Chemical Bonding
Content
Hydrogen Bond
Hydrogen bond is formed when a slightly positive hydrogen atom, attached covalently to strong electronegative atom A (e.g., F, O
or N) is held by a non-bonded electron pair of another atom B. The coordination number of hydrogen becomes two and it acts as a
bridging atom between A and B.
Generally hydrogen bond is formed with only F, O and N atoms. Sometimes less electronegative atoms such as Cl, S etc.,
also take part in the formation of hydrogen bond. Hydrogen bond is denoted by dotted lines (...........). It can be defined as :
The attractive force that binds a hydrogen atom, which is already covalently attached with strongly electronegative atom of gain
element is known as hydrogen bond. The bond energy of hydrogen bond is 3--10 kcal/mole.
Types of Hydrogen Bonding
Three types of hydrogen bonding exist :
(i) Intermolecular hydrogen bonding (ii) Intramolecular hydrogen bonding
(iii) -Hydrogen bonding
(i) Intermolecular hydrogen bonding. Intermolecular hydrogen bonding exists between two or more molecules of the same or
different compounds.
(a) Homo-intermolecular hydrogen bonding. It is also termed as self-association. It refers to the association of two or more
identical molecules e.g., association in alcohol, association in water, association in NH3, association in HF etc.
(b) Hetero-intermolecular hydrogen bonding. It refers to the association of two different species. One which donates the lone
pair of electrons is called electron donor or hydrogen acceptor, while the other which donates proton is called proton donor.
Example of hetero-intermolecular hydrogen bonding are as follows :
Due to intermolecular hydrogen bonding the molecules are associated together to form a cluster, which results in the increase in
melting point and boiling point of the compound.
Intramolecular hydrogen bonding.
(ii) Intramolecular hydrogen bonding. When hydrogen bonding exists within the molecule it is called intramolecular hydrogen
bonding. In such type of hydrogen bonding two groups of the same molecule link through hydrogen bond, forming a stable five or
six membered ring structure e.g., salicylaldehyde, o-chlorophenol, acetylacetone, ethylacetoacetate etc.
This intramolecular hydrogen bonding was first called chelation (after the Greek word "Chela" meaning, claw) because in the same
molecule the formation of a ring hydrogen bonding is a pincer like action resembling the closing of a Crab's claw. Some more
examples of intramolecular hydrogen bonding are :
(iii) -Hydrogen bonding. Sometimes -electrons of an olefinic or aromatic system act as proton acceptor in hydrogen bonding.
Such type of hydrogen bonding is called -Hydrogen bonding. These are again of two types :
(a) Intermolecular -hydrogen bonding. It occurs between two or more molecules of the same or different compounds. For
example,
(b) Intramolecular -hydrogen bonding. This occurs between the groups of the same molecule. Hydrogen of a group (e.g., --
OH) forms hydrogen bond with -electrons of an aromatic ring or an olefinic bond. For example,
Properties of Hydrogen Bond
(i) Hydrogen bond is a bond of hydrogen between two electronegative atoms only. It never involves more than two atoms.
(ii) Bond energy of hydrogen bond is in the range of 3-10 kcal/mole. Thus hydrogen bond is a weaker bond than a covalent bond
(bond energy of a covalent bond is 50-100 kcal/mole). But it is stronger than Van der Waal's forces (1 kcal/mol).
(iii) In the formation of hydrogen bond electron pair is not shared. In this respect it is different from the covalent bond.
(iv) The strength of hydrogen bond depends upon the electronegativity of the atom A to which hydrogen atom is attached with a
covalent bond. As the electronegativity of A increases, strength of hydrogen bond increases. Thus HF will form most strong
hydrogen bond as fluorine is the most electronegative atom.
(v) Typical hydrogen bond is linear. Angular hydrogen bonds exist in solids or in intramolecular hydrogen bonding.
(vi) The bond length A -- H and B -- H are generally different except HF2
-
(F -- H -- F) which is a symmetrical ion.
Effect of Hydrogen Bonding on Physical Properties of the Molecules
Unusual physical properties of H2O, HF, NH3, and alcohols can easily be explained on the basis of hydrogen bonding.
(1) Physical States
i) H2O and H2S. As we have already stated that the ease of formation of a hydrogen bond decreases as the electronegativity of the
atom attached to the hydrogen decreases. Oxygen is more electronegative than sulphur. There is a considerable hydrogen bonding
in H2O while in H2S the same is absent. H2O molecules are associated together in which hydrogen atom acts as bridge between two
oxygen atoms and the intermolecualr distance decreases, therefore, H2O exists as liquid.
(ii) HF and HCl. In HF, molecules are associated through hydrogen bonding and it is a liquid at ordinary temperature.
In HCl, due to less electronegativity of chlorine atom and its large size, hydrogen bond does not exist hence molecules of HCl are
not associated as in HF. Therefore, HCl is a gas at ordinary temperature.
(2) Melting Points and Boiling Points
(i) M.P. and B.P. of Hydrides of Oxygen, Fluorine and Nitrogen
The melting points and boiling points of compounds in a group of the periodic table increase with the increasing molecular
weights. This is evident from the melting point and boiling point curves of IV A group hydrides. [(a) and (b)].
The melting point and boiling point of CH4, SiH4, GeH4, and SnH4 decrease with decreasing molecular weights. But in case of VA,
VIA and VIIA groups the melting point and boiling point of H2O, NH3 and HF are exceptionally high than the hydrides of other
members of their groups.
The melting point and boiling point of the hydrides of the elements of IVA, VA, VIA and VIIA groups can be represented as
below :
IVA CH4 < SiH4 < GeH4 < SnH4
VA NH3 > PH3 < AsH3 < SbH3
VIA H2O > H2S < H2Se < H2Te
VIIA HF > HCl < HBr < HI
It is clear from these plots that there is a sudden increase in melting point and boiling point of HF, H2O and NH3. The existence of
hydrogen bond in these molecules exceptionally increases their melting point and boiling point. Boiling point of water is higher
than that of hydrogen fluoride because the extent of association through hydrogen bonding in water is more than hydrogen fluoride.
