Lecture4 Characteristics of Covalent Bond (1)

7/31/2019 Lecture4 Characteristics of Covalent Bond (1)

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Uneven sharing of electrons

produces partial negative charges onthe atom having greater than half

share, leaving partial positive charge

on the atom having less than halfshare. There will be surplus of

negative charge around one nucleus

and deficiency around the other.


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Bond polarity depends on electronegativity

difference. The greater the electronegativity

difference, the greater the unevenness ofsharing, the greater share is acquired by the

initially more electronegative atom, thus

electron density will be displaced in the

direction of the more electron attracting atom.

The greater the difference between the electron

attracting power of the two atoms, the greater

is the ionic character of the covalent bonds.


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Electronegativity:F > O > Cl, N > Br > C, H


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Question: The electronegativity difference of

C - I bond is zero but the bond is polarizable .

Why?

In polar environment this covalent bond canacquire an appreciable degree of ionic character.

This is due to the polarizability of the outermostelectrons of iodine as a consequence of the bigsize of iodine atom. The distance between thenucleus and the outermost shell is great so that

the outermost electrons are not as stronglydrawn towards the nucleus as are the electronsof small atoms.


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Formal positive charge effectively increases

the electronegativity, thus increases polarity.

Calculation of Formal charge:

FC= group # - ( # of shared electrons) # of

unshared electrons


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Example: 1. N in ammonia

FC= 5- 3- 2 = 0

Example: 2. N in ammonium

FC = 5

4 = +1 more polar compared toammonia

N

H

HH

HH

H

N

H


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The formal positive charge effectively

increases the electronegativity of nitrogen.

There will be a tendency to pull the electron

closer to nitrogen resulting in increased partial

positive charges on each of the four hydrogen

atoms


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a. Determine which C O is more polar and explain.

or( FC = 0) (FC = +1)

The excess positive charge increases its electronattracting power so that there will be a greater

displacement of electron density towards oxygen.

H

H

H

C O H

H

H

H

C O H

H

+


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b. Which C H is more polar?

or

The C-H bond of chloroform is more polar than the C-H bond in methane because in chloroform adjacent

atoms are chlorine atoms which possess greaterelectron-attracting capacity than hydrogen atoms inmethane. The 3 chlorine will reduce the electron densityaround C, increasing the partial positive charge on Cwhich in turn will attract electron from hydrogen

increasing the partial positive charge on this atom.

Cl

Cl

Cl

C H

H

H

H

C H


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c. Which C H is more polar?

or

The C-H bond in acetylene is more polar than C-H bondin ethylene because C in acetylene is an sp ( s and pcharacter) while C in ethylene is an sp2 ( 1/3 s and 2/3 p).

The s character of hybrid orbital is a measure of theelectron-attracting capacity of the atom. The electron-attracting capacity of C in acetylene is greater than inthat of ethylene. The electron attracting power as a resultof s character of the hybrid orbital, responsible for the

polarity of C-H bond in acetylene is also known asORBITAL ELECTRONEGATIVITY.

CH CH CH2

CH2


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d. Which C H is more polar?

or

The C-H bond in ethylene is more polar than the C-H bondin ethane because oforbital electronegativity.

Importance of bond Polarity:1. It contributes to the characteristic of bond that is itinfluenced bond length and energy.2. It influences molecular polarity, thus in the end

determines the physical properties like melting point, boilingpoint, solubility.

3. It influences chemical reactivity.

C C

H

H

H

H

CH

H

H

C H

H

H
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2. BOND LENGTH
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The bond length is the distance

between those two atoms. Thegreater the number of electrons

between two atoms, the closer the

atoms can be brought towards oneanother, and the shorter the bond.This is measured in Angstrom units

which 1A is equal to 10-8 cm. It hasbeen shown that bond length is

correlated with bond polarity,

hybridization of orbital and


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When two atoms approach each

other, their interaction isinfluenced by:

a. repulsion between twoelectron, clouds

b. repulsion between the two

nuclei

c. attraction between the nucleus

of each, and


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Factors influencing bond length

a. Bond Polarity

Increasing polarity decreasing bond lengthType of Single Bonds Bond Length

C-C 1.538

C-N 1.471

CO 1.430

CF 1.380


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b. Hybridization

Increasing s ( near the nucleus) character of hybrid orbital

decreasing bond length or orbital electronegativity leads to

reduction of bond distance.

Ethane CH3-CH3 sp3 sp3 1.538

Propylene CH3-CH=CH2 sp3 sp2 1.501

Methylacetylene sp3-sp 1.459Ethylene CH2=CH2 sp

2-sp2; p-p 1.339

Acetylene sp-sp; p-p; p-p 1.207

C CHCH3

CH CH


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c.atomic size

small atoms will form shorter bond.

Example: H-H < CH4-CH4


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d. bond order

decreasing bond order increasing bond

length.

Triple bond< double bond< single bond

sp sp2 sp3


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e. and electron delocalization

illustrated by system containing atoms in the trigonal stateof hybridization.

Example:

a. CH2 = CH - Cl

b. O = C

c. CH3 CH = CH2
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3. BOND ENERGY orBOND STRENGTH
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energy required to break the bond and at the

same time the energy release when the bondis formed.

This is express in terms of kilo calories permole

Bond energy is variable depending on length,the shorter the bond the stronger the bond.


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Bond strengthening are attributed to:

a. orbital hybridization

sp3 sp3sp3- sp2sp3 spsp2 sp2 Bond energy increases bond length decreases

sp2 spsp - sp


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b. bond polarity - when bond polarity increases

bond energy also increases

- C C sp3sp3 0

= C

C = sp2

sp2

0- C = C - sp2 sp2 1 increasing pi bond increasing= C C= sp sp 0 bond energy- C = C- sp-sp 2


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c. pi bonding

- C C sp3sp3 0

= C C = sp2 sp2 0

- C = C - sp2 sp2 1 increasing pi bond

= C C= sp sp 0 increasing bond

-C = C- sp-sp 2 energy


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d.Reduction of bond strength is observed

when there is loss of overlap of hybrid

orbitals. The loss of overlap is a consequence

of forced bending of orbitals. A vivid exampleof loss of overlap as a result of forced bending

of sp3 orbitals is seen in cyclopropane. Forced

bending results in angular strain due tocompression of the tetrahedral bond angle.
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4. BOND order
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The bond order is equal to the number of

bonds between two atoms.


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The BO is an indication of the bond length,the greater the bond order, the shorter thebond and the greater the bond strength.


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5. BOND ANGLE
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a. governed by hybridization of central atom

sp3 - 109.5o

sp2 120 osp - 180 o


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b. influenced by the presence of lone pairs

CH4 sp3 109o28

NH3

sp3 106o47

H2O sp3 104o31

As the number of lone pair increases, the bond angle

decreases.

Lecture4 Characteristics of Covalent Bond (1)

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