Lewis -Kossel theory

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Synopsis

Molecule: An atom or group of atoms which can exist independently at normal conditions is called a 'Molecule'.

Chemical Bond: The attractive force that holds two constituent atoms (or) oppositely charged ions together in different chemical species is called 'Chemical Bond'.

Valency: A particular number of atoms are present per molecule depending on the combining capacity of the constituent elements. This combining capacity is called 'Valency'.

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Why do atoms combine?

Atoms are less stable and more energetic hence, they form molecules to lose energy by participating in a c & b. Formation of bond is accompanied by decrease in potential energy. So, chemical bond formation is always exothermic. In a chemical bond, both attractive and repulsive forces exist in equilibrium. H+H→H2+ 434.72KJ

Cl+Cl→Cl2+ 239.1KJ

Also, a completely filled electronic configuration is highly stable. Hence, all other elements participate in chemical bond in order to acquire stable

configuration.

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KOSSEL - LEWIS THEORY

Also called Electronic Theory of Valency. Proposed on the basis of Bohr's Atomic theory. Valence shell: The outer most energy level in an atom

Valence electrons: The electrons in outer most energy level Core electrons: The electrons in the inner energy levels Kernel: Nucleus plus inner electrons is called kernel.

Lewis symbols: Valence electrons are represented by dots and this representation is known as Lewis symbols Ex: : ∙Li, ∙Be∙, ∙B˙∙ etc...,

Stability: Octet configuration is most stable. All elements try to acquire eight electrons in their outer most orbit (octet configuration). Duplet configuration: 'He' has only two electrons in the valence shell, it is highly stable and chemically inert (Duplet Configuration). Examples

Molecules with incomplete octet LiCl, BeH2, BCl3.

Molecules with expanded octet PF5, SF6, H2SO4

Covalent and Electrovalent Bond Atoms attain stable electronic configuration (Duplet and Octet configuration) either by transfer (or) by sharing of electrons. Electrovalent bond: The Number of electrons transferred is called Electrovalency and resultant bond formed is called Electrovalent bond. Covalent bond: The number of electrons shared is called Co-Valency and the resultant bond formed is called Covalent Bond.

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DIAGRAM

Lewis structure of molecules:

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Lewis structure of molecules: Formation of bonds

Electronegative element + Electronegative element → Covalent bond

Electropositive element + Electronegative element → Ionic bond

Covalent and Ionic bonds are extreme representation, though one type generally predominates. In most substances, the bond type is some where between these extreme forms.

Types of bonds

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Ionic or Electrovalent bond

The strong electrostatic attraction between two oppositely charged ions which are formed due to transfer of electrons from one atom to another is called Ionic Bond (or) Electrovalent Bond.

∙ Generally, Ionic Bond is formed between a Metal and a Non metal.

∙ The atom which gains electrons turns into an anion,and which loses

electrons turns into a cation. ∙ Formation of Ionic bond is a redox process because one atom undergoes

oxidation other one undergoes reduction.

Mg→Mg+2 + 2e−

2F + 2e−→2F−

Mg+2 + 2F−→MgF2(or)Mg2+2F−

∙ The maximum electrovalency in the formation of Ionic Bond is 3.

∙ In MgF2 Electrovalency of Mg = 2, F=1

∙ In Na2O electrovalency of Na = 1, O = 2

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Ionic Bond: Favourable Factors

Cation Formation Low ionization potential Potassium (IP=495.57 kJ/mole) forms the cation more readily than sodium (IP=519.82 kJ/mole) Low charge on the ion

Al+3<Mg+2<Na+

Large atomic size (K+>Na+

Cs+ ion has the biggest cation among alkali metal ions. Its compounds are

more ionic than the corresponding compounds of other alkali metals Formation of cation with Inert gas configuration.

Of the two cations Zn++(2, 8, 18) and Ca++ (2,8,8), Ca++ is more readily

formed and it is more stable than Zn++ and gives compounds with more ionic

character.

Anion Formation: High Electro-negativity and Electron affinity.

F2 > O2 > N2

Small atomic size F−>Cl−>Br−>I−

Low charge on the ion. Note: Formation of anion carrying less negative charge is easy. No Bond is 100% ionic in nature. It has some percentage of covalent character which is explained on the basis of Fajan's rule.

LiCl is more covalent in nature and so dissolves in non-polar solvents like

alcohol or ether.

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Fajan's Rules

Fajans rules are used to predict whether a chemical bond will be covalent (or) ionic. Based on the phenomenon of Polarisation They are used to know the relative ionic natures. Depends on the charge of the cation and anion, and the relative size of cation and anion. a) Ionic nature & Size

∝ size of cation

∝ 1/ size of anion b) Ionic nature & Charge

∝ 1/ charge on cation

∝ 1/ charge on anion Cations with inert gas configuration form ionic compounds while those cations

with psuedo inert gas configuration favour covalent bond formation.

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Fajan's Rules: Determining Nature of Bond

A compound is more ionic (or) less covalent if it contains a) Large cation and small anion.

CaF2 > CaCl2 > CaBr2 > CaI2

LiF < NaF < KF < RbF <CsF

b) Cation carrying less positive charge and anion carrying less negative charge. NaF>MgF2>AlF3

c) Cation having inert gas configuration.