Since CH4 cannot form hydrogen bond, its melting point and boiling point are the lowest among the hydrides of carbon family.
(ii) Melting point of ortho- nitrophenol and para- nitrophenol
Intramolecular hydrogen bonds do not involve molecular association and in these, change in physical properties is quite different
than in case of intermolecular hydrogen bonds. For example the ortho isomers have lower melting point (Intramolecular hydrogen
bonding) than respective para isomers (Intermolecular hydrogen bonding) as shown above.
Melting points of substituted Nitro compounds
Solubility
(3) Solubility. Alcohols, glycol, glycerol and sugars are soluble in water due to the formation of hydrogen bond with water
molecules. Dimethylether, (CH3)2O, is miscible in water as it can form hydrogen bond with water molecule but dimethyl sulphide,
(CH3)2S, is immiscible as it cannot form hydrogen bond with water molecules since the electronegativity
of sulphur is less.
(4) Viscosity. Viscosity increases with the extent of hydrogen bonding in molecules. The viscosity of water is 10.05 millipoise,
methanol 6 millipoise and dimethyl ether 2.3 millipoise. Since both H2O and CH3OH are hydrogen bonded the viscosities are high,
but when there is substitution of second methyl group to produce the non-hydrogen-bonded dimethyl-ether, (CH3)2O, the viscosity
drops to a low value. Polyhydroxy alcohols such as ethylene glycol, CH
2
OH.CH
2
OH and glycerol, CH
2
OH.CHOH.CH
2
OH which
have extensive hydrogen bonding exhibit much higher viscosities.
(5) Molecular weights. The association of two or more molecules by intermolecular hydrogen bonding affect the apparent
molecular weight. In case of carboxylic acids (RCOOH) it is observed that the apparent molecular weights are
higher than the formula weights. The apparent molecular weight decreases with increasing temperature due to dissociation of dimer
into monomer.
A monomer-dimer hydrogen-bonded equilibrium is the simplest explanation of these results. Increase in temperature increases the
average kinetic energy of the molecules, breaking more hydrogen bonds and shifting the equilibrium to left.
(6) Dielectric constants and Dipole moments
The formation of hydrogen-bond, A -- H...B leads to an increased polarity of the bond A -- H, and hence, to a larger dielectric
constant and greater dipole moment.
(7) Low density of ice than water
In the crystal structure of ice the oxygen atom of water is surrounded by four hydrogen at oms, two attached with covalent bonds
and two with hydrogen bonds. Thus in ice every water molecule is associated with four other water molecules in tetrahedral
pattern. Ice has an open structure with large empty space due to existence of hydrogen-bonds. When ice melts a number of
hydrogen bonds are broken and the space between water molecules decreases and the density of water increases, therefore from 0°
to 4°C, it is maximum. Above 4°C the increase in kinetic energy of the molecules disperse them and the result is that the density
now decreases with increasing temperature.
(8) Stability of unusual structures
Generally, the organic compounds with two -- OH groups on the same carbon atom are unstable and soon liberate water molecule.
For example :
The stability of compounds like chloral hydrate, CCl3 -- CH(OH)2 can be explained on the basis of intramolecular hydrogen-
bonding.
(9) Chain, Sheet and Three dimensional structures
Hydrogen bonding leads to the formation of chains (HCN, HF, HCOOH), sheet (orthoboric acid, oxyde acid) and three
dimensional network (water, KH2PO4) structures.
(10) Dissociation constants of carboxylic acids
The dissociation constant of an acid depends on the stability of its ion. If the stability of anion of an acid is increased due to
intramolecular hydrogen-bonding, the acid strength is greatly enhanced i.e., pKa value decreases. The carboxylate ion of o-hydroxy
benzoic acid is stabilised by intermolecular hydrogen-bonding, thus o-salicylic acid is more stronger (pKa = 2.89) than benzoic
acid (pKa = 4.17).
It can be seen that two hydrogen-bonds would be expected to bring more stabilization than one hydrogen bond, and 2,6-dihydroxy
benzoic acid is much more stronger (pKa = 2.30) than o-salicylic acid. Similarly, o-salicylic acid (pKa = 2.98) is much stronger
due to intramolecular hydrogen bonding than its meta (pKa = 4.08) and para (pKa = 4.58) isomers.
Fajan's Rule
It is generally assumed that covalent and ionic bonds are entirely distinct but this is probably not a totally valid assumption. Bonds
intermediate between ionic and covalent do occur through a process of deformation or polarization.
Ionic polarization is favoured by a number of factors which are summarized in the four Fajan's rules.
(1) Small cation _ In a small cation there is greater concentration of positive charge over a small surface area so it will cause
greater deformation of an anion than would he caused by a large cation. Thus small cations have high polarizing power.
The effect of cationic size upon covalent character is shown in table below. It is clear from the table that with the increase in the
size of the cation the covalent character decreases.
This is justified by the low melting point of beryllium chloride which is more covalent than the chlorides of the alkaline earth
metals.
Effect of cationic size upon covalent character
(2) Large anion _ The large anion has polarizability. The outermost orbitals of the anions are shielded from the nucleus by a
number of completely occupied orbitals hence they are readily polarized by a small cation.
If the size of the cation and the charge on both the ions is kept constant and only the size of the anion is increased, more covalency
will be noticed. This is shown in table from the decrease in the melting points of calcium halides.
Effect of anion size upon covalent character
(3) Large charge on either of the ions_ It is understandable that the electrostatic forces which cause polarization will be
considerably increased if the ions are highly charged. The increased nuclear charge will attract the ions to a greater extent
causing greater deformation and hence covalent character. The decrease in melting points with the increase of charge of the
ions is shown in table to justify this generalization.
