3.10.2 Bond Order (No. of Bonds)

This topic explains bond order — what it is, how it is calculated, and what it tells us about the stability of a molecule.

Definition

Bond Order: The total number of bonds formed when two atoms' atomic orbitals overlap with each other. It is half of the difference between the number of bonding electrons and antibonding electrons.

Main Idea

Bond order is a measure of the number of chemical bonds between two atoms. It is calculated using the number of electrons in bonding molecular orbitals (BMOs) and antibonding molecular orbitals (ABMOs). A positive bond order indicates a stable molecule, while a bond order of zero indicates an unstable molecule that does not exist.

Core Concept

Bond order = No. of electrons in BMOs - No. of electrons in ABMOs2

• Bond order = 0: Molecule is unstable and does not exist.
• Bond order > 0: Molecule exists and is stable.
• Higher bond order: Stronger bond, shorter bond length.

Detailed Explanation

What is Bond Order?

• Bond order is the total number of bonds formed when two atoms' atomic orbitals overlap with each other.
• It is half of the difference between the number of bonding electrons and antibonding electrons.
• It is a measure of the stability of a molecule.

Formula for Bond Order

Bond order = No. of electrons in BMOs - No. of electrons in ABMOs2

Example: Hydrogen Molecule (H₂)

• No. of electrons in the bonding orbitals = 2
• No. of electrons in the antibonding orbitals = 0

Bond order = 2 - 02 = 1

Conclusion: H₂ has a bond order of 1, meaning there is a single bond between the two hydrogen atoms. The molecule is stable.

Interpretation of Bond Order

Key Points

• If bond order is zero, the molecule is unstable, and it does not exist.
• A positive value of bond order reveals that the molecule exists and is stable.
• Higher bond order → stronger bond → shorter bond length → more stable molecule.

Examples of Bond Order

Explanation of Examples

• Total electrons = 4
• Configuration: σ 1s² σ^* 1s²
• Bonding electrons = 2
• Antibonding electrons = 2
• Bond order = 2 - 22 = 0
• Conclusion: He₂ is unstable and does not exist.

• Total electrons = 14
• Configuration: (σ 1s)² (σ^* 1s)² (σ 2s)² (σ^* 2s)² (π 2p_y)² = (π 2p_z)² (σ 2p_x)²
• Bonding electrons = 10
• Antibonding electrons = 4
• Bond order = 10 - 42 = 3
• Conclusion: N₂ has a triple bond and is very stable.

• Total electrons = 16
• Bonding electrons = 10
• Antibonding electrons = 6
• Bond order = 10 - 62 = 2
• Conclusion: O₂ has a double bond and is stable.

• Total electrons = 18
• Bonding electrons = 10
• Antibonding electrons = 8
• Bond order = 10 - 82 = 1
• Conclusion: F₂ has a single bond and is stable.

Relationship Between Bond Order and Bond Strength

• Higher bond order → stronger bond → shorter bond length → higher bond energy.
• Lower bond order → weaker bond → longer bond length → lower bond energy.

• N₂ (bond order 3) has a very strong triple bond.
• O₂ (bond order 2) has a double bond.
• F₂ (bond order 1) has a single bond.

Formula

Bond order = N_b - N_a2

• N_b = number of bonding electrons
• N_a = number of antibonding electrons

Interpretation

Examples

Relationship with Bond Strength

• Higher bond order → stronger bond → shorter bond length.
• Lower bond order → weaker bond → longer bond length.

Important Points

• 1. Bond order = N_b - N_a2.
• 2. Bond order = 0 → unstable; > 0 → stable.
• 3. Higher bond order → stronger bond → shorter bond length.
• 4. Examples: H₂ (1), He₂ (0), N₂ (3), O₂ (2), F₂ (1).
• 5. Paramagnetic: Unpaired electrons; Diamagnetic: No unpaired electrons.

Remember

• - Bond order: N_b - N_a2
• - Bond order = 0: Unstable
• - Bond order > 0: Stable
• - Higher bond order: Stronger bond, shorter length
• - Examples: H₂ (1), He₂ (0), N₂ (3), O₂ (2), F₂ (1)

Common Mistakes

• 1. Forgetting the formula: N_b - N_a2
• 2. Mixing up bonding and antibonding electrons: Bonding = lower energy; Antibonding = higher energy.
• 3. Thinking bond order can be negative: It cannot; it is always zero or positive.
• 4. Confusing bond order with bond length: Higher bond order → shorter bond length.
• 5. Ignoring magnetic properties: Unpaired electrons → paramagnetic.

Quick Revision

• Bond order = N_b - N_a2.
• Bond order = 0 → unstable; > 0 → stable.
• Higher bond order → stronger bond → shorter bond length.
• Examples: H₂ (1), He₂ (0), N₂ (3), O₂ (2), F₂ (1).
• Paramagnetic: Unpaired electrons; Diamagnetic: No unpaired electrons.

60-Second Revision

• Bond order: N_b - N_a2
• Bond order = 0: Unstable
• Bond order > 0: Stable
• Higher bond order: Stronger bond, shorter length
• Examples: H₂ (1), He₂ (0), N₂ (3), O₂ (2), F₂ (1)

QUALITY VALIDATION

• Topic: 3.10.2 Bond Order (No. of Bonds)
• Chapter: 3 (Chemical Bonding)

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• Important concepts covered: PASS
• Definitions checked: PASS
• Equations checked: PASS
• Derivations checked: N/A
• Examples checked: PASS
• Unsupported knowledge added: NO
• Unnecessary repetition: NO
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