometer.
1A0 = 10–8cm = 10–10m
1nm = 10–7cm = 10–9m = 10A0
• 1 pm = 10–10 cm = 10–12 m.
· Bond length is experimentally determined by spectroscopic, x-ray diffraction of e– diffraction methods.
· Covalent radius of AB is calculated by using pauling emperical equation dAB = rA+rB + C(XA–XB)
· For bond involving atoms of II period C = 0.08
· For Si, P, S bonded to more electronegative atoms not belongs to I period C = 0.06
· Examples for some bond lengths :
The bond length increases if the sizes of bonded atoms increases.

· The bond length decreases if the number of bonds between the atoms increases.
Example : i.e, Bond length

· The bond length increases if the P-character in a hybrid orbital increases.
· The C–H bond length in C2H2 < C2H4 < C2H6
· The magnitude of the bond length between the same two atoms in different substance is generally same.
· The O–H bond length in H2O, H2O2, C2H5O–H etc is same and it is equal to 96pm.
· The bond length in a hetero nuclear diatomic molecule is slightly less than the sum of the covalent radii of bonded atoms.
· A–B bond length is less than (Covalent radius of A+Covalent radius of B)
· The carbon, carbon bond length in
Benzene is 139pm
Graphite is 142pm
Diamond is 154pm
• For halogens :
Bond length size of atom F2 < Cl2 < Br2 < I2
2) Bond angles :
· The internal angle between the lines joining the centre of the nucleus of one atom to the centres of the nucleii of two other atoms combined to it is known as the bond angle.
· The bond angles are measured with methods like X-ray diffraction and spectroscopic methods etc.
· The bond angle is 1800 in BeF2, BeCl2, CO2, CS2, XeF2, ZnCl2, HgCl2, HCN, N2O, C2H2, I3– ion.
· The bond angle is 1200 in BF3, BCl3, SO3, CH2O, C2H4
· The bond angle is 109028| in CH4, CCl4, SiH4, SiCl4, NH4+
· The bond angle is 900 XeF4, BrF5, SF6
· Bond angles are 900 and 1800 in ClF3, BrF3, ICl3
· Bond angles are 900 and 1200 in PCl5, AsCl5, SbCl5
· The bond angle is 1070 in NH3, H3O+
· The bond angle in
H2O – 10405|
NF3 – 102030|
OF2 – 1020
Cl2O – 1110
3) Bond enthalpy:
· The energy required to convert one mole of diatomic molecules in the gaseous state into their constituent atoms in known as the bond dissociation energy.
· The energy released when one mole of diatomic molecules are formed in gaseous state by the combination of required number of gaseous atoms is known as the bond enrgy.
· The bond energy is numerically equal to the bond dissociation energy for a substance containing diatomic molecules.
· The bond energies of chlorine, nitrogen and oxygen are 242.4KJ/mole, 946 KJ/mole and 498 KJ/mole respectively.
· The bond energy depends on the mode of cleavage of the bond. The bond energy is less for homolytic cleavage and high for heterolytic cleavage.
· The bond energy for the cleavage of C2H5Br into and radicals is 280.9KJ/mole.
· The bond energy for the cleavage of C2H5Br into C2+H5 and Br– ions is 764.9KJ/mole.
· The bond energy is very high for least reactive substances and most stable substances.
· The bond energy increases if the number of bonds between the atoms increases.
Ex :
bond energy is 341.1KJ/mole.
bond energy is 610.7kJ/mol
bond energy is 827.6kJ/mol.
· The bond energy increases if the P-character in the orbital involved in bond formation increases.
The increasing order of bond energy for the atomic orbitals is s < sp < sp2 < sp3
· The bond energy increases if the number of lone pairs of electrons on bonded aotm decreases.
Example :
The bond energy is 145.9kJ/mol
The bond energy is 163kJ/mol
The bond energy is 341.1 kJ/mol
· The bond energy increases if the size of the atom ‘X’ decreases in the molecule HX.
· The bond energies of H–I, H–Br, H–Cl and H–F are 298.3, 366.1, 431 and 568.2 kJ/mol respectively.
· The heat energy required to break all the bonds in a molecule is known as dissociation energy (or) heat of dissociation.
· The heat of dissociation of methane is 1663.6kJ/mol.
The heat of dissociation of ethane is 2836.6kJ/mol.
· The bond energy in a binary compound which contains poly atomic molecules having more than one bond of same kind is known as average bond energy.
Average bond energy =
· The heats of dissociation of methane and ethane are 1663.6 and 2836.6 kJ/mol.
• The average C–H bond enthalpy in methane = = 415.9 KJ/mol
• The C–C bond enthalpy in ethane
= 2836.6 – (6×415.9) = 341.2 kJ/mol
• Incase of H2O molecule, the enthalpy needed to break the two O–H bonds is not the same.