Effect of cationic charge upon covalent character
Cation with non-inert gas atom structure
(4) Cation with non-inert gas atom structure _ The cations with the inert gas electron configuration are most effective in
shielding the nuclear charge from its surface while the cations with non-inert gas atom structure have positive fields at their
surfaces and consequently will possess high polarizing powers. Thus the cation should possess an electronic configuration which is
not that of an inert gas.
Hence, if the charge and size are kept nearly constant, cations with 18-electron structure cause greater anion deformation
than those with 8-electron arrangements. It is shown in table by the comparison of the melting points of anhydrous chlorides of IA
and IB group of periodic table.
Effect of 8 and 18 electronic shell upon the covalent character
On the basis of these important general rules it becomes possible to predict the type of bond that a given element is likely to
prefer.
Rules(1), (3) and (4) indicate that the cations which are large and have small charge and possess inert gas electronic configuration
should possess least polarizing power e.g., the large alkali metal ions. Thus the large alkali metal ions will prefer to form ionic
bond.
Similarly, the small halide ions will favour an ionic bond because they will form ions having the least polarizability. Accor ding to
rules (2) and (3) the most stable anions are those which are small and have only a small charge.
The fourth Fajan's rule suggests that, in general, the non-transition elements are more ionic than the transition elements because
their cations have lower polarizing power and so the cations are more stable.
Applications of Fajan's Rule
(1) Melting point
(2) Diagonal relationship
We know that chemistry of lithium, berylium and boron resembles with that of magnesium, aluminium and silicon, respectively.
The diagonal relationship observed between the following pairs of elements can also be explained with the help of Fajan's rules.
On moving to the right across a period in the periodic table the charge of the cation increases and the size decreases. Consequently'
the polarizing power will also increase. In a vertical group, with the increase power of the cation will correspondingly decr ease. If
both moves are made simultaneously, as in a diagonal relationship then two elements of similar polarizing power may result, e.g.,
polarizing power of Be2+ and Al3+ is almost similar as their ionic potential ( ) are also almost similar [ of Be2+ = 6.48 and of
Al3+ = 6.0]. Such elements will form bonds of a similar type in the corresponding compounds. This explains almost identical
chemical and physical properties of the above mentioned pairs of elements.
(3) Non- polar character and colour
The increase in nonpolar character of inorganic salts is manifested in the appearance or enhancement of colour. Thus:
Colour deepening tendency polarization of anion size of anion
(i) The oxides of colourless cations are usually white but the corresponding sulphides are likely to be deeply coloured if the cation
is one which has a tendency to polarize anions. With a few exceptions, the white metal sulphides are only those of alkali and
alkaline earth metals.
(ii) In a series of halides of ions such as Ag
+
, the fluorides and chlorides are colourless ions is usually an indication of an
appreciable amount of nonpolar character or some other unusual structural feature. An appreciable amount of polarization leads to
intense absorption bands.
(4) Solubility
Solubility of salts in polar solvents like water is affected by polarization. The example of silver halides may be considered in which
there is polarizing cation and increasing polarizable anions.
Silver fluoride is quite ionic and soluble in water. Less ionic silver chloride is soluble only after complexation with ammonia.
silver bromide and silver iodide are insoluble even with the addition of ammonia. Increasing covalency from fluoride to iodide is
expected and decreases solubility in water is observed. However, many other factors are involved in solubility in addition to
covalency.
(5) Chemical reactions_ Stability of metal carbonates
Chemical reactions can often be correlated in terms of the polarizing power of a particular cation. For example, in alkaline earth
carbonates, there is a tendency towards decomposition with the evolution of carbon dioxide.
With the increase in ionic potential of metal ion [charge/radius], its tendency to accept oxide (O2_) to form metal oxide increases.
Hence, the stability of metal carbonates increases down the group in periodic table, radius of the metal ion increases, ionic
potential decreases and stability of the metal carbonate increases as shown in table.
Along the period (from left to right) charge on the metal ion increases, ionic potential increases and stability of the metal car bonate
decreases e.g., stability of K2CO3 > CaCO3 > CuCO3.
Further, the effect of d electrons is also evident. Cd2+ and Pb2+ are approximately of the same size as Ca2+ but both CdCO3 and
PbCO3 decosmpose at approximately 350°C.
(6) Acidic, Basic and Amphoteric character of oxides
If ï < 2.2 metal oxide is basic e.g., MnO, CrO, Na2O, MgO etc.
If ï = 2.2 to 3.2 metal oxide is amphoteric e.g., MnO2, CrO2 etc.
If ï > 3.2 metal oxide is acidic e.g., MnO3, CrO3, Mn2O7 etc.
With the increase in ionic potential of metal ion, polarizing power increases, covalent character of M — O bond increases, bond
does not dissociate on hydrolysis and the acidic character increases.
Attempt has also been made to correlate the enthalpy of carbonates, sulphates, nitrates and phosphates with increasing charge and
size function of the cation. It may therefore be stated in general that size and charge are important factors governing the polarizing
power of ions and consequently, many of their chemical properties.
Results of polarization
Polarization results in increasing covalent character in predominantly ionic bonds which may in turn affect the melting and boiling
points of ionic compounds, their decomposition temperature, solubility, colour, chemical reactions etc. Possible correlation of these
properties with polarization has been discussed earlier.
Solubility of compounds
Solubility of diffrent compounds in solvents depends upon many fectors. It is very complicated property because many factors
control this property simultaneously for example
(a) Nature of solute (b) Nature of solvent (c) Temprature of reaction
(d) Pressure (e) Lattice energy (f) Solvation energy etc.
When any solute dissolves in solvent to give the saturated solution, heat evolved or absorbed to the surroundings is known as heat
of solution.