NaCl>CuCl A compound is less ionic (or) more covalent if it contains. a) Small cation and large anion. b) Cation carrying high positive charge and anion carrying high negative charge. Cation having pseudo inert gas configuration. In a binary compound AB, if the electronegativity difference between the elements A and B Electronegativity difference = 1.7, AB is 50% ionic Electronegativity difference > 1.7, AB is an ionic compound Electronegativity difference < 1.7, AB is a covalent compound. Exception: HF is covalent compound, even though electronegativity

difference is 1.9

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Crystal structures of ionic compounds

Lattice Arrangement.The arrangement of ions in the crystal of anionic solid Unit Cell The smallest part of the crystal of an ionic compound that represents its lattice arrangement Crystal of an ionic solid is a combination of different unit cells. Lattice points are the points that represents ions packed in a crystalline substance . In different cubic unit cells there are mainly four kinds of lattice points as in table.

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Co-ordination number

Co-ordination Number is the number of oppositely charged ions, that surround an ion at nearest possible distances in an ionic crystal Example: Co-ordination number of Na+ and Cl− in NaCl is six. Co-ordination number of Cs+ and Cl− in CsCl is eight.

The co-ordination number depends on radius ratio of the ionic crystal.

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Energy changes in ionic bond formation

Lattice energy: The amount of energy released when the required number of oppositely charged gaseous ions present at infinite distances come close and form one mole of ionic solid. The ions are more stable in crystal lattice than in free state. Relation with radii: Lattice energy is inversely proportional to the sum of the radii of cation and anion.

Lattice energy ∝1(rc+ra) Calculation of Lattice energy: Some terms used Born-Haber cycle is used to calculate the lattice energy. Born-Haber cycle is based on Hess law of constant heat summation. Hess law is a deduction from the law of conservation of energy. Hess law of constant heat summation the total amount of heat energy released or absorbed in a process remains constant, whether the process occurs in one stage or in several stages. Hess law is applicable to both physical changes and chemical changes. Heat of formation:

The heat energy released when one gram mole of NaCl(s) is formed by the

combination of Na(s) and Cl2(g) is known as the heat of formation of NaCl(s). It

is represented by Δ H.

Heat of Sublimation: The heat energy absorbed when one gram mole of Na(s) changes to Na(g) is

known as the heat of sublimation of Na(s). It is represented by S.

Ionisation energy:

The heat energy absorbed when one mole of Na(g) atoms change

to Na+(g) ions is known as the ionisation energy of Na(g). It is represented by I.

Heat of dissociation:

The heat energy absorbed when one mole of Cl2(g) molecules change to two

moles of Cl(g) atoms is known as heat of dissociation of Cl2(g). It is

represented by D. Electron Affinity: The heat energy released when one mole of Cl(g) atoms change to Cl−(g) ions

is known as electron affinity of Cl(g) It is represented by E. Electron affinity is

also called electron gain enthalpy (ΔegH ) with opposite sign.

Electron affinity = -electron gain enthalpy.

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Calculation of Lattice Energy

Lattice Energy:

The lattice energy of NaCl(s) is the heat energy released when one mole

of NaCl(s) is obtained by the combination of Na+(g) and Cl−(g) ions. It is

represented by U. Calculation: The lattice energy of NaCl(s) can be calculated by using the following formula

−ΔH = +S+I+D2−E+U

(The negative sign indicates the heat released and positive sign indicates the heat absorbed) U = - 769.96 KJ /mole ∙ Lattice energy,U = −NAZ+.Z−.e2r + NBe2rn

attractive repulsive force force [N = Avogadro number; A= Madelung constant, depends on geometry of the crystal varying between 1.763(CsCl) and 2.519(fluorite structure Z(+) & Z− are the charges on the positive and negative ions; e = charge on an electron; r = inter-ionic distance; B= repulsion coefficient depends on the structure and is proportional to the number of nearest neighbours; n= Born exponent] The equilibrium distance between ions is determined by the balance between the attractive and repulsion terms. At equilibrium,

dUdr = NAZ+Z−e2r204πε0−nNBrn+10 = 0

Rearranging this B = AZ+Z−e2rn−104πε0n

Substituting in the expression of U U=NAAZ+Z−e2rπε0 r0(1+1n)

This equation is called Born-Lande equation.

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Factors affecting lattice energy

The force of attraction between two oppositely charged ions is given by well known coulomb's law as, F = q1q2(4πε0)r2

where q1 and q2 are the charges of ions in coulombs,

4πε0 is the permitivity factor and r is the distance between ions .

Hence, lattice energy depends upon i) Charge on ions: Higher the charge on ions, greater will be the force of attraction between them and greater will be the lattice energy. The order of lattice energy for different solids is as follows- bi - bivalent solids > uni-bivalent solids > uni-univalent solids MgO > CaCl2 > NaCl ii) Size of ions: Smaller the size of the ions, lesser will be the distance between them, greater will be the force of attraction and so, greater will be the

lattice energy.

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Polarising Power of Cation

It is generally represented by ϕ and also known as ionic potential charge

density. It can be represented as ϕ = chargeoncationradiusof cation

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DIAGRAM

Polarisation

When cation and anion come closer to each other,the electron cloud of anion is attracted towards the cation, some partial sharing of electrons take place, the anion is distorted and the effect is known as polarisation. More the effect of polarisation, more is the sharing of electrons and more is

the covalent character of ionic bond.

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Properties of ionic compounds

Ionic compounds exist as solids. Since electrostatic forces of attraction are extending in all directions, each ion tends together as many opposite kind of ions around itself. Ionic compounds have high melting points and high boiling points. (Melting point of NaCl = 803∘C ) Ionic compounds dissolve in polar solvents.