Average bond enthalpy of H2O =
= 464.5kJ/mol
4) Bond order :
· Bond order = [Nb–Na]
Nb = number of e– in bonding molecular orbitals
Na= number of e– in antibonding molecular orbitals.
• Bond order is useful for understanding the stabilities of molecules.
· Note:
If (i) Nb > Na bond order is +ve, molecule / ion is stable
(ii) Nb < Na bond order -ve, molecule/ ion is unstable
(iii) Nb=Na bond order is zero, molecule / ion is unstable (not exists)
• In the Lewis description of covalent bond, the Bond Order is given by the number of bonds between the two atoms in a molecule.
• Isoelectronic molecules and ions have idenical bond orders.
Ex: a) F2 and O2–2 have bond order =1.
b) N2, CO and NO+ have bond order = 3.
• With increase in bond order, bond enthalpy increases and bond length decreases.

· Bond orders of 1,2,3 correspond to single, double, triple bond respectively.
· Molecule/ ions with one (or) more unpaired e– are paramagnetic and with all paired e– are diamagnetic.
5) Resonance :
· The concept of resonance was introduced by “Linus Pauling”.
· A single Lewis structure cannot describe a molecule accurately.
· A no. of structures with similar energy, position of atomic nuclei, bonding, non – bonding pairs of electrons are taken as canonical structures of the real structure which describe the molecule accurately.
· Resonance is represented by a double headed arrow («)
· The more accurate structure of a molecule is called Resonance hybrid.
· Resonance can stabilise the molecule & average the bond parameters.
· Resonance stabilizes the molecule, because energy of resonance hybrid is less than the energy of any its one canonical structure.
· Resonance energy increases, stability of the molecule also increases.
· The canonical forms have no real existance.
· There is no equilibrium between canonical structure.
Ex : O3, C6H6, N2O5, PO4–3, SO4–2 ………… like molecules exhibit resonance structure.
Ex(1) : Resonance structure of ozone

· In Canonical structure O–O bond length is 148pm, O=O bond length is 121pm. But In resonance hybrid both O–O bond lengths are 128pm. due to Resonance structure the bond length is intermediate between single & double bond.
Ex(2) :

· In benzene all C–C bond lengths are equal (139pm) but not 154pm (C–C single bond length);
134 pm (C = C double bond length)

