Chapter 1:- Structure and Reactivity 1st year Book
Organic Chemistry
(Page 3)

Bond Parameters

  1. Bond Lengths :

    When in a molecule or ionic crystal, atoms are brought close to each other, the attraction takes place and therefore the potential energy of the system keeps on decreasing till a particular distance. If the atoms are further brought close to each other, the repulsion starts and the potential energy of the system begins to increase. When attractive and repulsive forces are equally balanced at a state of the stable equilibrium position, the potential energy of the system is minimum. At this state of equilibrium, the average distance between the centers of the nuclei of the two bonded atoms is called ‘bond length‘.

                                                                              Bond length is measured in either  or pm(1Aº=10−10m and 1pm=10−12) by x-ray diffraction methods or spectroscopic methods. In an ionic crystal, the bond length is the sum of their ionic radii and in a covalent compound, it is the sum of the covalent radii of the bonded atoms.

Factors Governing the Bond length:

Bond lengths depend on the following factors :

(i) Size of the atoms:

In molecules of the type ABn, the A−B bond length increases with an increase in the size of the bonded atom B. For example, the bond lengths in halogen acids are in the order: HF<HCl<HBr<HI.

(ii) Multiplicity of the bonds:

The bond length decreases with the multiplicity of the bonds. Thus, in carbon compounds, the C−C bond lengths changes in the order:  C−C>C=C>C≡C.

(iii) Type of hybridization:

An s- orbital is smaller in size, thus greater the s- character in hybrid orbitals, the shorter will be their size. Thus, the bond length decreases with the increase of s- character in hybrid orbitals involved in bonding.
For example, sp3C−H bond >sp2C−H bond >spC−H bond.

(iv) Type of compounds:

The normal length of the carbonyl group (>C=O) in ketone is about 1.20A°, while the value found in CO2 is 1.15 A∘. This indicates the value of the bond length of the same bond varies with different functional groups of the compounds. Bond Parameters

(v) Resonance effects:

In benzene, there are three C-C, single bonds, and three C−C, double bonds in the ring. It should have three bond lengths of equal lengths (1.54A°) and the remaining three bonds shorter but also equal in length (1.34 A°). However, spectroscopic measurements show that the C−C, bond lengths in benzene are not different, they are identical and equal in length (1.40 A°). This is due to the resonance effect. This effect results in partial C−C, double bonds in benzene. That is why C−C, bond lengths in benzene are neither equal to pure C−C, single bond nor pure C−C, double bond.
For the sake of simplicity, some bond lengths measured in organic compounds are given in table 1.03.
Table 1.03
Bond Length (A) Bond Length (A) Bond Length (A)
C−C 1.54 C−F 1.42 C−N 1.47
CC 1.40 G−Cl 1.77 C−O 1.43
C=C 1.21 C−Br 1.91 C=O 1.20
C−H 1.12 C−I 2.13 C−S 1.82
N−H 1.03 O−H 0.97


  1. Bond Energies:

They are expressed in kJmol−1 and are two types :

(i) Bond dissociation energy and
(ii) Bond energy

(i) Bond dissociation energy :

It is designated by D. This is the energy required to break a particular bond in a polyatomic molecule, in the gaseous phase into neutral separate atoms. Greater the bond dissociation energy, the stronger the bond. Since a particular type of bond present in different molecules(O-H bond in H2O, alcohols, phenols, and carboxylic acids) or even in the same molecule(eg. O-H bond in H2O and C−H bond in CH4 ) do not possess the same bond energies, bond energies are usually the average value.

(ii) Bond energy :

It is designated by E. In polyatomic molecules, e.g., in CH4, all the four C−H bonds are equivalent, but the energy required to break the first bond (CH4⟶C˙H3+H˙) is not the same as that for the second bond ( C˙H3⟶C˙H2+H˙ ) and so on for the fourth bond as it is clear from table 1.04.Bond energies, dissociation energy

The table shows that the values of the dissociation energies of the C−H bond vary in the first four compounds. Thus, the energy of a bond depends on the nature of the rest of the molecule. In practice, it is usual to take the average of all the different values, and this average value is called bond energy. For diatomic molecules, D and E are identical.

Factors Governing the Bond energy:

Various factors are responsible for the differences in the energy of a given bond in different compounds.

(i) Size of the atoms:

Greater the size of the atoms, the greater the bond length and the smaller the bond energy.

(ii) Multiplicity of the bonds:

The bond energy increases with an increasing multiplicity of the bonds. Thus, in carbon compounds, the C−C bond energy changes in the order: C−C<C=C<C≡C.

(iii) Angular strain:

The greater the number of lone pairs of electrons present on the bonded atoms, the greater will be the angular strain due to repulsion between the atoms, and hence, the smaller will be the bond dissociation energy.
Bond: C−C :N−N: :O:−:O: :F::−::F:
Lon Pair of electrons: 0 1 2 3
Bond Energy(kcal): 83.0 39.0 35.0 36.0


Another factor that is also responsible for influencing the value of bond energy is steric effects.

3. Bond angles:

The angle between the lines representing the directions of the bonds is called the ‘bond angle’. It is expressed in degrees (∘), minutes (′), and seconds(‘). e.g. HCH bond angle in CH4, H2C=CH2 and HC≡CH are 109∘28′,120∘
and 180∘ respectively.
Factors Governing the Bond angles:
Various factors are responsible for the differences in bond angles of different compounds having the same structures or having the same hybridization.
(i) Type of hybridization: Greater the s-character in hybrid orbitals involved in bonding, the greater the bond angles.
For example
Hybridization: sp3 sp2 sp
Bond Angle: 109.5°   < 120°   < 180°


(iii) Multiplicity of the bonds:

The bond angle increases with an increase in the multiplicity of the bonds of the molecules. Example, Molecule CH3−CH3  Molecule  Bond angle :CH3−CH3109∘28∘>CH2=CH2>CH≡CH120∘

(iv) Size of the central atoms:

In hydrides, the bond angle decreases with an increase in the size of the central atom of the molecules. Example,
Hydrides: H2O>H2S>H2Se>H2Te
Bond angle: 104.592.591.089.5

(v) Electronegativity of bonded atom:

In the covalent molecule of the type AXn, the X−A−X bond angle decreases with the increase in electronegativity of bonded atom X. If ‘ A ‘ is the central atom.
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