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Properties of Ionic compounds

Solubility: They are generally soluble in polar solvents (having high value of di-electric constant). The solubility of ionic compounds decreases with increase in covalent character of ionic compounds. It is also governed by Lattice energy: More the lattice energy, lesser is the solubility, eg. sulphates and phosphates of Ba and Sr are insoluble in water due to high lattice energy. Heat of hydration: More the heat of hydration, more is the solubility. eg. AlCl3 though covalent in nature is soluble in water due to high value of heat of hydration Electrical Conduction: Ionic compounds are good electrical conductors in molten state (or) in aqueous solutions. Ionic compounds undergo chemical reactions quickly in aqueous solutions.

NaCl+AgNO3→AgCl+NaNO3

white ppt. Isomerism: Ionic compounds do not exhibit space isomerism because ionic bond is a non-directional bond. Melting and Boiling points: The melting and boiling points of ionic compounds are very high due to the presence of strong electrostatic forces of attraction between the ions. If the anion is common in ionic compounds, the ionic compound with its cations having high atomic weight will have more melting point. Example: Melting point of BaCl2 is very high compared with BeCl2. Melting point of CsF is very high compared with LiF. If cation is common in the ionic compounds, compound with high lattice energy will have high melting point and the compound with low lattice energy will have low melting point. Example:

Melting point of NaF is more than that of NaI.

Melting point of CaF2 is more than that of CaI2 .

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Isomorphism

NaF and MgO are isomorphous due to similar electronic structure

Na+ F− Mg++ O−−

2, 8 2,8 2,8 2,8 Similarly K2S and CaCl2 are isomorphous

K+ S2− K+ Cl− Ca++ Cl−

2, 8, 8 2, 8, 8 2, 8, 8 2, 8, 8 2, 8, 8 2, 8, 8

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Variable electrovalency

It is due to unstable atomic core and inert pair effect. When valence electrons from a metal atom are removed, it leaves behind the atomic core or atomic kernel. The latter with 2 or 8 electrons are stable. In case of transition elements, the atomic cores are not very stable. They may lose one or more electrons thus, exhibiting variable electrovalency eg. iron and copper etc.

Fe ( 2,8,14,2) 1s2, 2s2p6, 3s2p6d6, 4s2

Fe++ ( 2,8,14) 1s2, 2s2p3, 3s2p6d6

Fe+++ ( 2,8,13) 1s2, 2s2p6, 3s2p6d5

The energy difference between 4s and 3d subshell is not very high.

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Covalent Bond

Covalent Bond is the bond formed between two atoms by the mutual sharing of their valence electrons. Only unpaired electrons present in the atom involve in covalent bonds. One unpaired electron involves in one covalent bond only and so the number of covalent bonds in which an atom involves is equal to the number unpaired electrons present in it. The atom forms less number of covalent bonds in its ground state. The atom forms more number of covalent bonds in its excited state. Example:- Phosphorus atom has three unpaired electrons in its ground state. So, it can form only three covalent bonds in its ground state. Phosphorus atom has five unpaired electrons in its excited state. So, it can

form five covalent bonds in its excited state. Single bond is formed by the mutual sharing of one pair of electrons between two atoms. Single bond is represented as '_____' Double bond (Dicovalent bond) is formed by the mutual sharing of two pairs of electrons between two atoms. Triple bond (Tricovalent bond) is formed by the mutual sharing of three pairs

of electrons between two atoms. It is represented as ≡

Double and triple bonds are called multiple bonds. Pure covalent bond is formed by the sharing of electron pairs between two like atoms. The word like atoms stands for atoms of the same element (or) atoms having same electro-negativity value. Pure covalent bond is present in H2, Cl2, O2, N2, P4, S8 etc. Polar covalent bond is formed by the mutual sharing of electron pairs between two dissimilar atoms.The word dissimilar atoms stands for atoms of different elements (or) atoms having different electro - negativity values. Polar covalent bond is present in HF, HCl, ICl, H2O, CO2, SO2, BeCl2, SO3, BCl3, NH3, CH4, PCl5, SF6 etc. The unpaired electrons present in the valence shell of an atom involve in covalent bonds, the remaining electron pairs present in its valence shell are called lone pairs of electrons (or) non bonded pairs of electrons. The lone

pairs of electrons present in the atom involve in co-ordinate covalent bonds.

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Favourable conditions for the formation

1) Small size of cation 2) Large size of anion 3) Large charge on cation and anion 4) Generally, a covalent bond is formed between 2 atoms of similar E.N

values

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Formal charge

Formal charge is a factor based on a pure covalent bond formed by the sharing of electron pairs equally by neighboring atoms. Formal charge may be regarded as the charge that an atom in a molecule would have if all the atoms had the same electro-negativity. It may or may not approximate the real ionic charge. In case of polyatomic ions, the net charge is possessed by the real ion as a whole and not by any particular atom. It is, however, feasible to assign a formal charge on an atom in a polyatomic molecule or ion. where

QF = [NA−NM ] = [NA−NLP−1/2NBP]

NA = number of electrons in the valence shell in the free atom

NM = number of electrons belonging to the atom in the molecule

NLP =number of electrons in unshared pairs, i.e.number of electrons in lone

pairs

NBP = number of electrons in bond pairs, respectively.