· This is applicable only for hetero nuclear molecules.
But not applicable for homonuclear molecules.
• Because it can be explain on the basis of electronegativity of bonded atoms.
· In a heteronuclear molecule, the shared electron pair between the two bonded atoms get displaced towards more E.N atom than other. The resultant covalent bond becomes polar covalent bond.
· Polarisation of the molecule possesses the dipolemoment.
Dipole moment (m) :
· In a polar covalent bond the more electronegative atom acquires negative charge and the less electro negative atom acquires positive charge. So a dipole is formed
Example : ________
· If two charges are separated by a distance, it is known as a dipole. The dipole is represented as + __
· In a polar covalent bond, the product of charge on the atom and the bond length is known as dipole moment of the bond (or) bond moment.
· It is represented by m(bond)
m(bond) = q x d
· The dipole moment is a vector quantity.
· In a poly atomic molecule containing polar covalent bonds, the vector sum of the dipole moments of all bonds present in it is known as the dipole moment of the molecule. It is represented by ‘m’
· The units for dipole moment are debye units (D)
1D = 10–18e.s.u–cm = 3.34×10–30 coulomb – meter
· The dipole moment can also be expressed in coulomb-metre.
· The dipole moments of water and ammonia are 1.85D and 1.47D.
· The dipole moment of the molecule depends mainly on the electro negativity difference between the bonded atoms rather than on the bond length.
Example : 1. Dipole moment of HF is more than that of HCl.
2. Dipole moment of H2O is more than that of H2S.
· The dipole moment is zero for symmetrical molecules.
· The dipole moment is not zero for unsymmetrical molecules (non- symmetrical molecules).
Symmetrical molecules (m = 0) :
• For symmetrical molecules m = 0 even though bond is polar.
• AB2 type Linear molecules.
Eg : CO2, CS2, BeCl2, ZnCl2, HgCl2, BeF2, XeF2
• AB3 type Planar triangle molecules.
Eg : BF3, BCl3, SO3 etc.
• AB4 type tetrahedron molecules.
Eg : CH4, CCl4, SiH4, SiCl4
• AB5 type trigonal bipyramid molecules.
Eg : PCl5, AsCl5, SbCl5
• AB6 type octahedron molecules.
Eg : SF6, SeF6
• P4, S8, C2H2, C2H4, C2H6, C6H6, XeF4, p- dichloro benzene, trans-1,2 dichloro ethene etc.
Non symmetric molecules : (m ¹ 0)
• AB type diatomic molecules.
Eg : CO, NO, HF, HCl, HBr, HI, ICl etc.
• ABC type linear molecules.
Eg : HCN
• A-A-B type linear molecules.
Eg :N=NO
• AB2 type angular molecules.
Eg : Cl2O, F2O, ClO2, H2O, H2S, SnCl2, SO2.
• AB2C type planar triangle molecules.
Eg : CH2O, COCl2
• AB3 type pyramidal molecules.
Eg : NH3, PH3, NF3, PF3, AsCl3
• AB3C (or) AB2C2 type tetrahedron molecules. (distorted tetrahedron)
Eg : CH3Cl, CHCl3, CH2Cl2
• O3, ICl3, ClF3, BrF5, IF5 , chloro benzene, m-dichloro benzene, o-dichloro benzene, cis-1,2-dichloro ethene etc.
· NCl3 is not a symmetrical molecule but its dipole moment is zero because (i) the electronegativity values of nitrogen and chlorine are equal (ii) N-Cl bond is not a polar bond.
· The dipole moment of ozone is not zero because ozone molecule contains a polar bond.
in its structure.
· The dipole moment of carbon monoxide molecule is very low (0.12D) than the expected value (2.7D) because of the presence of dative bond in its structure.
· In an angular molecule
m = 2 x bond moment x cos
where q is the bond angle in the molecule.
· Bond angle inversely related to dipole moment.
· The percentage of ionic nature in a bond
· The dipole moment is very high for polar solvents.
· Water, liquor ammonia, liquid sulphur dioxide, hydro fluoric acid etc are polar solvents.
· The dipole moment is zero (or) very low for non polar solvents.
· Benzene, carbon disulphide, carbon tetra chloride etc are non polar solvents.
· ‘m’ of NH3(1.47) > NF3(0.23)
· ‘m’ of CH4(0) = CCl4(0) < CHCl3(1.04)
Order of dipole moments of HX :
HF > HCl > HBr > HI
(1.78) (1.07) (0.79) (0.38)
· ‘m’ of H2O(1.85) > H2S(0.95)
• Based on the phenomenon of Polarisation.
Fajan proposed following rules. They are used to know the relative natures.
• Fajan’s rules are used to predict whether a chemical bond will be covalent (or) ionic. It depends on the charge of the cation and anion, and the relative size of cation and anion
a) Ionic nature a size of cation
a1/ size of anion
b) Ionic nature a1/charge on cation
a1/charge on anion
· A compound is more ionic if it contains large cation and small anion (Cs+>Rb+>K+>Na+>Li+)
· A compound is more ionic if it carries less positive charge and anion carries less negative charge (Na+CI->Mg2+O2->Al3+N3-).
· A compound is more ionic if cation carries inert gas configuration.
· A compound is less ionic or more covalent, if it contains small cation and large anion.
· A compound is less ionic or more covalent, if cation carries more charge, anion carries high negative charge.
· A compound is more covalent, if cation is having
pseudo inert gas configuration.
· Ionic bond is formed easily
a) Between Alkali metals and Halogen atoms.
b) Between Alkaline earth metals and Halogen atoms.
c) Between Alkali metals and any chalcogen atoms.
d) Between Alkaline earth metals and O,S atoms.
e) Between Aluminium metal and O,F atoms.
f) Between Alkali metals and Hydrogen atom.
g) Between Alkaline earth metals and Hydrogen atom.
· The absolute value of the difference in electro negativity of two bonded atoms, provides a rough measure of the bond type.
· The condition for the formation of ionic bond is, Electro negativity difference between the bonded elements is more than 1.7
· Bond formed between very low electro negative element, very high electro negative elements is ionic.(Ex. Bond between Alkali metals and Halogens).
· Bond formed between atoms of two elements of very high electro negativity is a covalent bond (F-F).
· Bond formed between the elements with very low electro negativity is a metalic bond (Bond between metal atoms).
· If electro negativity difference between bonded atoms is zero, it is 100% pure covalent (F-F,H-H,CI-CI).
· If electro negativity difference is greater than zero, it is polar covalent bond ( < 1.7 ).
· The higher the oxidation state, the lesser is the formation of ionic bond.
Ex. SnCI2 (Sn =+2) is more ionic than SnCI4 (Sn=+4 )
MnO (Mn=+2) is more ionic than Mn2O7(Mn=+ 7)
· Most predominent ionic bond is formed by Cs and F in CsF
· Among the Alkali metals, least tendency to form ionic bond is shown by Li.
· There is no 100% ionic compound
·. This Theory was proposed by Gillespie and Nyholm.
· This theory explains about shapes, without taking the hybridisation concept into consideration.
· The shape of a molecule depends on the number of electron pairs in the valence shell (bonded or nonbonded) around the central atom.
· Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.
· The valence shell is taken as a sphere with the electron pairs localising on the spherical surface at maximum distance from one another.
· The electron cloud of a bond pair of electrons is localised in space between the bonded atoms,
· The bond is referred to as”localised electron pair bond”.
· The orbitals involved in bond are localised orbitals.
· A lone pair of electrons takes up more space round the central atom than a bond pair.
· The orbitals (bonded, Non bonded)changes their orientation, to minimise the repulsions.
· Order of repulsions, between various kinds electron pair is
· In VSEPR Theory, single bond, double bond, Triple bond,dative bond- is counted as only one bonded electron pair, because all the electropairs in the same’ bond are oriented in the same direction.
· Deviation of normal Tetrahedral Bond angle in NH3 is due to repulsion between lone pair and bond pair.
· Deviation of normal Tetrahedral bond angle in water molecule is due to the repulsion between lone pair and lone pair.
According to VSEPR theory, molecules are classified in to two types.
(i) molecules in which the central atom has no lone pair
On Central Type of Shape Bond
Atom Molecule Angle
L.P. b.p.