Whether the heat of solution is positive or negative depends on the nature of solute and solvent. When any solid are dissolved in
water enthalpy of over all reaction is depends upon strength of two energies, that is
(i) Energy required to break down one mole ionic crystal lattice into their respective ions known as lattice energy.
(ii) Energy liberated when the ions are solvated or hydrated by solvent there is process of neutralisation, known as solvation energy
or hydration energy. This energy actually takes into account of both solvent. Solvent interaction (energy required to make a hole in
water) and solvent solute interaction. These two are combined together because experimentally they are hard to seprate. Energy
change involved in dissolution of a salt represented by born haber cycle.
In above process U is the lattice energy of the crystal, DH solvation is the energy liberated when positive and negative ions gets
solvated and DH solution is the observed heat of solution at in finite dilution. The total energy change from MX(s) to their
respective solvated ions is actually independent of the path. Above representation of ionic compound in water is born-Haber type
cycle. In case of KCl, overall process can be imagined to occur in two consecutive steps
First step involves vapourizing solid requires energy i.e., work must be done to seprate the ions while second step is exothermic
because process of solvation where ion-dipole attraction liberate energy. When solvation or hydration energy is greater than the
lattice energy the overall process is exothermic so liberation of energy takes place, normally solute is soluble in solvent while if
lattice energy is more than hydration energy, process is endothermic and its difficult for solubility of compound in solvent.
Factors affecting Hydration (solvation) energy
(a) Size of ions-
As the size of ions decreases, more in teraction between ions and H2O takes place so hydration energy increases. For example
hydration energy of:
In transition series (d-Block) size of ions are almost same due to balancing of shielding effect and nuclear attraction force,
therefore their hydration energy is almost constant. Transition metal ion can form strong bond with water due to presence of vacant
d-orbital so they contain high solvation energy.
(b) Charge on ions _ As the charge on ions increases, attraction between ions and dipole (H2O) increases so hydration energy
increases.
(c) Dielectric constant of solvent — Measurement of the tendency to attract solute particals by solvent is known as DEC. As the
DEC increases force of attraction between ions decreases so solvent with high DEC is responsible for solubility of compound.
Factors affecting Lattice energy
(a) Size of ions — As the size of ions decreases their is compact packing of atoms so system is stable with high lattice energy for
example LiF contains high lattice energy.
(b) Charge on ions — As the charge on ions increases, their is high attraction between ions so stability increases which increases
lattice energy for example Al2O3 > Mgcl2 > NaCl.
(c) Bond character _ Ionic compounds have high lattice energy than covalent compound due to more stability.
Factors affecting solubility of ionic compounds
Solubility of compounds mainly control by solvation energy and lattice energy with the support of some minor factor like DEC of
solvent, hydrogen bonding etc but it is very difficult to set a trend in periodic table for solubility of different compounds because
both solvation-energy and lattice energy decreases with increase in size, thing is which decreases more with respect to other,
actually this decides the solubility of compound.
Generally, ionic substances dissolve more readily in solvents composed of molecules containing electrostatic dipoles.
Solids are usually crystalline in nature. In the crystal lattice the ions are arranged in a symmetrical fashion and are held in thei r
relative positions by strong electrostatic forces resulting from the charge upon the ions. To breakdown the arrangement of ions in
the crystal, these forces must be overcome. Thus, for a substance to be readily soluble, more energy must be provided for the
separation of the ions from the crystal than was liberated in building up the ionic lattice. In other words, the energy of solvation
must be of greater magnitude than the lattice energy. (The interaction that takes place when a substance is introduced into a solvent
is called solvation and the energy change involved in this process is known as the solvation energy).
Thus, both the solvation energy and the lattice energy affect solubility of ionic compounds but in an opposite manner. The
important factors affecting solubility of ionic compounds are discussed below.
1. Nature of solvents. As the DEC of solvents increases solubility of ionic compounds increases because it weakens the force of
attraction between ions. Force of attraction between ions given by formula
DEC of water is very high (81), it is the best solvent for ionic compounds. Hydrogen peroxide has higher DEC (92) than water but
it not used as a solvent for ionic compounds because it undergo decomposition at room temprature so work as pwerful oxidising
agent.
H2O2 H2O + [O]
Nonpolar solvents like benzene, ether and CCl4 fail to solvate the ions so these are not use to dissolve ionic compounds.
2. Size of ions. Both lattice energy and solvation energy depends upon size of ions on following manner
where r
+
and r
-
are radii of cation and anion. It is clear from the above relations that decrease in the size of the ions affect the two
energies in a similar manner. However, the two energies are influenced to different extents and the predominating energy affects
the solubility to different extents.
In case of compounds containing large anions e.g., I
-
, SO4
2-
, CO3
2-
, PO4
3-
etc., the solubility will decrease with increase of
cationic size. For example, in case of sulphates of alkaline earth metals, the decrease in solvation energy is more rapid than the
decrease in lattice energy with the increase in the size of cation. The solubility of alkaline earth metal sulphates decreases in the
following order :
MgSO4 > CaSO4 > SrSO4 > BaSO4
Similarly, the solubility of alkali metal iodides decrease in the following order :
LiI > NaI > KI > RbI > CsI
In case of compounds containing small anions e.g., F—, the decrease in lattice energy is more rapid than the decrease in
solvation energy with increasing size of the cation. Thus the solubility of the fluorides of alkai metals increase as the size of cation
increases as given below :
CsF > RbF > KF > NaF> LiF
Solubility of II group carbonates in H2O
BeCO3 > MgCO3 > CaCO3 > SrCO3 > BaCO3 (Large size anion)
BeF2 > BaF2 > SrF2 > CaF2 > MgF2
(Small size anion but exceptionally high solubility of BeF2 due to high hydration energy of Be
+2
)
Ba(OH)2 > Ca(OH)2 > Mg(OH)2 > Be(OH)2 (Small size anion)
Solubility in water
LiI > LiBr > LiCl > LiF
CaI2 > CaBr2 > CaCl2 > CaF2
(Due to decrease in lattice energy by increase in size)
3. Ionic Charge. With the increasing ionic charge, the lattice energy increases much more rapidly than the solvation energy. Thus,
solubility of ionic compounds decreases very sharply as the ionic charge increases.
