chemical banding

8/4/2019 Chemical Banding

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Hydrogen bond is formed when a slightly positive hydrogen atom, attached covalently tostrong electronegative atom A (e.g., F, O or N) is held by a non-bonded electron pair ofanother atom B. The coordination number of hydrogen becomes two and it acts as a bridgingatom between A and B.

Generally hydrogen bond is formed with only F, O and N atoms. Sometimes lesselectronegative atoms such as Cl, S etc., also take part in the formation of hydrogenbond. 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 withstrongly electronegative atom of gain element is known as hydrogen bond. The bond energyof 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 twoor more molecules of the same or different compounds.

(a) Homo-intermolecular hydrogen bonding. It is also termed as self-association. It refersto the association of two or more identical molecules e.g., association in alcohol, associationin water, association in NH3, association in HF etc.

(b) Hetero-intermolecular hydrogen bonding. It refers to the association of two differentspecies. One which donates the lone pair of electrons is called electron donor or hydrogenacceptor, 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.


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(ii) Intramolecular hydrogen bonding. When hydrogen bonding exists within the moleculeit is called intramolecular hydrogen bonding. In such type of hydrogen bonding two groups ofthe same molecule link through hydrogen bond, forming a stable five or six membered ringstructure 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 hydrogenbonding is a pincer like action resembling the closing of a Crab's claw. Some more examplesof intramolecular hydrogen bonding are :

(iii) -Hydrogen bonding. Sometimes -electrons of an olefinic or aromatic system act asproton acceptor in hydrogen bonding. Such type of hydrogen bonding is called -Hydrogenbonding. 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,


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(b) Intramolecular -hydrogen bonding. This occurs between the groups of the samemolecule. Hydrogen of a group (e.g., -- OH) forms hydrogen bond with -electrons of anaromatic 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 neverinvolves more than two atoms.

(ii) Bond energy of hydrogen bond is in the range of 3-10 kcal/mole. Thus hydrogen bond is aweaker bond than a covalent bond (bond energy of a covalent bond is 50-100 kcal/mole). Butit 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 isdifferent from the covalent bond.

(iv) The strength of hydrogen bond depends upon the electronegativity of the atom A towhich hydrogen atom is attached with a covalent bond. As the electronegativity of Aincreases, strength of hydrogen bond increases. Thus HF will form most strong hydrogenbond as fluorine is the most electronegative atom.

(v) Typical hydrogen bond is linear. Angular hydrogen bonds exist in solids or inintramolecular 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 thebasis of hydrogen bonding.

(1) Physical States

i) H2O and H2S. As we have already stated that the ease of formation of a hydrogen bonddecreases as the electronegativity of the atom attached to the hydrogen decreases. Oxygen ismore electronegative than sulphur. There is a considerable hydrogen bonding in H

2O while in

H2S the same is absent. H2O molecules are associated together in which hydrogen atom acts


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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 liquidat ordinary temperature.

In HCl, due to less electronegativity of chlorine atom and its large size, hydrogen bond doesnot exist hence molecules of HCl are not associated as in HF. Therefore, HCl is a gas atordinary 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 increasewith the increasing molecular weights. This is evident from the melting point and boilingpoint curves of IV A group hydrides. [(a) and (b)].

The melting point and boiling point of CH4, SiH4, GeH4, and SnH4 decrease with decreasingmolecular 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 theirgroups.


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The melting point and boiling point of the hydrides of the elements of IVA, VA, VIA andVIIA 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 ofHF, H2O and NH3. The existence of hydrogen bond in these molecules exceptionallyincreases their melting point and boiling point. Boiling point of water is higher than that ofhydrogen fluoride because the extent of association through hydrogen bonding in water ismore than hydrogen fluoride. Since CH4 cannot form hydrogen bond, its melting point andboiling point are the lowest among the hydrides of carbon family.