0 2 AB2 Linear 180°
0 3 AB3 Trigonal 120°
0 4 AB4 Tetrahedral 109°28′
0 5 AB5 Trigonal
bipyramidal 90,120° 0 6 AB6 Octahedral 90°
0 7 AB7 Pentagonal 72°, 90°

(ii) molecules in which the central atom has one (or) more lone pairs

On Central Type of Shape Example
Atom Molecule
L.P. b.p.

1 2 AB2E V-Shape SnCl2, PbCl2
1 3 AB3E Pyramid NH3, PH3
2 2 AB2E2 Angular H2O,SO2,H2S
1 4 AB4E See Saw SF4
2 3 AB3E2 T -shape ClF3
3 2 AB2E3 Linear XeF2
1 5 AB5E Square BrF5
2 4 AB4E2 Square planar XeF4

• L.P.– No. of lone pairs & B.P.– No. of bond pairs
S. Shape of Examples
No. Molecule/Ion
1. Linear BeCI2, HgCI2, CO2,CS2, XeF2,
C2H2, HCN, N2O.
2. Angular SO2, H2O, H2S, SnCI2, O3,
Cl2O, CIO2,OF2, NO2.
3. Trigonal planar BBr3, BCI3,SO3,CH2O,C2H4. NO31– ,CO32–.
4. Tetrahedron CH4, CCI4,SiH4.SiCI4,NH4+, BF4 –,ClO4 –
5. Distorted CHCI3, CH3- CH2-Cl (or) tetrahedron C2H5 – Cl , CH3Cl, CH2Cl2
6. Trigonal PCl5, AsCls, SbCl5 etc…. bipyramid
7. Octahedron SF6, SeF6, TeF6,
8. Pyramid NH3,PH3,PCI3,NCI3,H3O+,CIO3–
9. Square Pyramid BrF5,IF5
10. T- Shape BrF3, CIF3, ICI3
11. Square planar XeF4, [Cu(NH3)4]2+
12 pentagonal IF7
13. Seasaw SF4
· VBT was proposed to explain the directional nature of Covalent Bond.
· This explain the formation of covalent bond, based on”Orbital Overlapping” concept.
· VBT was proposed by “Heitler and London”, and further extended by “Pauling and Slater”, based on wave mechanics.
· A Covalent Bond is formed by the overlapping of Orbitals with unpaired electrons with opposite spin.
· Greater the extent of overlapping, the stronger is the bond formed.
· Orientation of resultant covalent bond depends on that particular orbital which has more directional character.
· The extent of overlapping of orbitals depends on their shape and sizes.
· Relative strength of overlapping of orbitals is in the order p-p>s-p>s-s
· Example for s-s overlapping is H2 molecule.
· Examples for p-p overlapping
F2,CI2,Br2 ,I2,O2,N2… etc.
· Examples for s-p overlapping
· An imaginary line, which connect the nuclie of two bonded atoms is called as “Inter neculear axis”.
• Two types of bonds are formed on acount of
A) Sigma (s) bond B) Pi (p) bond
• Sigma Bond(): A covalent bond formed by the overlapping of orbitals lying along the internuclear axis, in axial fashion or linear manner or end to end or head to head is called sigma bond.
• Pi Bond(): A Covalent bond formed by the incomplete sidewise or lateral overlapping of atomic orbitals lying perpendicular to internuclear axis is called as Pi Bond.(p and d atomic orbitals are commonly participate in Pi bond formation)
• In Case of Sigma bond, electron cloud is Cylindrical and Symmetrical along the internuclear axis.
• In case of Pi-Bond electron cloud is above and below the internuclear axis.
• A Pi bond is formed only after the formation of Sigma Bond.
• Every single bond is the sigma bond
• A double bond contains one sigma and one pi
• A Triple bond contains One sigma and two pi
• Strength of the Sigma bonds follows the order
p-p > s-p > s-s.
• Strength of the bonds follow the order
Triple Bond > Double bond > Single Bond.
Strength of s,p bonds:
· The strength of a bond depends up on the extent of overlapping.
s bond is stronger than p–bond
· s bond formed by large extent of overlapping of bonding orbitals along the internuclear. Ex is [head on head]
· p bond formed by small extent of overlapping of bonding orbitals.
· In the formation of multiple bonds between two atoms of a molecule, pi bonds(s) is formed after the formation of s bond.
1) High ionisation enthalpy :
Atoms having high ionisation enthalpy do not form cations easily. These elements prefer to form covalent bonds Ex:- Cl–Cl
2) Comparable Electronegativity :
Atoms with equal (or) nearly equal electronegativities tend to share equally a pair of electrons with opposite spins form a covalent bond [by better over lap of atomic orbitals]
· Covalent substances exist as solids, liquids and gases.
· Covalent substances have low melting and boiling points.
· Covalent substances dissolve in non – polar solvents.