Effect of Ionic charge on Solubility
4. Polarization of Anions. Large anions are more polarizable than small anions (Fajan's Rule). Polarization of anion increases
covalent character in the molecule, hence decreases solublility in water. This explains the order of solubility of silver hal ides in
water as given below :
AgF > AgCl > AgBr > AgI
Cations with 18-electrons structure polarize anions more than those with 8-electron arrangements. Polarization of anion decreases
the solubility. This explains the following order of solubility
KCl > AgCl
NaCl > CuCl
Compound Cationic structure Cationic radius Å Solubility
Effect of Temperature
5. Effect of Temperature. Formation of a solution may be exothermic or endothermic process and may be represented as follows :
Exothermic : Solute + Solve Solution + Heat ....(1)
Endothermic : Solute + Solvent + Heat Solution ....(2)
The addition of heat (i.e., a rise in temperature) in equation (2) causes more of the solute to dissolve. This is the case with most of
the solid- liquid solutions, where the solubility increases with rise in temperature. In equation (1) the solublility decreases with the
rise in temperature. For example, the solubilities of KNO3, NaNO3, KCl, NH4Cl etc. increases with the rise in temperature whereas
that of Na2SO4 decreases with the rise in temperature.
6. Hydrogen bonding. It is very minor factor which control solubility of compounds. In the system where solute and solvent
particals shows association with H-Bonding are more soluble than other combination. For example NH3 is more soluble in H2O
than PH3, ROH are more soluble in H2O than RSH due to association. In case of ROH & ROR their solubility in water is almost
same because both show H-Bonding with H2O. Diols and triols are much more soluble than mono hydroxyderivatives because they
form more effective hydrogen bonding.
(1) Types of Covalent Bonds
It has already been discussed that the formation of a covalent bond involves the overlapping of half -filled atomic orbitals. The
covalent bonds can be classified into two different categories depending upon the type of overlapping. These are :
(a) Sigma covalent bond (b) Pi covalent bond.
(a) Sigma ( ) bond. This type of covalent bond is formed by the axial overlapping of half-filled atomic orbitals. The atomic
orbitals overlap along the inter-nuclear axis and involve end to end or head on overlap. The electron cloud formed as a result of
axial overlap is cylindrically symmetrical about inter-nuclear axis. The electrons constituting sigma bond are called sigma
electrons. There can be three types of axial overlap as discussed below :
(i) s-s overlap. It involves mutual overlap of half-filled s-orbitals of the atoms approaching to form a bond. The bond formed is
called s-s bond.
(ii) s-p overlap. It involves mutual overlap of half-filled s-orbital of the one atom with half-filled p-orbital of the other. The bond
so formed is called s-p bond.
(iii) p-p overlap. It involves mutual overlap of half-filled p-orbitals of the two atoms. The bond so formed is called p-p bond.
The s-s, s-p and p-p overlaps have been shown diagramatically in Figure below.
Strength of three types of sigma bonds. The strength of three types of sigma bonds varies as follows :
p-p > p-s > s-s
It is because of the fact, that p-orbitals allow overlap to a greater extent as compared to p-s which is larger as compared to s-s
overlap.
(b) Pi ( ) Bond. This type of covalent bond is formed by the lateral or sidewise overlap of the atomic orbitals. The orbitals
overlap takes place in such a way that their axes are parallel to each other but perpendicular to the internuclear axis. The pi bond
consists of two charge clouds above and below the plane of the atoms involved in the bond formation. The electrons involved i n
the p-bond formation are called -electrons.
It may be noted that :
(i) Sigma bond is stronger than pi bond. It is because of the fact that overlapping of atomic orbitals can take place to a greater
extent during the formation of sigma bond whereas overlapping of orbitals occurs to a smaller extent during the formation of pi
bond.
(ii) Pi bond between the two atoms is formed only in addition to a sigma bond. It is because of the fact that the atoms constituting
a single bond prefer to form a strong sigma bond rather than a weak pi bond. Thus, pi bond is always present in molecules having
multiple bonds, i.e., double or triple bond. In other words, a single bond cannot be a pi bond.
(iii) The shape of molecule is controlled by the sigma frame work (orientations of sigma bonds) around the central atom. Pi bonds
are superimposed on sigma bonds hence they simply modify the dimensions of the molecule.
Compraison between sigma and pi bonds. The various points of distinction between sigma and pi bonds are given in Table
below.
Table below. Comparison of Sigma and Pi Bonds
he formation of sigma and pi bonds, in oxygen (O2) molecule.
Oxygen atom (8O) has two half-filled p-orbitals in its valence shell as is evident from its electronic configuration (1s
2
, 2s
2
, 2px
2
,
2py
1
>, 2pz
2
1). One of the half-filled p-orbital overlaps axially with half-filled p-orbital of the other oxygen atom to form s bond.
The other half-filled p-orbitals of the two atoms overlap sidewise to form a bond which is denoted as p -p bond. The
formation of molecule is shown in Fig. below.
Thus, O = O bond consists of one s bond and one p bond.
Bonding Parameters
Covalent bonds are characterised by the following parameters, bond energy, bond length and bond angle.
(a) Bond Energy
It has already been pointed out that the formation of a bond occurs as a result of decrease of energy. Therefore, same amount of
energy is required to break the bond between the two atoms. For example, the energy released during the formation of bonds
between the gaseous hydrogen atoms to form one mole of hydrogen moleculs is 433 kJ mol
-
1. This energy involved in making or
breaking of bonds is referred to as bond energy. Thus, bond energy may be defined as the amount of energy required to break
one mole of bonds of same kind so as to separate the bonded atoms in the gaseous state.