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(ii) Melting point of ortho- nitrophenol and para- nitrophenol

Intramolecular hydrogen bonds do not involve molecular association and in these, change inphysical properties is quite different than in case of intermolecular hydrogen bonds. For

example the ortho isomers have lower melting point (Intramolecular hydrogen bonding) thanrespective para isomers (Intermolecular hydrogen bonding) as shown above.

Melting points of substituted Nitro compounds

(3) Solubility. Alcohols, glycol, glycerol and sugars are soluble in water due to the formationof hydrogen bond with water molecules. Dimethylether, (CH3)2O, is miscible in water as itcan form hydrogen bond with water molecule but dimethyl sulphide, (CH3)2S, is immiscibleas it cannot form hydrogen bond with water molecules since the electronegativity

of sulphur is less.


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(4) Viscosity. Viscosity increases with the extent of hydrogen bonding in molecules. Theviscosity of water is 10.05 millipoise, methanol 6 millipoise and dimethyl ether 2.3millipoise. Since both H2O and CH3OH are hydrogen bonded the viscosities are high, butwhen there is substitution of second methyl group to produce the non-hydrogen-bondeddimethyl-ether, (CH3)2O, the viscosity drops to a low value. Polyhydroxy alcohols such as

ethylene glycol, CH2OH.CH2OH and glycerol, CH2OH.CHOH.CH2OH which have extensivehydrogen bonding exhibit much higher viscosities.

(5) Molecular weights. The association of two or more molecules by intermolecularhydrogen 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 increasingtemperature 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 morehydrogen 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 hydrogenatoms, two attached with covalent bonds and two with hydrogen bonds. Thus in ice everywater molecule is associated with four other water molecules in tetrahedral pattern. Ice has anopen structure with large empty space due to existence of hydrogen-bonds. When ice melts anumber of hydrogen bonds are broken and the space between water molecules decreases andthe density of water increases, therefore from 0 to 4C, it is maximum. Above 4C theincrease in kinetic energy of the molecules disperse them and the result is that the densitynow decreases with increasing temperature.


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(8) Stability of unusual structures

Generally, the organic compounds with two -- OH groups on the same carbon atom areunstable and soon liberate water molecule. For example :

The stability of compounds like chloral hydrate, CCl3 -- CH(OH)2 can be explained on thebasis of intramolecular hydrogen-bonding.

(9) Chain, Sheet and Three dimensional structures

Hydrogen bonding leads to the formation of chains (HCN, HF, HCOOH), sheet (orthoboricacid, 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 anionof an acid is increased due to intramolecular hydrogen-bonding, the acid strength is greatlyenhanced i.e.,pKa value decreases. The carboxylate ion of o-hydroxy benzoic acid isstabilised 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) thano-salicylic acid. Similarly, o-salicylic acid (pKa = 2.98) is much stronger due tointramolecular hydrogen bonding than its meta (pKa = 4.08) and para (pKa = 4.58) isomers.

Fajan's Rule


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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 occurthrough a process ofdeformation or polarization.

Ionic polarization is favoured by a number of factors which are summarized in the fourFajan's rules.

(1) Small cation _ In a small cation there is greater concentration of positive charge over asmall surface area so it will cause greater deformation of an anion than would he causedby 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 fromthe 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 thanthe 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 anionsare shielded from the nucleus by a number of completely occupied orbitals hence they arereadily 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


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the anion is increased, more covalency will be noticed. This is shown in table from thedecrease 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 forceswhich 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 greaterdeformation and hence covalent character. The decrease in melting points with theincrease of charge of the ions is shown in table to justify this generalization.

Effect of cationic charge upon covalent character

4) Cation with non-inert gas atom structure _ The cations with the inert gas electronconfiguration are most effective in shielding the nuclear charge from its surface while thecations with non-inert gas atom structure have positive fields at their surfaces andconsequently will possess high polarizing powers. Thus the cation should possess anelectronic configuration which is not that of an inert gas.