· Covalent substances are bad conductors both in molten and aqueous state.
· Covalent reactions are very slow.
· Covalent bond is directional, hence they exhibt space isomerism.
· Covalent compounds contained discrete molecules, hence they possess molecular formula (also empirical formula).
· Weak Vanderwaal forces of attraction are existing between covalent molecules, they require 1 K Cal of energy to break these forces.
· Covalent attractions are minimum in gases maximum in solids.
· As molecular weight increases, Vanderwaal forces increases, melting and boiling points increases.
· Diamond, Graphite, Silicon, Silica,Boron Nitride, Boron Carbide, Silicon Carbide are examples for covalent polymers.
· Covalent substances do not contain free electrons and free ions.
· Graphite contains free electrons so it is good conductor of electricity.
· Due to the presence of free ions in the covalent compounds like HCI, HBr, HI, HNO3 aqueous medium they are good conductors of electricity.
· CCI4,CHCI3, CH2CI2, etc. can not give precipitate with silver nitrate due to the absence of Chloride ions.
· Non polar solvents like CCI4, C6H6,CS2, etc. have lower dielectric constant values. Hence covalent substances are readily soluble in these solvents.
(HCI,HBr,etc are soluble in water like solvent which have high dielectric constant values, though they are covalent.)
· Shape and Stability of molecules can be explained by the concept of hybridisation.
· The concept was proposed by Pauling.
· Hypothetical phenomenon of intermixing of atomic orbitals of almost same ene.rgy and their redistribution into equal number of identical orbitals is known as Hybridisation of orbitals.
· It is a concept applicable to central atom of a molecule, but do not corresponds entire molecule.
Salient features of hybridisations :
The main features of hybridisation are as under :
· In hybridisation, we mix a certain number orbitals, not a number of electrons.
· The number of hybrid orbitals forms equals the number of atomic orbitals that get hybridised.
· Only orbitals of similarly energy can be mixed to from hybrid orbitals.
· The hybrid orbitals formed are equivalent in energy and shape.
· The hybrid orbitals are more effective in forming stronger s bonds that lead to the formation of more stable molecules.
· The hybrid orbitals are directed in space so as to have minimum repulsions between electron pairs.
· Those atomic orbitals of central atom, which are not involved in hybridisation are called as unhybrid orbitals.
· Hybrid Orbitals are utilised in sigma bond formation.
· If necessary according to requirement, un hybrid orbitals are utilised in pi bond formation.
Important conditions for hybridisation :
· The orbitals present in the valence shell of the atom are hybridised.
· The orbitals undergoing hybridisation should have comparable energy.
· Promotion of electron is not an essential condition prior to hybridisation.
· It is not necessary that only half filled orbitals participate in hybridisation. In some cases, even filled orbitals of valence shell take part in hybridisation.
Hybridi No.of %s %p %d Geometry Bond sation H.O’s Angle sp 2 50 50 – Linear 180°
sp2 3 33.3 66.7 – Trigonal Planar 120°
(1/3rd) (2/3rd) –
sp3 4 25 75 – Tetra hedral 109°.28’
(1/4th) (3/4th) –
sp3d 5 20 60 20 Trigonal 120°;90°
(1/5th) (3/5th) (1/5th) bipyramidal
sp3d2 6 16.6 50 33.3 octahedral 90°
(1/6th) (3/6th) (2/6th)
sp3d3 7 14.3 42.9 42.9 Pentagonal 72°;90°
(1/7th) (3/7th) (3/7th) bipyramidal

(H.O.’S Means Hybrid Orbitals)
· sp hybridisation is present in
BeCl2,BeF2,CO2,CS2,HgCl2, C2H2,HCN,
triple bonded cabon atoms in Alkynes
· sp2 hybridisation is present in
BF3, BCl3,SO3, SO2 HCHO, O3,C2H4,
C2H6,NO3–, CO2–3
· sp3 hybridisation is present in
CH4,C2H6, CCl4, SiCl4, CH3 Cl, CH2Cl2, SiH4,BF4– NH3,NCI3,NH4+,PCI3,PCI4+,H2O,H3O+,CIO-, ClO2–, CIO3,CIO–4,CI2O,Diamond,SiO2,SiC,all the carbon atoms in Alkanes.
· sp3d hybridisation is present in
PCl5, AsCl5,SbCl5,CIF3,ICI3,BrF3, XeF2, I3–
· sp3d2 hybridisation is present in
SF6,SeF6,TeF6,BrF5,IF5,XeOF4,XeF4 … etc
· sp3d3 hybridisation is present in
IF7, XeF6 ….etc