The magnitude of bond energy reflects the strength of the bond. Its magnitude depends upon the following factors :
(i) Size of the participating atoms. Larger the size of the atoms involved in bond formation, lesser is the extent of overlapping and
consequently, smaller is the value of bond energy.
For example, bond energy of Cl—Cl bond is 237 kJ mol
-
1 whereas that of H—H bond is 433 kJ mol
-
1.
(ii) Multiplicity of bonds. The magnitude of bond energy increases with the multiplicity of bonds even though the atoms involved
in the bond formation are same. It is because of the fact that with the multiplicity of bonds the number of shared electrons between
the atoms increases. As a result, the attractive force between nuclei and electrons also increases and consequently, the magnitude of
bond energy increases. For example, bond energy of C — C bond is 348 kJ/mol
-
1 but that of C = C bond is 619 kJ mol
-
1. The
average bond energies of some bonds are given in Table below.
(iii) Number of lone pairs of electrons. Greater the number of lone pair of electrons present on the bonded atoms, greater is the
repulsive interactions between them and smaller is the bond energy. Let us compare the bond energies of some of the single bonds
Table below. Bond Energies of Some Common Bonds
Bond Length
(b) Bond Length
It has already been discussed that two bonded atoms in a molecule remain held up at a certain distance from each other. They
cannot approach too close because it leads to repulsive interactions and potential energy of system increases. This minimum
distance between the bonded atoms is referred to as bond length. Thus, bond length may be defined as the average distance
between the centres of nuclei of the two bonded atoms in a molecule. Bond length is usually expressed in Angstrom units (Å) or
picometers (pm) and it can be determined experimentally by X-ray diffraction and and other spectroscopic techniques.
Bond length depends upon the size of the atoms and nature of bonds.
(i) Bond length increases with the increase in the size of the atoms, e.g., bond length between hydrogen and chlorine atoms in HCl
molecule is 127 pm whereas bond length between carbon and chlorine atoms is C—Cl bond 177 pm.
(ii) Bond length decreases with the multiplicity of bonds. It is because of the fact that larger the number of electrons shared by the
two atoms greater will be attractive force between electrons and the nuclei and consequently, lesser is the bond length. For
example, bond length of C—C bond is 154 pm whereas that of C = C bond is 134 pm. Bond lengths of some common bonds are
given in Table below.
Table below. Bond Lengths of Some Common Bonds
(c) Bond Angle
We know that covalent bonds are formed by overlapping of atomic orbitals. Due to directional character of atomic orbitals , the
covalent bonds in a molecule are oriented in specified directions. The bond angle is defined as the average angle between the
lines representing the orbitals containing the bonding electrons.
Bond angle is expressed in degree/minute/seconds. For example, H—C—H bond angle in CH4 molecule is 109° 28¢. Similarly,
F—B—F bond angle in BF3 is 120° and H—N—H bond angle in NH3 molecule is 107°.
The bond angles in CH4, NH3, H2O and BF3 molecules are shown below in Figure below.
Concept of Hybridisation
It has already been pointed out that covalency of an element is equal to the number of half-filled orbitals present in the valence
shell of its atoms. On applying this concept to carbon, we find that the valency of carbon should be equal to 2 because it has only
two half-filled orbitals in the valence shell.
6C : 1s
2
, 2s
2
, 2px
1
, 2py
1
, 2pz
0
In the same way, the valency of beryllium should be zero and that of boron should be one as is evident from their ground stat e
configurations.
4Be : 1s
2
, 2s
2
, 2p0
5B : 1s
2
, 2s
2
, 2px
1
, 2py0, 2pz
0
Contrary to this, carbon atom always exhibits a valency of four while beryllium and boron exhibit valency of two and three
respectively.
In order to explain the observed valencies of Be, B and C, it is assumed that these atoms acquire excited states before participating
in bonding. In the excited state the electron pair present in 2s-orbital gets unpaired and one of the electrons is promoted to vacant
2p-orbital. For example, the simple excited state configurations of Be, B and C are shown below :
Thus, concept of excitation or promotion of electrons could very well explain the valency of beryllium, boron and carbon as 2, 3
and 4 respectively.
The energy required for excitation is compensated by the energy released during bond formation.
Let us now study the formation of methane from the excited state configuration of carbon.
Carbon uses its four half-filled orbitals for the axial overlap with 1s orbitals of four different H atoms as shown below:
It is quite evident that the four C—H bonds in CH4 (methane) should not be equivalent. Three C—H bonds should be sigma p-s
bonds and one C—H bond should be sigma s-s bond. These bonds should be different so far as their (i) strength, (ii) bond length
and (iii) bond angle are concerned. The experimental facts about methane reveal that all the C—H bonds in the molecule are
equivalent. Their bond length is same i.e., 109 pm and the bond angle H—C—H is 109° 28'. In order to explain these experimental
facts and the equivalency of bonds, a concept called hybridisation is introduced.
According to this concept, valence orbitals of the atom intermix to give rise to another set of equivalent orbitals before the
formation of bonds. These orbitals are called hybrid orbitals or hybridised orbitals and the phenomenon is referred to as
hybridisation. Thus, hybridisation may be defined as the phenomenon of intermixing of atomic orbitals of slightly different
energies of the atom(by redistributing their energies) to formnew set of orbitals of equivalent energies and identical shape.
In case of carbon atoms, one orbital of 2s-level and three orbitals of 2p-level intermix at the time of participation in bonding to
produce four equivalent sp
3
hybridised orbitals. These hybrid orbitals overlap axially with 1s orbitals of four H atoms to form four
equivalent C—H bonds in methane.