Qf = Formal charge

(or) F.C = V - L - B/2 V = No. of valence electrons L = No. of lone pairs of electrons B = No. of electrons involved in bonding Example of PH3 molecule : The Lewis dot formula of PH3 is as above Formal Charge of P:

QF = [NA−NM ] = [NA−NLP−1/2NBP]

{5−2−1/2(6) } = {5−5} = 0

Formal Charge of H:

QF = [NA−NM ] = [NA−NLP−1/2NBP]

{1−0−2/2} = 0

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Properties of Covalent Compounds

Exist as solids, liquids and gases. Exhibit space isomerism because covalent bond is a directional bond. Melting and boiling points: Low melting and boiling points. The melting and boiling points of covalent compounds are proportional to their molecular weights. Melting and boiling of Covalent Polymers: The melting and boiling points of covalent polymers are very high. Examples for covalent polymers: Diamond,Graphite, Boron, Silicon, Silica, Boron nitride, Boron Carbide and Silicon carbide. Electrical conductance: They are not electrical conductors either in molten state or in aqueous solutions. Covalent substances do not contain free electrons or free ions. So, they are not electrical conductors. However, Graphite contains free electrons. So, it is a good conductor of electricity. HCl, HBr, HI, HNO3 and H2SO4 etc. are covalent compounds. They have both

ionic and covalent character. The aqueous solutions of these substances are good electrical conductors because they contain free ions. Chemical Characteristics: Covalent substances involve in chemical reactions slowly because molecules take part in the reactions. CCl4 ,CHCl3 ,CHCl2 etc. cannot give white precipitate with AgNO3 in aqueous solutions because they do not contain chloride ions as they are covalent compounds. Solubility: Dissolve in non-polar solvents. The solubility of a covalent substance is more in a solvent having high dielectric constant and less in a solvent having low dielectric constant. Dielectric Constant: The dielectric constant of a solvent is a measure of its capacity to break the covalent molecules into ions. Polar solvents have high dielectric constants. Non polar solvents have low dielectric constants. Dielectric constant is high for solvents like water, ammonia, hydrofluoric acid etc. Dielectric constant is low for solvents like Benzene, Carbon disulphide, carbon tetrachloride etc. Vander Waals Force of Attraction: The force of attraction between the molecules is known as Vander Waals force of attraction. It is of the order of 1K.cal/mole. Maximum in solids and minimum in gases.

It is proportional to the molecular weight of the substance.

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Ionic character of a covalent bond

Based on Electro-negativity Difference: a) Pauling equation:

% ionic character = 100[1−exp14(XA−XB)]

b) Hannay and Smith equation:

% ionic character = 16(XA−XB)+3.5(XA−XB)2

Where XA and XB are the electronegativities of atoms.

Based on Dipole Moment: % ionic character = observed dipole moment of bondcalculated dipole moment of bond×100

Eg: Calculate the percentage ionic character of HCl molecule. The bond length is 1.275 Ao and observed dipole moment is 1.03 D. Solution For 100% ionic character the dipole moment should be

μ = e×d = 4.8×10−10×1.275×10−8 = 6.12 D

The observed dipole moment is = 1.03 D

∴ % ionic character = 1.036.12×100 = 16.83%

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Co-ordinate covalent bond (or) dative bond

Co-ordinate covalent bond is also known as dative bond or semi polar bond. Co-ordinate Covalent Bond is the bond formed between two atoms by the sharing of an electron pair which is contributed by only one of the two atoms.

Co-ordinate Covalent bond is represented by an arrow (→ ) the head of which

is close to the atom which accepts the electron pair. For all practical purposes, the co-ordinate covalent bond is treated as a single covalent bond. The formation of co-ordinate covalent bond occurs only after the formation of covalent bond. Examples: Co-ordinate covalent bond is present in molecules like:

N2O, O3, H2O2, BF3, N2O4, N2O5, CO, F3B−NH3, B3N3H6, Al2Cl6

Ions like: H3 O+, NH+4, NO−3, BF−4, N2H+5, C6H5NH+3

∙ In a hydrated cation the bond between water molecule and cation is dative

bond. Every water molecule involves in one dative bond only.

[Al(H2O)6]+++ ion contains 6 dative bonds.

[Be(H2O)4]++ ion contains 4 dative bonds.

The bond between ligand and metal atom (or)metal ion in complex compounds is dative bond.

∙ [Cu(NH3)4]SO4 contains 4 dative bonds.

∙ [Ni(H2O)6]Cl2 contains 6 dative bonds.

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Co-ordinate covalent bond (or) dative bond

Earlier, to account for the octet of the central atom there is a practice of writing dative bonds in the structures of SO2, SO3, Cl2O6, Cl2O7, H2SO4, H3PO4, SO2−4, PO3−4, ClO−4, ClO−3 ,ClO−2, SO2−3 etc. But as per the modern concepts, the dative bonds in the structures of all these examples must be shown as double bonds. ∙ In any anion, negative charge is present on more electronegative

atom.(except in C¯¯¯N ion)

∙ In any cation, positive charge is present on less electronegative atom (except

in NO+). Covalency: The total number of covalent bonds and dative bonds made by an atom is known as its covalency. Example: Covalency of Nitrogen in NO−3 ion is 4. Covalency of sulphur in SO2−4 ion is 6 Covalency of Oxygen in H3O+ ion is 3. According to Sidgewicks covalency rule, maximum covalency of the elements present in Second period is four Third period is six Fourth period is six. Fifth, sixth and seventh periods is eight. Central Atom: The atom common in all the bonds is known as central atom. The central atom obeys octet rule in NH3, H2O, CCl4, CH4, SiH4, CHCl3 etc. The octet is not complete in the valency shell of central atom in BeCl2, BCl3, AlCl3 etc. Expanded octet is present in the valency shell of central atom in PCl5, AsCl5, SO2, SO3, SF6, IF7 etc. Paramagnetism: A molecule or ion is paramagnetic, if it contains odd number of electrons or unpaired electrons. NO, NO2, ClO2, O2 molecules are paramagnetic because they contain unpaired electrons. A molecule (or) ion contains odd electron bond, if it is paramagnetic. This is an old concept. H+2 ion contains single electron bond.