H = No.of Hybrid Orbitals
V = valency of central atom (equal to its group number)
M = No. of Mono valent species attached to central atom[ like F, CI , Br , I ,H ..]
C = Charge on cation
A = Charge of Anion
• Based on Number of Hybrid Orbitals one can identify the type of Hybridisaiton.
• If No. of Hybrid Orbitals are 2 ® sp Hybridisation.
• If No. of Hybrid Orbitals are 3 ® sp2 Hybridisation.
• If No. of Hybrid Orbitals are 4 ® sp3 Hybridisation
• If No.of Hybrid Orbitals are 5 ®sp3d Hybridisation
• If No. of Hybrid Orbitals are 6 ® sp3d2 Hybridisation
• If No of Hybrid Orbitals are 7 ® sp3d3 Hybridisation
· Execited carbon atom configuration is 2S12P1x 2p1y 2p1z. Making use of these four unpaired electrons for bonding Carbon exhibits tetracovalency.(forms four covalent bonds)
· A carbon atom is said to be involved in Sp3. when it is bonded with four single bonds (which are sigma)
like ——
· A carbon atom is said to be involved in sp2 when it is bonded with two single bonds and a double bond
like =
· A carbon atom is said to be involved in sp when it is bonded with one single bond and a triple bond or if carbon is simultaneously bonded with two double bonds.
like – c c – (or) = c =
· In Alkanes, each and every carbon is involved in Sp3 hybridisation.
· In alkenes, the two double bonded carbon atoms are invovled in sp2 hybridisation.
· In alkynes, the two triple bonded carbons are invovled in sp hybridisation.
· Number of Hybrid orbitals in Alkanes = 4n
Number of Hybrid orbitals in Alkenes = 4n-2 (for alkenes with one double bond)
· Number of Hybrid orbitals in Alkynes = 4n-4 (for alkynes with one triple bond)
· An Alkanes with n-number of Atoms contains (n-1) number of Sigma bonds.
· It was proposed by sidg wick.N.V.
· A bond formed by the sharing of a pair of electrons between two atoms contributed by one atoms alone is called as Co-Ordinate Covalent Bond or Dative Bond.
· The chemical species which contributes electron pair is a donor and which receives is an acceptor. Hence the bond also called donor-acceptor bond.
· Donor must possess a non bonded electron pair, and acceptor, must possess atleast one vacant orbi;tal of suitable energy.
· Dative bond is formed, only after the formation of covalent bonds.
· Dative bond is present in following molecules ® N2O,O3, BF3,N2O4,N2O5.CO,B3N3H6,Al2Cl6, CuSO4, K4[Fe(CN)6], [Cu(NH3)4] SO4… etc
Ions – NH H3O+ NO BF PCI N2H
SiF, AICl.
· Note: In hydrated cation the bond between water molecule and cation is dative bond
® [AI(H2O)6]3+ ® has 6 – dative bonds
[Be(H2O)4]2+ ® has 4 – dative bonds
· When H2O is converted to H3O+, No change in hybridisation (here the shape changes from angular to pyramidal).
· When NH3 is converted to NH, No change in hybridisation (here the shape changes from pyramidal to tetrahedral).
· When BF3 is converted to BF, Hybridisation changes from Sp2 to Sp3.(here the shape changes from planar triangle to tetrahedral).
· When PCI5 is converted to PCl, hybridisation changes from Sp3d to Sp3d2.(here the shape changes from trigonal bipyramidal to octahedral).
Molecular orbital theory:
· Molecular orbital theory or Hund-Mulliken theory-Accoriding to this theory the atomic orbitals combine to form the molecular orbitals. The number of molecular orbitals formed is equal to the number of atomic orbitals involved and belong to the molecule.
• The molecular orbitals are formed by LCAO method (linear comebination of atomic orbitals) i.e. by addition or subraction of wave functions of individual atoms thus
yMO = yA + yB
yb = yA + yB
ya = yA – yB
yb = bonding molecular orbital
ya = Anti bonding molecular orbital
· Molecular orbital of lower energy is known as bonding molecular orbital and of higher energy is known as antibonding molecular orbitals.
· Molecular orbitals are characterised by a set of quantum numbers.
· Aufbau rule, Pauli’s exclusion principle and Hund’s rule are applicable to molecular orbitals, during the filling electrons.
· Their shape is governed by the shape of atomic orbitals.

• The increasing order of relative energies of M.O having upto 14 electrons.