Hence, if the charge and size are kept nearly constant, cations with 18-electron structurecause 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 ofperiodic table.

Effect of 8 and 18 electronic shell upon the covalent character


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On the basis of these important general rules it becomes possible to predict the type ofbond that a given element is likely to prefer.

Rules(1), (3) and (4) indicate that the cations which are large and have small charge andpossess inert gas electronic configuration should possess least polarizing power e.g., the largealkali 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 havingthe least polarizability. According to rules (2) and (3) the most stable anions are those whichare small and have only a small charge.

The fourth Fajan's rule suggests that, in general, the non-transition elements are more ionicthan the transition elements because their cations have lower polarizing power and so thecations 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.


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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 verticalgroup, with the increase power of the cation will correspondingly decrease. If both moves aremade simultaneously, as in a diagonal relationship then two elements of similar polarizingpower may result, e.g., polarizing power of Be2+ and Al3+ is almost similar as their ionicpotential ( ) are also almost similar [ of Be2+ = 6.48 and of Al3+ = 6.0]. Such elements willform bonds of a similar type in the corresponding compounds. This explains almost identicalchemical 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 orenhancement 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 arelikely to be deeply coloured if the cation is one which has a tendency to polarize anions. Witha 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 isusually an indication of an appreciable amount of nonpolar character or some other unusualstructural 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 ofsilver halides may be considered in which there is polarizing cation and increasingpolarizable 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 withthe addition of ammonia. Increasing covalency from fluoride to iodide is expected anddecreases solubility in water is observed. However, many other factors are involved insolubility 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 particularcation. For example, in alkaline earth carbonates, there is a tendency towards decompositionwith 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


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the group in periodic table, radius of the metal ion increases, ionic potential decreases andstability of the metal carbonate increases as shown in table.

Along the period (from left to right) charge on the metal ion increases, ionic potentialincreases and stability of the metal carbonate decreases e.g., stability of K2CO3 > CaCO3 >CuCO3.

Further, the effect of d electrons is also evident. Cd2+ and Pb2+ are approximately of the samesize as Ca2+ but both CdCO3 and PbCO3 decosmpose at approximately 350C.

(6) Acidic, Basic and Amphoteric character of oxides

If < 2.2 metal oxide is basic e.g., MnO, CrO, Na2O, MgO etc.

If = 2.2 to 3.2 metal oxide is amphoteric e.g., MnO2, CrO2 etc.

If > 3.2 metal oxide is acidic e.g., MnO3, CrO3, Mn2O7 etc.

With the increase in ionic potential of metal ion, polarizing power increases, covalentcharacter of MO bond increases, bond does not dissociate on hydrolysis and the acidiccharacter increases.

Attempt has also been made to correlate the enthalpy of carbonates, sulphates, nitrates andphosphates with increasing charge and size function of the cation. It may therefore be statedin general that size and charge are important factors governing the polarizing power of ionsand consequently, many of their chemical properties.

Results of polarization

Polarization results in increasing covalent character in predominantly ionic bonds which mayin turn affect the melting and boiling points of ionic compounds, their decompositiontemperature, solubility, colour, chemical reactions etc. Possible correlation of these propertieswith polarization has been discussed earlier.

Solubility of compounds

Solubility of diffrent compounds in solvents depends upon many fectors. It is verycomplicated property because many factors control this property simultaneously for example


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(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 absorbedto the surroundings is known as heat of solution.

Whether the heat of solution is positive or negative depends on the nature of solute andsolvent. 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 ionsknown as lattice energy.

(ii) Energy liberated when the ions are solvated or hydrated by solvent there is process ofneutralisation, known as solvation energy or hydration energy. This energy actually takes intoaccount of both solvent. Solvent interaction (energy required to make a hole in water) andsolvent solute interaction. These two are combined together because experimentally they arehard to seprate. Energy change involved in dissolution of a salt represented by born habercycle.