NO, NO2, ClO2 and O2 molecules contain three electron bond in them.

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Hydrogen bond

Hydrogen bond is a weak electrostatic force between positively charged hydrogen atom of a polar molecule and a highly electronegative atom carrying negative charge. The highly electronegative atom may be present in the same molecule (or) in a different molecule. Represented by a dotted line (----) The length of the hydrogen bond varies from 1.76 A∘ to 2.75 A∘and depends on the substance

The energy of hydrogen bond varies from 2 to 10 K.cals/mole or 10-50 kJ / mole. It is weaker than covalent bond but stronger than Van der waals force of attraction. Common examples Most electronegative atoms like Fluorine, Oxygen, Nitrogen only can involve in hydrogen bond. Chlorine atom very rarely involves in hydrogen bond. Intramolecular Hydrogen bond: Present in the same molecule. Leads to ring formation (or) chelation. Examples: o- Chlorophenol, o- Nitro-phenol, o- Nitro-aniline, o- Hydroxy benzaldehyde (Salicylaldehyde), o- Hydroxy benzoic acid ( Salicylic acid), Less water soluble, are steam volatile, have low boiling points. Intermolecular Hydrogen bond Present between different polar molecules. Examples: Water, ammonia, hydrofluoric acid, ortho phosphoric acid, ortho boric acid, p-nitro phenol, p- hydroxy benzaldehyde, p-hydroxybenzoic acid, Primary

alcohols (CH3OH,C2H5OH ), fatty acids (HCOOH,CH3COOH), and

primary amines, (F.......H−F)−

exist as associated molecules.

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Hydrogen Bonds: Associated and Normal Liquids

Liquids having hydrogen bonds between their molecules are called associated liquids. E.g. Water, ammonia, hydrofluoric acid, methylalcohol and ethyl alcohol etc., are associated liquids. Liquids which do not contain hydrogen bonds between their molecules are called normal liquids. E.g. Benzene, carbon disulphide, carbon tetrachloride, acetone, ether, bromine, nitrobenzene etc., are normal liquids. Associated liquids have higher boiling points than normal liquids. Liquids having very low boiling points are called volatile liquids. Eg: ether, acetone, benzene etc. Boiling Points of Associated Liquids: The boiling point of an associated liquid depends on strength of hydrogen bond present in it and number of hydrogen bonds present in one mole of it.

The order of boiling points

NH3>PH3<AsH3<SbH3

H2O>H2S<H2Se<H2Te

HF>HCl<HBr<HI

CH4<SiH4<GeH4<SnH4

The boiling points of NH3, H2O, HF are more than those of PH3, H2S,

and HCl respectively because intermolecular hydrogen bonds are present

in NH3, H2O, HF. Hydrogen bonds are not present in PH3, H2S, HCl.

The boiling point gradually increases from CH4 to SnH4 because there are no

hydrogen bonds in CH4, SiH4, GeH4, SnH4

HF Exception: The boiling point of water (100∘C) is more than that of hydrofluoric acid (19.4∘C though the hydrogen bond in HF is very strong. This is due to the presence of more number of hydrogen bonds in one mole of water than in one mole of HF.

Hydrofluoric acid exists as H6F6 / (HF)6 molecules in vapour state. Water

exists as only H2O molecules in vapour state, so heat of vaporisation of

hydrofluoric acid is less than that of water. Ammonia and Water: The boiling point of water is more than that of ammonia though one mole of ammonia contains more hydrogen bonds than one mole of water.This is due

to the presence of very strong hydrogen bonds in water than in ammonia.

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Hydrogen Bonds: Impact on Structure

Formic Acid: Dimer formation The molecular weight of formic acid (or) acetic acid determined by using its solution in a non-polar solvent like benzene is twice the expected value. This is due to the dimerisation of acid molecules in the solution. The dimer formation takes place with the help of hydrogen bonds. Ice In ice, every water molecule involves in four hydrogen bonds. The ice is a tetrahedral three dimensional polymer. DNA

The two helical strands in the DNA molecule are joined by Hydrogen bonds.

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Hydrogen Bonds: Water Solubility and Hydrated Salts

Solubility of Covalent Substances: Covalent substances like Glucose, Urea, Sugar,Ammonia, Ethyl alcohol etc. dissolve freely in water because they form hydrogen bonds with water. Substances having intermolecular hydrogen bonds are highly water soluble. They have high boiling points and they are not steam volatile. Hydrated Cations and Anions: In a hydrated cation, the bond between water molecule and cation is dative bond. In a hydrated anion, the bond between water molecule and anion is hydrogen bond. Hydrated Salts: Hydrated salts having even number of water molecules: In a hydrated salt having even number of water molecules generally the water molecules are bonded only to the cation. Examples: BeSO4.4H2O ; AlCl3.6H2O

MgCl2.6H2O ; FeCl3. 6H2O

CaCl2.6H2O ; Fe(NO3)3.6H2O

Hydrated salts containing odd number of water molecules: In a hydrated salt having odd number of water molecules, one water molecule is bonded to the anion and the remaining water molecules are bonded to the cation. Examples:

CuSO4.5H2O ; NiSO4.7H2O

MgSO4.7H2O ; FeSO4.7H2O

In some hydrated salts (Eg: BaCl2.2H2O;SrCl2.2H2O etc) the water molecules

are not bonded either to the cation or to the anion.