• For more than 14 electrons

Electronic configuration and molecular behaviour :
• The distribution of electron among various molecular orbitals is called the electronic configuration of the molecular.
• The filling of electrons in the molecular orbitals by applying Aufbau principle and Hund’s rule.
• These is useful to get information about the molecule as follows.
Stability of molecules :
i) Nb > Na, the molecule is stable
Reason : More bonding orbitals are occupied & the bonding influence is stronger that results stable molecule.
ii) Nb < Na, the molecule is unstable
Reason : If the antibonding influence is stronger therefore the molecule is unstable.
Bond order:
• The relative stability of a molecule can be determined on the basis of bond order. It is defined as the number of covalent bonds in a molecule. It is equal to one half of the difference between the number of electorns in the bonding and antibonding molecular orbitals.

Nb – Number of bonding electrons
Na – Number of antibonding electrons
· Note:
If (i) Nb > Na bond order is +ve, molecule / ion is stable
(ii) Nb < Na bond order -ve, molecule/ ion is unstable
(iii) Nb=Na bond order is zero, molecule / ion is unstable (not exists)
Nature of bond :
• The bond orders of 1,2 or 3 correspond to single, double or triple bond. But bond order may be fractional in some cases.
Magnetic nature :
· If molecular orbitals (bonding or antibonding) occupies paired electrons are diamagnetic and unpaired are paramagnetic.
Bond length :
· The bond order between two atoms in a molecules may be taken as an approximate measure of the bond length.
· Bond length decreases with increase of bond order.
AsH3 SeH2 HBr
SbH3 TeH2 Hi
B.P: HF > HI > HBr > HCl
H2O > TeH2 > SeH2 > H2S
Bonding in some diatomic molecules and ions
a) Hydrogen molecule:
Total number of electrons = 2, filling in molecular orbitals we have (s1s)2
Bond order =
• Hence there is a single bond between two hydrogen atoms and due to absence of unpaired electorns it is diamagnetic
b) Helium molecule (He2):
• The total number of electrons = 4 and filling in molecular orbitals we have (s1s)2(s*1s)2 Bond order =
• Hence He2 molecule can not exist
c) Boron molecule (B2)
• The total number of electrons = 10 and filling in molecular orbitals we have
(s1s)2(s*1s)2(s2s)2(s*2s)2(p2px1= p2py1)
• Bond order =
• It is a paramagnetic.
d) Carbon molecule (C2)
• The total number of electrons = 12 and filling in molecular orbitals we have
(s1s)2(s*1s)2(s2s)2(s*2s)2(p2px2= p2py2)
• Bond order =
• It is a diamagnetic.
e) Nitrogen molecule (N2)
• The total number of electrons = 14 and filling in molecular orbitals we have
• Bond order =

• It is a diamagentic
f) Oxygen molecule (O2)
• Total number of electrons = 16 and electronic configuration is

• Bond order (O = O)
• As shown by electronic configuration the O2 molecule contains two unpaired electrons, hence it is paramagnetic in nature
g) ion:
• Total number of electrons (16-1) = 15, Electronic configuration

• Bond order = .
• It is paramagnetic
h) (Super oxide ion)
• Total number of electrons (16+1) = 17. Electronic fonguration

• Bond order
• It is paramagnetic
i) Peroxide ion ()
Total number of electrons (16+2) =18. The electronic configuration is

• Bond order = .
• It is diamagnetic
Atomic and Molecular orbitals – Main differences:
Atomic orbitals Molecular orbitals
1) They belong to one 1) They belong to all the
specific atom only atoms in a molecule
2) They are the internal 2) They result when
characteristic of an atomic orbital of
atom energies combine similar
3) They have simple 3) They have complex
shapes of geometries shapes
4) The atomic orbitals are 4) The molecule orbitals
named as s,p,d,f..etc. are named as s,p.
5) The stabilities of these 5) The stabilities of these
orbitals are less than orbitals are either more
bonding and more than or less than the atomic
the antibonding orbitals orbitals
Difference between s and p MO’s
s – molecular orbitals p – molecular orbitals
1) Formed by the end on 1) Formed by the sidewise
overlap along the overlap perpendicular
internuclear axis to inter nuclear axis
2) Overlapped region is 2) Over lapped region is
very large small
3) Rotation about the 3) Rotation about the inter
internuclear axis is nuclear axis is un-
symmetrical symmetrical.
4) Strong bonds are 4) Weak bonds are
favoured favoured