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 ofsolution at in finite dilution. The total energy change from MX(s) to their respective solvatedions is actually independent of the path. Above representation of ionic compound in water isborn-Haber type cycle. In case of KCl, overall process can be imagined to occur in twoconsecutive steps

First step involves vapourizing solid requires energy i.e., work must be done to seprate theions while second step is exothermic because process of solvation where ion-dipole attraction


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liberate energy. When solvation or hydration energy is greater than the lattice energy theoverall process is exothermic so liberation of energy takes place, normally solute is soluble insolvent while if lattice energy is more than hydration energy, process is endothermic and itsdifficult 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 sohydration energy increases. For example

hydration energy of:

In transition series (d-Block) size of ions are almost same due to balancing of shielding effectand nuclear attraction force, therefore their hydration energy is almost constant. Transitionmetal ion can form strong bond with water due to presence of vacant d-orbital so they containhigh 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 decreasesso 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 systemis 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 ionsso stability increases which increases lattice energy for example Al2O3 > Mgcl2 > NaCl.

(c) Bond character _ Ionic compounds have high lattice energy than covalent compound dueto more stability.

Factors affecting solubility of ionic compounds

Solubility of compounds mainly control by solvation energy and lattice energy with thesupport of some minor factor like DEC of solvent, hydrogen bonding etc but it is verydifficult to set a trend in periodic table for solubility of different compounds because bothsolvation-energy and lattice energy decreases with increase in size, thing is which decreasesmore with respect to other, actually this decides the solubility of compound.


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Generally, ionic substances dissolve more readily in solvents composed of moleculescontaining electrostatic dipoles.

Solids are usually crystalline in nature. In the crystal lattice the ions are arranged in asymmetrical fashion and are held in their 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 mustbe provided for the separation of the ions from the crystal than was liberated in building upthe ionic lattice. In other words, the energy of solvation must be of greater magnitude than thelattice energy. (The interaction that takes place when a substance is introduced into a solventis called solvation and the energy change involved in this process is known as the solvationenergy).

Thus, both the solvation energy and the lattice energy affect solubility of ionic compoundsbut in an opposite manner. The important factors affecting solubility of ionic compounds arediscussed below.

1. Nature of solvents. As the DEC of solvents increases solubility of ionic compoundsincreases because it weakens the force of attraction between ions. Force of attraction betweenions given by formula

DEC of water is very high (81), it is the best solvent for ionic compounds. Hydrogenperoxide 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 todissolve ionic compounds.

2. Size of ions. Both lattice energy and solvation energy depends upon size of ions onfollowing manner

where r+ and r- are radii of cation and anion. It is clear from the above relations that decreasein the size of the ions affect the two energies in a similar manner. However, the two energiesare influenced to different extents and the predominating energy affects the solubility todifferent extents.

In case of compounds containing large anions e.g., I-, SO42-, CO3

2-, PO43- etc., the

solubility will decrease with increase of cationic size. For example, in case of sulphates ofalkaline earth metals, the decrease in solvation energy is more rapid than the decrease in


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lattice energy with the increase in the size of cation. The solubility of alkaline earth metalsulphates 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 ismore rapid than the decrease in solvation energy with increasing size of the cation. Thus thesolubility of the fluorides of alkai metals increase as the size of cation increases as givenbelow :

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 ofBe+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 morerapidly than the solvation energy. Thus, solubility of ionic compounds decreases very sharplyas the ionic charge increases.


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Effect of Ionic charge on Solubility

4. Polarization of Anions. Large anions are more polarizable than small anions (Fajan'sRule). Polarization of anion increases covalent character in the molecule, hence decreasessolublility in water. This explains the order of solubility of silver halides in water as givenbelow :

AgF > AgCl > AgBr > AgI

Cations with 18-electrons structure polarize anions more than those with 8-electronarrangements. Polarization of anion decreases the solubility. This explains the followingorder of solubility

KCl > AgCl

NaCl > CuCl

Compound Cationic structure Cationic radius Solubility


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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 NH 3 is moresoluble 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 H 2O. 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-filledp-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 andp-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 :
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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. Theorbitals 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 in 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


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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 (1 s2, 2s2,

2px2, 2py1>, 2pz21). 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 asp -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,bondenergy 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.