VBT and VSEPR theory

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VALENCY SHELL ELECTRON PAIR REPULSION THEORY (VSEPR THEORY)

A = Central atom in the compound. B = Atom linked to the central atom. E = Lone pairs of electron. VSEPR theory was proposed to explain the shapes of molecules and ions. Localised Orbital is the orbital which contains the bonded pair of electrons. Central Atom is atom which forms maximum number of bonds. Orientation: The lone pairs of electrons and bonded pairs of electrons on central atom are oriented in such away, that there is least repulsion between them. The repulsion between electron pairs on central atom follows the order. l.p. - l.p. > l.p. - b.p. > b.p. -b.p. In VSEPR theory, the number of electron pairs in single bond (or) double bond (or) triple bond (or) dative bond is counted as only one pair because all the electron pairs in the same bond are oriented in the same direction. The number of lone pairs of electrons (l.p) and bond pairs of electrons (b.p) on central atom determines the shape of molecule or ion.

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Valence Bond Theory (VBT)

V.B.T was proposed by Heitler and London to explain the shapes of covalent molecules, their bond angles and bond lengths and was extended by Pauling and Slater to explain the directional nature of covalent bonds. Covalent bond is formed by the overlapping of two orbitals having unpaired electrons. The two orbitals involving in overlapping must belong to two different atoms and the electrons present in them must have opposite spins. If two atoms A and B approaching each other having nuclei NA and NBand electrons present in them are represented by eA and eB Attractive forces: (i) Nucleus of one atom and its own electron i.e.NA - eA and NB - eB. (ii) Nucleus of one atom and electron of other atom i.e. NA - eB, NB - eA Repulsive forces: (i) electrons of two atoms like eA - eB. (ii) Nuclei of two atoms NA - NB. Dative bond is formed by the overlapping of an orbital having a pair of electrons and a vacant orbital.The greater the extent of overlapping, the stronger is the bond formed. The overlapping of atomic orbitals follows the order. p-p >s-p > s-s sigma bond: A covalent bond formed with its electron cloud concentrated along the inter-nuclear axis and having a cylindrical symmetry is known as sigma bond (σ).Linear overlapping of atomic orbitals results in the formation of

sigma bond (σ). pi bond: A covalent bond formed by the side wise overlap of atomic orbitals of atoms

already bonded through a σ bond and in which the electron cloud is present

on either side of the inter nuclear axis is known as a pi bond (π).

Lateral overlapping of atomic orbitals results in the formation of pi bond (π).

π bond is formed only after the formation of σ bond.

Any type of orbitals can involve in σ bond formation.

Only p or d - orbitals can involve in π bond information.

Single bond is equal to one σ bond.

Double bond is a combination of one σ bond and one π bond.

Triple bond is a combination of one σ bond and two π bonds.

Strength of the bonds follows the order σp−p>σs−p>σs−s>π

σ bond between 1s orbitals is exceptionally stronger.

Strength of the bonds follows the order

triple bond > double bond > single bond.

3

Dipole moment

If a covalent bond is formed between two disimilar atoms, eg. A and B, one of the atoms ( A or B ) must be more electronegative than the other . If A is more electronegative than B, the shared pair of electrons is drawn near A leaving a positive charge on B and hence making the molecule dipolar ( A− B +) . The percentage of polar character is given in terms

of dipole moment ( μ ).

The dipole moment is defined as the product of electric charge e and the

distance d between the two atoms of a polar molecule.( μ = e ×d ) Dipole

moment is a vector quantity with direction same as that of the line joining positive and negative centers. Polar and Non Polar molecules:

Thus molecules having dipole moment (μ =0) are called non polar molecules

and molecules with μ > 0 are polar . Greater is μ , greater is the polarity. For

poly-atomic molecules with two or more bonds, net dipole moment is the resultant of vector addition of individual moments . Eg. in the figure shown, the resultant dipole moment,

μ→ = μ1−→ + μ2−→

Units: While expressing dipole moments, generally charge is given in electrostatic units (esu ) and distance in angstrom units (1 Ao = 10−10 m). Thus, dipole moment of an electron separated from unit positive charge by a distance 1

Ao would be( 4.80 x10−10 esu ) × ( 10−8 cm ) = 4.8 x 10−18 esu cm = 4.8 Debye.

S.I units of dipole moment = (coulomb) (metre)

1D = 3.33564 × 10−30 C.m.

4

DIAGRAM

Applications of dipole moment

Cis and trans isomers can be distinguished by dipole moments, usually cis isomers have higher dipole

moment and hence, higher polarity.

5

DIAGRAM

Applications of dipole moment

Dipole moment is greatest for ortho isomer, zero for para isomer and less than that of ortho for meta isomer o > m > p

6

Applications of dipole moment

Determination of Ionic Character: Ionic character can be determined using dipole moment for e.g. experimental

dipole moment for HCl is 1.03 D and μ =d × e

as d= 1.26 and, e = 4.8 × 10−10 esu,

⇒μ = 6.05 Debye. Thus the ionic character

1.036.05≈16×100%

Thus we can say % ionic character

experimentaldipole momenttheoreticaldipole moment×100

Determination of Hybridization: Hybridization can be determined by dipole moment for eg.

i) If a molecule AB2 has μ = 0, the σ orbitals used by A ( Z < 21 ) must be sp

hybridised e.g BeF2.

ii) If a molecule AB3 has μ = 0, the σ orbitals used by A ( Z< 21 ) must be

Sp2 hybridised e.g. BF3 iii) If a molecule AB4 has μ = 0, the σ orbitals used by A ( Z < 21 ) must be

Sp3 hybridised e.g. CCl4

7

Bond angles

The internal angle between the lines joining the centre of the nucleus of one atom to the centres of the nuclei of two other atoms combined to it Measured with methods like X-ray diffraction, electron diffraction etc.