· Hydrogen bond in defined as the electrostatic attraction between the positively charged hydrogen atom in a molecule and negatively charged electronegative atom of the same molecule or another molecule.
· Hydrogen bond was presented by Latimer and Rodebush.
· Hydrogen bond is indicated by a dotted line( ——- ).
· The length of the hydrogen bond depends on the substance under investigation. It varies from 176 pm to 275 pm.
· The energy of hydrogen bond varies from 2 to 10 K.cals/mole.(80 kj/mole,far less than energy of covalent bond, it is 418 kj/mole.
· Hydrogen bond is weaker than covalent bond and stronger than Vanderwaal forces of attraction.
· Most electronegative atoms like Fluorine, Oxygen, Nitrogen only can involve in hydrogen bond. Chlorine atom very rarely involves in hydrogen bond.
Intra molecular hydrogen bond
· Hydrogen bond present in the same molecule is known as intra molecular hydrogen bond.
· The intra molecular hydrogen bond is present in substances like,
o-Chlorophenol, o-Nitrophenol, o-Nitroaniline
o-Hydroxybenzoic acid (Salicyclic acid)
· The intra molecular hydrogen bond is also found in o-Chlorophenol.
· Substances having intramolecular hydrogen bonds.
1. are less water soluble.
2. are steam voltile.
3. have low boiling points.
inter molecular hydrogen bond
· Hydrogen bond present between different molecules is known as inter molecular hydrogen bond.
· The inter molecular hydrogen bonds are present in substances like,
Water, Ammonia, Hydrofluoric acid, ortho Phosphoric acid, ortho Boric acid, p-Nitrophenol, p-Chlorophenol, p-Hydroxybenzaldehyde, p-Hydroxybenzoic acid, primary alcohols (CH3OH, C2H5OH), Fatty acids (HCOOH, CH3COOH) and Primary amines.
· Substances having inter molecular hydrogen bonds exist as associated molecules.
Ex :
Consequences of hydrogen bond :
· Liquids having hydrogen bonds between their molecules are called associated liquids
· Water, ammonia,hydrofluoric acid methyl alcohol etc., are associated liquids
· Liquids which do not contain hydrogen bonds between their molecules are called normal liquids
· Benzene, Carbon disulphide, Carbontetrachloride Acetone, Ether, Bromine, Nitrobenzene etc., are normal liquids
· Associated liquids have higher boiling points than normal liquids
· Liquids having very low boiling points are called volatile liquids
· The boiling point of an associated liquid depends on
1. strength of hydrogen bond present in it
2. number of hydrogen bonds present in one mole of that compound.
· The order of boiling points
1. PH3<AsH3<NH3<SbH3<BiH3
2. H2O>H2S<H2Se<H2Te
4. CH4<SiH4<GeH4<SnH4
· The boiling points of NH3, H2O, HF are more than those of PH3, H2S and HCl respectively because inter molecular hydrogen bonds are present in NH3& H2O and HF Hydrogen bonds are not present in PH3,H2S, HCl.
· The boiling point gradually increases from CH4 to SnH4 because there are no hydrogen bonds in CH4,SiH4,GeH4,SnH4.
· The —- hydrogen bond is stronger than — hydrogen bond. But then also boiling point of water is more than B.P of HF ( 19.4oc) This is due to
a) The no. of hydrogen bonds in water molecule per mole is more than the no.of hydrogen bonds in HF molecule.
b) In HF molecule hydrogen bond is present in both liquid state & vapour state. Therefore its heat of vapourisation is less.
· The boiling point of water is more than that of ammonia although NH3 contains more hydrogen bonds than one mole of water this is due to the presence of very strong hydrogen bonds in water than in ammonia
· The molecular weight of formic acid (or)acetic acid determined by using its solution in a non polar solvent like benzene is twice the expected value. this is due to the dimerisation of acid molecules in the solution the dimer formation takes place with the help of hydrogen bonds
· In ice every water molecule involves in four(4) hydrogen bonds.
· The ice is a tetrahedral three dimensional polymer.
· The two helical strands in the DNA molecule are joined by hydrogen bonds
· Covalent substances like Glucose, urea,sugar, ammonia, alcohol etc .dissolve freely in water because they from hydrogen bonds with water.
· Substance having inter moleculer hydrogen bonds are highly water soluble .they have high boiling points and they are not steam volatile.
· In a hydrated cation, the bond between water molecule and cation is dative bond.
· In a hydrated anion the bond between water molecule and anion is hydrogen bond.
· In a hydrated salt having even number of water molecule generally the water molecule are bonded only to the cation
· Hydrated salts having even number of water molecules.
BeSO4.4 H2O : AICI3. 6 H2O
MgCI2 . 6H2O : FeCI3 .6H2O
CaCI2 . 6H2O : Fe(NO3)3 .6H2O
· In a hydrated salt having odd number of water molecule one water molecule is bonded to anion and the remaing water molecule are bonded to the cation.
· Hydrated salts containing odd number of water molecules.
Example: CuSO4.5H2O: NiSO4.7H2O
MgSO4.7H2O: FeSO4.H2O
· In some hydrated salt (Eg.BaCI2, 2H2O; SrCI2.2H2O etc.) the water molecules are not
bonded either to the cation or to the anion.
· Most of the hydrogen bonds are asymmetric.
Ex :- X–H ……….. :Y
i.e., H–atom is not located exactly between X and Y atoms, but much closer to X than Y.
· H–bonds are linear (or) slightly bent, to maximising attraction between H, Y and minimising repulsion between X,Y.
· Valence angle (a) between H and Y–Z bond varies between 1000 to 1400

Here, Y is more electronegative than X, Z.

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