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(b) Bond Length

It has already been discussed that two bonded atoms in a molecule remain held up at a certaindistance from each other. They cannot approach too close because it leads to repulsiveinteractions and potential energy of system increases. This minimum distance between thebonded 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 lengthis usually expressed in Angstrom units () or picometers (pm) and it can be determinedexperimentally 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 betweencarbon and chlorine atoms is CCl bond 177 pm.

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 ClCl bond is 237 kJ mol-1 whereas that of HH 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


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(ii)Bond length decreases with the multiplicity of bonds. It is because of the fact that largerthe number of electrons shared by the two atoms greater will be attractive force betweenelectrons and the nuclei and consequently, lesser is the bond length. For example, bondlength of CC bond is 154 pm whereas that of C = C bond is 134 pm. Bond lengths of somecommon 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 todirectional character of atomic orbitals, the covalent bonds in a molecule are oriented inspecified 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, HCH bond angle in CH4molecule is 109 28. Similarly, FBF bond angle in BF3 is 120 and HNH bondangle 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, wefind that the valency of carbon should be equal to 2 because it has only two half-filledorbitals in the valence shell.

6C : 1s2, 2s2, 2px

1, 2py1, 2pz

0

In the same way, the valency of beryllium should be zero and that of boron should be one asis evident from their ground state configurations.

4Be : 1s2, 2s2, 2p0


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5B : 1s2, 2s2, 2px

1, 2py0, 2pz0

Contrary to this, carbon atom always exhibits a valency offourwhile beryllium and boronexhibit valency oftwo and three respectively.

In order to explain the observed valencies of Be, B and C, it is assumed that these atomsacquire excited states before participating in bonding. In the excited state the electron pairpresent 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 ofberyllium, 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:


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It is quite evident that the four CH bonds in CH4 (methane) should not be equivalent. ThreeCH bonds should be sigmap-s bonds and one CH bond should be sigma s-s bond. Thesebonds should be different so far as their (i) strength, (ii) bond length and (iii) bond angle areconcerned. The experimental facts about methane reveal that all the CH bonds in themolecule are equivalent. Their bond length is same i.e., 109 pm and the bond angle HCH is 109 28'. In order to explain these experimental facts and the equivalency of bonds, aconcept called hybridisation is introduced.

According to this concept, valence orbitals of the atom intermix to give rise to another set ofequivalent orbitals before the formation of bonds. These orbitals are calledhybrid orbitals orhybridised orbitals and the phenomenon is referred to as hybridisation. Thus,hybridisation may be defined as thephenomenon of intermixing of atomic orbitals of

slightly different energies of the atom (by redistributing their energies) to form new 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 thetime of participation in bonding to produce four equivalent sp3 hybridised orbitals. Thesehybrid orbitals overlap axially with 1s orbitals of four H atoms to form four equivalent CHbonds in methane.

Hybridization

(1) In hybridization, orbitals of almost equal energies mix up to give new orbitals of anotherspace and identical energies. If the energy difference between orbitals is high e.g., 1s and 2por 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 chemicalenvironment. In different conditions an element may hybridize in different ways. Forexample, carbon shows sp3, sp2 and sp hybridization in CH4 , C2H4 and C2H2, respectively.

(4) In hybridization, orbitals are used and not electrons, so completely filled, half-filled orempty orbitals can also take part in hybridization. As far as possible, the electrons are shownin the orbitals in such a way that they remains minimum number of paired electrons.


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(5) The hybridized orbitals mutually repel each other. The repulsion depends on the numberof lone pairs and bond-pairs.