Hybridization
(1) In hybridization, orbitals of almost equal energies mix up to give new orbitals of another space and identical energies. If the
energy difference between orbitals is high e.g., 1s and 2p or 2s and 3p, hybridization is not possible. If the energy difference
between orbitals is very little e.g., between 2s and 2p or 3s and 3p, hybridization is possible.
(2) The number of orbitals before and after hybridization remains the same.
(3) The type of hybridization is not fixed for an element; rather, it depends on the chemical environment. In different conditions an
element may hybridize in different ways. For example, carbon shows sp
3
, sp
2
and sp hybridization in CH4 , C2H4 and C2H2,
respectively.
(4) In hybridization, orbitals are used and not electrons, so completely filled, half-filled or empty orbitals can also take part in
hybridization. As far as possible, the electrons are shown in the orbitals in such a way that they remains minimum number of
paired electrons.
(5) The hybridized orbitals mutually repel each other. The repulsion depends on the number of lone pairs and bond-pairs.
(6) Hybridization is an hypothetical concept which is utilised to explain the experimental observations.
(7) Due to force of repulsion, hybrid orbitals try to remain apart at a maximum possible distance. The order of decreasing force of
repulsion is as follows :
lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.
(8) Hybrid orbitals are more directional than the atomic orbitals, hence they form more strong bonds. We know that, greater the
overlapping, stronger is the bond formed. This probability is comparatively higher in more directional orbits. The order of bond
strength is as under :
sp < sp
2
< sp
3
< sp
3
d < sp
3
d
2
(9) The properties of hybrid orbitals taking part in hybridization are in the same ratio in which the orbitals unite. For example, sp
hybrid orbitals have 50% properties of s and 50% properties of p-orbital. sp
3
hybrid orbitals have 25% properties of s and 75%
properties of p-orbital.
(10) Single electron which take part in hybridisation always represent bond.
(11) Electron pair which take part in hybridisation express lone pair of electron.
(12) Single electron which do not take part in hybridisation actually represent bond of system and forms
p —p , d — p or d —d type bonding.
(13) The hybrid orbital has electron density concentrated on one side of the nucleus so one lobe is relatively larger than ot her.
(14) Hybridisation of any species can be find out by the following formula.
Number of hybrid orbital or steric number = number of Bonds + number of lone pair
Types of Hybridization and Shapes of Molecules
There are many different types of hybridisation depending upon the type of orbitals involved in mixing such as sp
3
, sp
2
, sp, sp
3
d,
sp
3
d
2
, etc.
Let us now discuss various types of hybridisation along with some examples with reference to the compounds of carbon, boron
and beryllium.
(i) sp
3
hybridisation. The type of hybridisation involves the mixing of one orbital of s-sub-level and three orbitals of p-sub-level of
the valence shell to form four sp
3
hybrid orbitals of equivalent energies and shape. Each sp
3
hybrid orbital has 25% s-character and
75% p-character. These hybridised orbitals tend to lie as far apart in space as possible so that the repulsive interactions between
them are minimum. The four sp
3
hybrid orbitals are directed towards the four corners of a tetrahedron. The angle between the sp
3
hybrid orbitals is 109.5° (Figure below).
sp
3
hybridisation is also known as tetrahedral hybridisation. The molecules in which central atom is sp
3
hybridised and is linked
to four other atoms directly, have tetrahedral shape. Let us study some examples of molecules where the atoms assume sp
3
hybrid
state.
1. Formation of methane (CH4). In methane carbon atom acquires sp
3
hybrid states as described below :
Here, one orbital of 2s-sub-shell and three orbitals of 2p-sub-shell of excited carbon atom undergo hybridisation to form four sp
3
hybrid orbitals. The process involving promotion of 2s-electron followed by hybridisation is shown in figure below.
As pointed out earlier the sp
3
hybrid orbitals of carbon atom are directed towards the corners of regular tetrahedron. Each of the sp
3
hybrid orbitals overlaps axially with half-filled 1s-orbital of hydrogen atom constituting a sigma bond figure below.
Because of sp
3
hybridisation of carbon atom, CH4 molecule has tetrahedral shape.
2. Formation of ethane (CH
3
—CH
3
). In ethane both the carbon atoms assume sp
3
hybrid state as shown in figure below. One of
the hybrid orbitals of carbon atom overlaps axially with similar orbitals of the other carbon atom to form sp
3
-sp
3
sigma bond. The
other three hybrid orbitals of each carbon atom are used in forming sp
3
-s sigma bonds with hydrogen atoms as described below :
Each C—H bond in ethane is sp
3
-s sigma bond with bond length 109 pm. The C—C bond is sp
3
-sp
3
sigma bond with bond length
154 pm.
(ii) sp
2
hybridisation. This type of hybridisation involves the mixing of one orbital of s-sub-level and two orbitals of p-sub-level
of the valence shell to form three sp
2
hybrid orbitals. These sp
2
hybrid orbitals lie in a plane and are directed towards the corners of
equilateral triangle (Figure below). Each sp
2
hybrid orbital has one-third s-character and two-third p-character. sp
2
hybridisation is
also called trigonal hybridisation. The molecules in which central is sp
2
hybridised and is linked to three other atoms directly
have triangular planar shape.
Let us study some examples of the molecules which involve sp
2
hybridisation.
Formation of boron trifluoride (BF3)
Formation of boron trifluoride (BF3). Boron (5B) atom has ground state configuration as 1s
2
2s
2
, 2p
1
. But in the excited state its
configuration is 1s
2
, 2s
1
, 2px
1
, 2py
1
. One 2s-orbital of boron intermixes with two 2p-orbitals of excited boron atom to form three sp
2
hybrid orbitals as shown in figure below.
The sp
2
hybrid orbitals of boron are directed towards the corners of equilateral triangle and lie in a plane. Each of the sp
2
hybrid
orbitals of boron overlaps axially with half-filled orbital of fluorine atom to form three B-F sigma bonds as shown in figure below.
