Bond Parameters - Chemistry Class 11 - NEET PDF Download

Table of Contents
1. What are Bond Parameters?
2. Bond Length
3. Bond Enthalpy (Bond Energy)
4. Bond Angle
5. Bond Order
6. Formal Charge
7. Dipole Moment
8. Resonance in Chemical Bonding
View more

What are Bond Parameters?

Covalent bonds are described by a set of measurable properties called bond parameters. The main bond parameters are bond length, bond angle, bond order and bond enthalpy (bond energy). These parameters give information about the geometry, strength and stability of molecules and help predict chemical behaviour.

Bond Length

Bond Length

• Definition: Bond length (or bond distance) is the equilibrium distance between the centres of the nuclei of two bonded atoms in a molecule.
• Measurement: Bond lengths are measured by spectroscopic methods, X-ray diffraction and electron diffraction techniques.
• Covalent radius: In a covalent bond each atom contributes a part of the internuclear distance; the contribution from an atom is called its covalent radius.
• Periodic trend: Covalent radii increase down a group and generally decrease across a period for s- and p-block elements.
• Multiplicity: For the same pair of atoms, bond length decreases as bond multiplicity increases (single > double > triple).

Factors affecting the bond length

• Size of the atoms: Larger atoms give longer bonds. Example: bond lengths in H-X follow HI > HBr > HCl > HF.
• Multiplicity of bond: More shared electron pairs bring nuclei closer; hence C≡C < C=C < C-C.
• Type of hybridisation: Greater s-character gives shorter bonds because s-orbitals are smaller. Example:
  Bond lengths: sp³ C-H > sp² C-H > sp C-H
  s-character: (25%)      (33%)      (50%)
• Resonance and delocalisation: Delocalisation produces bonds of intermediate length. Example: C-C in ethane = 1.54 Å, C=C in ethene = 1.34 Å, but in benzene each C-C = 1.39 Å due to resonance.

Bond Enthalpy (Bond Energy)

• Definition: Bond enthalpy is the enthalpy required to break one mole of a particular bond in the gaseous state; it is also called bond dissociation enthalpy.
• Interpretation: Larger bond enthalpy means a stronger bond and greater stability of that bond in the molecule.

Factors affecting bond energy

• Size of the atoms: Larger atoms form longer bonds with smaller overlap of orbitals; hence bond strength (bond dissociation enthalpy) decreases as atomic size increases.
• Multiplicity of bonds: For the same atoms, higher multiplicity (double, triple) gives greater bond dissociation enthalpy because atoms are closer and there are more electron pairs to break. Example: Bond dissociation enthalpies: H-H < O=O < N≡N.
• Number of lone pairs: Lone pairs on bonded atoms increase electron pair repulsion and can weaken a bond, lowering the bond dissociation enthalpy.

Bond Enthalpies

Bond Angle

Definition: The bond angle is the angle between the directions of two bonds that originate from the same atom. In orbital terms it is the angle between the hybrid orbitals (or atomic orbitals) containing the bonding electron pairs.

Factors affecting bond angle

• Hybridisation: Bond angles depend on the hybridisation state of the central atom.
  Examples: CH₄ (sp³), BCl₃ (sp²), BeCl₂ (sp)
  As s-character increases, the hybrid orbital is smaller and bond angle tends to increase.
• Lone pair repulsion: Lone pairs repel bonding pairs more strongly than bonding pairs repel each other, reducing bond angles. Example: NH3 has bond angle ≈ 107°, H₂O ≈ 105° though both central atoms are sp³ hybridised.
• Electronegativity of surrounding atoms: If the central atom is fixed, bond angle increases when the substituent atoms are less electronegative because bonding electron density is held more evenly; conversely, greater electronegativity of substituents can pull bonding electrons and slightly change bond angles.

Example of Bond Angles

Bond Order

• Definition: In Lewis representation, bond order is the number of chemical bonds between a pair of atoms (single = 1, double = 2, triple = 3).
• Fractional bond orders: For molecules with odd electrons or delocalisation, bond order can be fractional. Example: In nitric oxide (NO) the partial or three-electron bond is equivalent to a bond order of 1.5 in some descriptions.
• Correlation: As bond order increases, bond enthalpy increases and bond length decreases.

Lewis's structure of NO

Formal Charge

Definition: The formal charge on an atom in a Lewis structure is the difference between the number of valence electrons in the free atom and the number of electrons assigned to that atom in the structure, where each bonding pair is split equally between the two bonded atoms and lone pairs are assigned wholly to the atom on which they reside.