The bond angle is 180∘ BeF2, BeCl2, CO2, CS2, XeF2, ZnCl2, HgCl2, HCN, N2O, C2H2 , I−3 ion .

The bond angle is 120∘ BF3, BCl3, SO3, CH2O, C2H4 The bond angle is 109∘ 28′ in CH4, CCl4, SiH4, SiCl4, NH+4 The bond angle is 90∘ XeF4, BrF5, SF6 Bond angles are 90∘ and 180∘ ClF3, BrF3, ICl3 Bond angles are 90∘ and 120∘ PCl5, AsCl5, SbCl5 The bond angle is 107∘ NH3, H3O+ The bond angle in H2O - 104∘30′ NF3 - 102∘30′ OF2 - 102∘ Cl2O - 111∘

8

Covalent Radius

Half of the distance between 2 similar atom is joined by a covalent bond.

9

Van der Waals Radius

One half of the distance between the nuclei of adjacent atoms in the neighbouring molecules (in a solid). Van der Waals radius is larger than Covalent radius

As the size of the bonded atoms increases, their bond length also increases.

10

Bond Enthalpy

Amount of energy required to break one mole of bonds of a particular type between two atoms in gaseous state. Units: KJ/Mol. Relation with Bond Order, and Bond Length: Bond order is number of bonds between two atoms

Bond enthalpy∝1Bond length

Bond enthalpy∝Bond order

Bond length∝1Bond order

11

DIAGRAM

Resonance

Representation of molecule as more than one structure and none of them explains all the properties of the molecule singly is called resonance. Stability: Resonance stabilises the molecule. More the number of resonating structures, more will be the stability of the molecule. Bond length: Because of resonance, all the bond lengths are same. Resonance averages the bond characteristics as a whole.

Eg: All the bond lengths in benzene are identical.

Hybridisation

1

Hybridisation

1. The concept of hybridisation was introduced by Pauling. The intermixing of atomic orbitals of almost same energy and their

redistribution into an equal number of identical orbitals is known as hybridisation.

The orbitals of one and the same atom only involve in hybridisation. The orbitals formed in hybridisation process are called hybrid orbitals.

The no.of hybrid orbitals formed = no.of orbitals participating in hybridisation.

The hybrid orbitals symmetrically arranged around the nucleus such that they have maximum stability.

The order of repulsions will be in order of lp-lp > lp-bp > bp-bp repulsions. The orbitals involving in the hybridisation have different shapes but almost

same energy. The hybrid orbitals have same shape and same energy. The angle between any two hybrid orbitals in an atom is generally same. Hund's Rule: Electron filling in hybrid orbitals obey Hund's rule

Pauli's rule: The hybrid orbitals are involve only in 'sigma' bond formation.

2

sp - hybridisation

Diagonal or linear hybridization. One s- orbital combines with one p- orbital to give two identical orbitals called sp - hybrid orbitals.

The angle between the two sp-hybrid orbitals = 180∘ 50% s-character and 50% p-character. Most commonly found: Central atom in a linear molecule is generally sp.

3

sp2 -hybridisation

Trigonal hybridization. One s- orbital combines with two-p-orbitals to give three identical orbitals. The angle between any two sp2 -hybrid orbitals in an atom is 120∘. sp2 hybrid orbital will have 13rd s-character, 23rd p- character. The hybridization of central atom in a molecule having planar triagonal shape is sp2.

4

sp3 - hybridization

Tetrahedral hybridization (or) tetragonal hybridization. One s- orbital combines with three p-orbitals to give four identical orbitals.

The angle between any two sp3 -hybrid orbitals in an atom is 109∘.28′ sp3 - hybrid orbital will have 14th S- character, 34th P- character. Most commonly found

The hybridization of central atom in a molecule having tetrahedral shape is

sp3 .

5

sp3d -hybridization

4 One s-orbital, three p-orbitals and one d-orbital combines and gives five identical sp3d -hybrid orbitals.

The angle between two sp3d hybrid orbitals in an atom is 90o or 120o .

sp3d hybrid orbital will have 15 th s-character, 35 th p-character, 15 th d-character. The hybridization of central atom in a molecule having trigonal bi-pyramid

shape is sp3d.

6

sp3d2 - hybridisation

∙ One s-orbital , three p-orbitals and two d- orbitals combines to give six

identical orbitals called sp3d2 hybrid orbitals. ∙ The angle between any two sp3d2 hybrid orbital is 900.

∙ sp3d2 hybrid orbital will have shape of molecules / ions having only bond pair

of electrons.

7

Assigning the type of hybridization

I)Total molecular valency method: Step1: Total number of valence electron in the molecule (N) = ∑ individual valency of all constituent atoms.