Becasue of sp
2
hybridisation of boron, BF3 molecule has triangular planar shape.
2. Formation of ethylene (C2H4). Both the carbon atoms in ethylene assume sp
2
hybrid state. In acquiring sp
2
hybrid state, one 2s-
orbital and two 2p-orbitals of excited carbon atom get hybridised to form three sp
2
hybridised orbitals. However, one orbital of 2p-
sub-shell of the excited carbon atom does not take part in hybridisation. The promotion of electron and hybridisation in carbon
atom is shown in figure below.
As already indicated, the three sp
2
hybrid orbitals lie in one plane and are oriented by space at an angle of 120° to one another. The
unhybridised 2p-orbital is perpendicular to the plane of sp
2
hybrid orbitals as shown in figure below.
In the formation of ethylene, one of the sp
2
hybrid orbital of carbon atom overlaps axially with similar orbital of the other carbon
atom to form C—C sigma bond. The other two sp
2
hybrid orbitals of each carbon atom are utilised for forming sp
2
-s sigma bond
with two hydrogen atoms.
The unhybridised p-orbitals of the two carbon atoms overlap sidewise each other to form two p clouds distributed above and bel ow
the plane of carbon and hydrogen atoms figure below.
Thus, in ethylene, the six atoms (bonded by sigma bonds) lie in one plane while the p bond is projected perpendicular to the plane
of six atoms (two C atoms and four H atoms).
In ethylene molecule, the
C = C bond consists of one sp
2
-sp
2
sigma bond and one p bond. Its bond length is 134 pm. C—H bond is sp
2
-s sigma bond with
bond length 108 pm. The H—C—H angle is 117.5° while H—C—C angle is 121°.
(iii) sp-hybridisation. This type of hybridisation involves the mixing of one orbital of s-sub-level and one orbital of p-sub-level of
the valence shell of the atom to form two sp-hybridised orbitals of equivalent shapes and energies. These sp-hybridised orbitals are
oriented in space at an angle of 180° figure below. This hybridisation is also called diagonal hybridisation. Each sp hybrid orbital
has equal s and p character, i.e., 50% s-character and 50% p-character. The molecules in which the central atom is sp-hybridised
and is linked to two other atoms directly have linear shape.
Let us study some examples of molecules involving sp hybridisation.
1. Formation of beryllium fluoride (BeF2). Beryllium (4Be) atom has a ground state configuration as 1s
2
, 2s
2
. In the excited state
one of the 2s-electron is promoted to 2p-orbitals. One 2s-orbital and one 2p-orbitals of excited beryllium atom undergo
hybridisation to form two sp-hybridised orbitals as described in figure below.
The two sp hybrid orbitals are linear and oriented in opposite directions at an angle of 180°. Each of the sp-hybridised orbital
overlaps axially with half-filled orbital of fluorine atom to from two Be—F sigma bonds figure below.
Due to the sp-hybridised state of beryllium, BeF2 molecule has linear shape.
2. Formation of acetylene (CH CH). Both the carbon atoms in acetylene assume sp-hybrid state. In acquiring sp-hybrid state,
one 2s orbital and one 2p-orbital of excited carbon atom (1s
2
2s
1
2px
1
2py
1
> 2pz1) get hybridised to form two sp-hybridised orbitals
figure below.
The two sp-hybrid orbitals of carbon atom are linear and are directed at an angle of
180° whereas the unhybridised p-orbitals are perpendicular to sp-hybrid orbitals and
also perpendicular to each other as shown in figure below.
In the formation of acetylene, carbon atom uses its one of the sp-hybrid orbital for
overlapping with similar orbital of the other carbon to form C—C sigma bond. The
other sp-hybrid orbital of each C atom overlaps axially with 1s-orbital of H atom to
form C—H sigma bond. Each of the two unhybridised orbitals of both the carbon
atoms overlap sidewise to form two bonds. The electron clouds of one bond lie
above and below the internuclear axis whereas those of the other bond lie in front and back of the inter-nuclear axis.
The overlapping of orbitals has been shown in figure below.
The four clouds so formed further merge into one another to form a single cylindrical electron cloud around the internuclear axis
representing C—C sigma bond. It has been shown in figure below.
Gigure. Orbital diagram of BeF2.
Thus, in acetylene molecule,
C C bond consists of one sp-sp s bond alongwith two p bonds. The C C bond length is 120 pm. C—H bond is sp-s sigma bond.
The H—C—C angle is 180°, i.e., the molecule is linear.
Predicting The Hybrid State of Central Atom
Predicting The Hybrid State of Central Atom
The hybrid state of central atoms in simple molecule or in plyatomic ion can be easily predicted by the following considerati ons :
(i) Count the number of atoms/groups surrounding the central atom. Let it be (SA)
(ii) Find the valence electrons of central atom from its atomic number. Let it be (G)
(iii) Find the valency of the central atom in the species. Let it be (V).
V is equal to the number of monovalent groups such as H, Cl, Br, OH ....., etc., or double the numbers of divalent groups such as
O, S, ..... etc., directly bonded to central atom.
(iv) If the given species is anionic, the charge on the ion is represented by (a) whereas as if it is cation, the charge on it represented
by c.
Now SA, G, V, E are related as :
Now, hybrid state, shape, etc., can be prepicted from the following table :
In 1940, Sidgwick and Powell suggested that the shape of a molecule is related to the number of electrons in the outer shell of the
central atom. Irrespective of the fact whether bond pairs or lone pairs occupy the orbitals, the occupied orbitals repel each other and
consequently are oriented in space as far apart as possible. In this position molecule has minimum energy, and therefore, maximum
stability. The shape of the molecule and bond angles can be predicted if the distribution of orbitals about the central atom can be
ascertained.