(6) Hybridization is an hypothetical concept which is utilised to explain the experimentalobservations.

(7) Due to force of repulsion, hybrid orbitals try to remain apart at a maximum possibledistance. 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 strongbonds. We know that, greater the overlapping, stronger is the bond formed. This probabilityis comparatively higher in more directional orbits. The order of bond strength is as under :

sp < sp2 < sp3 < sp3d < sp3d2

(9) The properties of hybrid orbitals taking part in hybridization are in the same ratio inwhich the orbitals unite. For example, sp hybrid orbitals have 50% properties of s and 50%properties of p-orbital. sp3 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 ofsystem 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 onelobe is relatively larger than other.

(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 involvedin mixing such as sp3, sp2, sp, sp3d, sp3d2, etc.

Let us now discuss various types of hybridisation along with some examples with reference

to the compounds of carbon, boron and beryllium.


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(i)sp3 hybridisation. The type of hybridisation involves the mixing of one orbital ofs-sub-level and three orbitals ofp-sub-level of the valence shell to form four sp3 hybrid orbitals ofequivalent energies and shape. Each sp3 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 therepulsive interactions between them are minimum. The four sp3 hybrid orbitals are directed

towards the four corners of a tetrahedron. The angle between the sp3

hybrid orbitals is 109.5(Figure below).

sp3 hybridisation is also known as tetrahedral hybridisation. The molecules in which

central atom is sp3 hybridised and is linked to four other atoms directly, have tetrahedralshape. Let us study some examples of molecules where the atoms assume sp3 hybrid state.

1. Formation of methane (CH4). In methane carbon atom acquires sp3 hybrid states as

described below :

Here, one orbital of 2s-sub-shell and three orbitals of 2p-sub-shell of excited carbon atomundergo hybridisation to form four sp3 hybrid orbitals. The process involving promotion of2s-electron followed by hybridisation is shown in figure below.

As pointed out earlier the sp3 hybrid orbitals of carbon atom are directed towards the cornersof regular tetrahedron. Each of the sp3 hybrid orbitals overlaps axially with half-filled 1s-

orbital of hydrogen atom constituting a sigma bond figure below.


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Because of sp3hybridisation of carbon atom, CH4 molecule has tetrahedral shape.

2. Formation of ethane (CH3CH3). In ethane both the carbon atoms assume sp3 hybridstate as shown in figure below. One of the hybrid orbitals of carbon atom overlaps axiallywith similar orbitals of the other carbon atom to form sp3-sp3 sigma bond. The other three

hybrid orbitals of each carbon atom are used in forming sp3-s sigma bonds with hydrogenatoms as described below :

Each CH bond in ethane is sp3-s sigma bond with bond length 109 pm. The CC bond issp

3-sp3 sigma bond with bond length 154 pm.

(ii)sp2 hybridisation. This type of hybridisation involves the mixing of one orbital ofs-sub-level and two orbitals ofp-sub-level of the valence shell to form three sp2 hybrid orbitals.These sp2 hybrid orbitals lie in a plane and are directed towards the corners of equilateraltriangle (Figure below). Each sp2 hybrid orbital has one-third s-character and two-thirdp-

character. sp2

hybridisation is also called trigonal hybridisation. The molecules in whichcentral is sp2 hybridised and is linked to three other atoms directly have triangular planarshape.


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Let us study some examples of the molecules which involve sp2 hybridisation.

Formation of boron trifluoride (BF3). Boron (5B) atom has ground state configuration as1s2 2s2, 2p1. But in the excited state its configuration is 1s2, 2s1, 2px

1, 2py1. One 2s-orbital of

boron intermixes with two 2p-orbitals of excited boron atom to form three sp2 hybrid orbitalsas shown in figure below.

The sp2 hybrid orbitals of boron are directed towards the corners of equilateral triangle andlie in a plane. Each of the sp2 hybrid orbitals of boron overlaps axially with half-filled orbitalof fluorine atom to form three B-F sigma bonds as shown in figure below.