Formal Charge Formula

Example:

Calculate the formal charge on each O-atom of the O₃ molecule.

Sol. Lewis structure of O₃ is:

The atoms have been numbered as 1, 2 and 3.

Formal charge on end O-atom numbered 1

The formal charge on the central O-atom numbered 2

Formal charge on end O-atom numbered 3

Hence, the resonance representation showing formal charges is:

Example:

Write the formal charges on atoms in (i) carbonate ion (ii) nitrite ion.

Sol. (i) Lewis structure of CO₃^2- ion is

Formal charge on C atom

Formal charge on double-bonded O atom

Formal charge on single-bonded O atom

(ii) Lewis structure of NO₂^- ion is

Formal charge on N atom

Formal charge on double-bonded O atom

Formal charge on single-bonded O atom

Significance of formal charge: Formal charge helps to choose the most stable Lewis structure among several possibilities. The most stable structure is usually the one in which the magnitudes of formal charges are minimised and negative formal charges reside on the more electronegative atoms.

Dipole Moment

• If a covalent bond is formed between two dissimilar atoms A and B, the shared electron pair is displaced towards the more electronegative atom. This creates a partial negative charge on the more electronegative atom and a partial positive charge on the other - producing a dipole.
• Definition: Dipole moment (μ) is defined as the product of the magnitude of the partial charge (q) and the distance (r) between the centres of positive and negative charges: μ = q × r.
• Dipole moment is a vector quantity; its direction is from the positive centre to the negative centre.
• Molecules with net μ = 0 are non-polar; those with μ > 0 are polar. Greater μ indicates greater polarity.
• For polyatomic molecules, the net dipole moment is the vector sum of individual bond dipole moments.
• Units: In electrostatic units (esu) and centimetres the dipole moment of an electron separated from a unit positive charge by 1 Å (1 Å = 10^-10 m = 10^-8 cm) is (4.80 × 10^-10 esu) × (10-8 cm) = 4.8 × 10^-18 esu·cm = 4.8 Debye (D).

Applications of dipole moment

1. Dipole moment helps predict the molecular geometry when combined with other structural information.
2. It determines molecular polarity and influences physical properties like boiling point and solubility.
3. Dipole moment distinguishes symmetrical and non-symmetrical molecules: for example, CO₂ is symmetrical and has μ = 0 while H₂O is polar with μ ≈ 1.85 D.
4. Cis and trans isomers often have different dipole moments; cis isomers generally have higher μ and hence are more polar than trans isomers.
5. In substituted benzene derivatives, dipole moments of ortho, meta and para isomers often follow o > m > p.
6. Dipole moment can be used to estimate percentage ionic character of a bond by comparing experimental μ with the value expected for complete charge separation.
7. Dipole moment can help infer the hybridisation of the central atom: if an AB_n molecule has μ = 0 and the central atom A has low atomic number (z < 21), it indicates ideal symmetric hybridisation (examples below).

 

• If a molecule AB₂ has μ = 0, the central atom A (z < 21) is likely sp hybridised; example: BeF₂.
• If AB₃ has μ = 0, the central atom A (z < 21) is likely sp2 hybridised; example: BF₃.
• If AB₄ has μ = 0, the central atom A (z < 21) is likely sp3 hybridised; example: CCl₄.

Resonance in Chemical Bonding

Some molecules or ions cannot be represented by a single Lewis structure; instead several canonical (resonance) structures can be drawn. The actual structure is a hybrid (resonance hybrid) of these canonical forms and shows delocalisation of electrons over more than two atoms.

Example: Ozone (O₃) can be drawn as two canonical forms in which a double bond and a single bond appear alternately. Experiment shows both O-O bonds are identical and intermediate between single and double. The bonding electron pair is delocalised over the three atoms; neither canonical formula alone represents the true structure.

Structures (A) and (B) are the canonical structures and the actual molecule is the resonance hybrid (C). Delocalisation lowers the energy of the system; the difference in energy between the resonance hybrid and the most stable canonical form is called the resonance stabilization energy.

Some other examples of resonance:

Carbonate ion, CO₃^2- (three equivalent canonical forms):

All C-O bond lengths in carboxylate ions are equal because of resonance:

Benzene - the six C-C bonds are equivalent due to delocalisation of π-electrons:

Vinyl chloride - resonance and conjugation affect bond character and reactivity:

Resonance summary: Resonance provides a more accurate description of molecules with delocalised electrons, explains equalisation of bond lengths, influences stability and reactivity, and is quantified qualitatively by resonance stabilization energy.