Step2: If total electrons N > 8, No. of Hybridized orbitals = quotient of (N/8) Example1: PCl5 Total valence electrons = 5+(5 X 7)=40 40 / 8=5=no.of hybridized orbitals-->sp3d ∙ If remainder comes, that remainder should be divided by 2, still remainder

comes then sum of quotient and final remainder is equal to number of hybridized orbitals. Example 2:

XeOF2 Total valence electrons=8+6+(2 x 7)= 28 28 / 8 Quotient=3 Remainder=4 Remainder is divided by 2 Quotient=2 Hence, sum of Quotient = 3+2 = 5 = no.of hybridised orbitals --->sp3d

∙ Step3: If total valence electrons is either eight or less, dividing by 2,what

Quotient appears, should be considered as number of hybridized orbitals. Example: ,H2O Total valence electrons = 2 + 6 = 8 8 / 2 = 4 =sp3 II)Total number of Hybridized orbitals For neutral molecule = 1/2(V + M) V = no.of valence electrons M = no.of monovalent atom Total number of Hybridized orbitals (For cationic species) = 1/2(V + M - C); Total number of Hybridised orbitals (For anionic species) = 1/2(V + M + A); C = +Ve charge. A = -Ve charge.

8

Hybridisation in carbon compounds

The sp3 carbon atom involves in four single bonds

Molecular Orbital theory

1

Molecular orbital theory

Formation of Molecular Orbitals: Molecular orbitals are formed by LCAO method (linear combination of atomic orbitals) i.e. by addition or subtraction of wave functions of individual atoms.

ΨMO = ΨA+ΨB

Ψb = ΨA+ΨB

Ψa = ΨA−ΨB

Ψb = bonding molecular orbital

Ψa = Anti bonding molecular orbital

Order of Energies of Molecular Orbitals: Bonding orbitals < Non-bonding orbitals < Anti bonding orbitals. Bonding Molecular Orbital: Molecular orbitals of lower energy than atomic orbitals. Anti Bonding Molecular Orbital: Molecular orbitals of higher energy than atomic orbitals. Non Bonding Orbitals: Molecular orbitals which are not involved in bonding. Quantum Numbers, Rules and Shapes of Molecular Orbitals Molecular orbitals are characterized by a set of quantum numbers. Aufbau's principle, Pauli's exclusion principle and Hund's rule are applicable to molecular orbitals, during the filling of electrons.

The shape is governed by the shape of atomic orbitals.

2

Atomic and Molecular Orbitals

3

Atomic and Molecular Orbitals

4

Difference between σ and π MO’s

5

Stability of molecules

If Nb is the number of electrons occupying bonding orbitals and Na the number of electrons occupying anti-bonding orbitals, then (i) The molecule is stable if Nb is greater than Na (ii) The molecule is unstable if Nb is less than Na.

Note : KK σ1s2.σ∗1s2

6

Bond Order

Determines the relative stability of a molecule. : The number of covalent bonds in a molecule. Calculation: It is equal to one half of the difference between the number of electrons in the bonding and antibonding molecular orbitals.

Bond order=12[Number of bonding electrons - Number of antibonding

electrons] or = Nb−Na2 Bond Order 1 - Single bond Bond Order 2 - Double bond

Bond Order 3 - Triple bond But bond order may be fractional in some cases. The magnetic properties of molecules can also be ascertained.

7

Bond Order in Common Diatomic Molecules

Hydrogen molecule - Diamagnetic Total number of electrons = 2, filling in molecular orbitals we have

σ21s<σ∗01s

Bond order = Nb−Na2 = 2−02 = 1 Hence, there is a single bond between two hydrogen atoms and due to absence of unpaired electrons it is diamagnetic. Helium molecule (He2 ) The total number of electrons =4 and filling in molecular orbitals we have

σ21s<σ∗21s

Bond order = Nb−Na2 = 2−02 = 1

Hence, He2 molecule can not exist.

Nitrogen molecule (N2) The total number of electrons =14

σ21sσ∗21sσ22sσ∗22s {π22Pzπ22Py}

Bond order = (Nb−Nb)2 = 10−42 = 3 It is diamagnetic.

Oxygen O2

σ21sσ∗21sσ22sσ∗22s σ2Px{π22Pzπ22Py}{π∗12Pzπ∗12Py}

Total number of electrons =16 and electronic configuration is Bond order= 10−62 = 2 It is paramagnetic, due to unpaid electrons.

SuperOxide Ion (O−2) Total number of electrons (16 +1) = 17.

σ21sσ∗21sσ22sσ∗22s σ2Px {π22Pzπ22Py}{π∗22Pzπ∗12Py}

Bond order =(Nb−Na)2 = 10−72 = 1.5 It is paramagnetic.

Peroxide Ion O2−2

Total number ofelectrons (16 + 2) =18. The electronic configuration is σ21sσ∗21sσ22sσ∗22s σ2Px {π22Pzπ22Py}{π∗22Pzπ∗22Py}

Bond order = 10−82 = 1 It is diamagnetic.

[B2] no of electrons = 10

The electronic configuration is

σ21sσ∗21sσ22sσ∗22s{π12pyπ12pz}

it has 2 paired electrons . Hence, paramagnetic.

8

Oxygen molecule (O2)

Total number of electrons =16 and electronic configuration is Bond order= 10−52 = 2

It is paramagnetic.

9

Peroxide ion - (O2−2)

Total number ofelectrons (16 + 2) =18. The electronic configuration is Bond order = 10−82 = 1 It is diamagnetic. [B2] no of electrons = 10

The electronic configuration is

σ1s2σ∗1s2σ2s2σ∗2s2π2p1y = π2p1z

it has 2 paired electrons . Hence paramagnetic.