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Becasue ofsp2 hybridisation of boron, BF3 molecule has triangular planar shape.

2. Formation of ethylene (C2H4). Both the carbon atoms in ethylene assume sp2 hybrid state.

In acquiring sp2 hybrid state, one 2s-orbital and two 2p-orbitals of excited carbon atom get

hybridised to form three sp2

hybridised orbitals. However, one orbital of 2p-sub-shell of theexcited carbon atom does not take part in hybridisation. The promotion of electron andhybridisation in carbon atom is shown in figure below.

As already indicated, the three sp2 hybrid orbitals lie in one plane and are oriented by space atan angle of 120 to one another. The unhybridised 2p-orbital is perpendicular to the plane ofsp

2 hybrid orbitals as shown in figure below.

In the formation of ethylene, one of the sp2 hybrid orbital of carbon atom overlaps axiallywith similar orbital of the other carbon atom to form CC sigma bond. The other two sp2hybrid orbitals of each carbon atom are utilised for forming sp2-s sigma bond with two

hydrogen atoms.


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The unhybridisedp-orbitals of the two carbon atoms overlap sidewise each other to form twop clouds distributed above and below 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 isprojected perpendicular to the plane of six atoms (two C atoms and four H atoms).

In ethylene molecule, the

C= C bond consists of one sp2-sp2 sigma bond and one p bond.Its bond length is134 pm.CHbond is sp2-ssigma bond with bond length108 pm.The HCH angle is117.5while HCC angle is121.

(iii) sp-hybridisation. This type of hybridisation involves the mixing of one orbital ofs-sub-level and one orbital ofp-sub-level of the valence shell of the atom to form two sp-hybridisedorbitals of equivalent shapes and energies. These sp-hybridised orbitals are oriented in spaceat an angle of 180 figure below. This hybridisation is also calleddiagonal hybridisation.Each sp hybrid orbital has equal s andp 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 stateconfiguration as 1s2, 2s2. 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 formtwo sp-hybridised orbitals as described in figure below.


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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 tofrom two BeF 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-hybridstate. In acquiring sp-hybrid state, one 2s orbital and one 2p-orbital of excited carbon atom

(1s2

2s1

2px1

2py1

> 2pz1) get hybridised to form two sp-hybridised orbitals figure below.

The two sp-hybrid orbitals of carbon atom are linear and aredirected at an angle of 180 whereas the unhybridisedp-orbitals are perpendicular to sp-hybrid orbitals and alsoperpendicular to each other as shown in figure below.

In the formation of acetylene, carbon atom uses its one of thesp-hybrid orbital for overlapping with similar orbital of theother carbon to form CC sigma bond. The other sp-hybrid

orbital of each C atom overlaps axially with 1s-orbital of Hatom to form CH sigma bond. Each of the two


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unhybridised orbitals of both the carbon atoms overlap sidewise to form two bonds. Theelectron clouds of one bond lie above and below the internuclear axis whereas those of theother 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 cylindricalelectron cloud around the internuclear axis representing CC sigma bond. It has been shownin 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 120pm. CHbond is sp-ssigma bond. The HCCangle is 180, i.e., the molecule islinear.

Predicting The Hybrid State of Central Atom

The hybrid state of central atoms in simple molecule or in plyatomic ion can be easilypredicted by the following considerations :

(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 thenumbers of divalent groups such as O, S, ..... etc., directly bonded to central atom.


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(iv) If the given species is anionic, the charge on the ion is represented by (a) whereas as if itis 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 :


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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 pairsor lone pairs occupy the orbitals, the occupied orbitals repel each other and consequently areoriented in space as far apart as possible. In this position molecule has minimum energy, andtherefore, maximum stability. The shape of the molecule and bond angles can be predicted ifthe distribution of orbitals about the central atom can be ascertained.

chemical banding

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