Chemical Bonding – Bond Order & Magnetic Nature | Class 11 + JEE Main 2027

Bond Order and Magnetic Nature are important concepts in Class 11 Chemistry and JEE Main. These topics are studied using Molecular Orbital Theory (MOT), which explains how electrons occupy molecular orbitals and how this affects the stability, strength and magnetic behaviour of a molecule.

In this article, we will understand bonding and antibonding molecular orbitals, learn the bond order formula, distinguish between paramagnetic and diamagnetic species, and compare O₂, O₂⁺, O₂⁻ and O₂²⁻. We will also solve a JEE Main-style question and revise the most useful exam tricks.

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1. Molecular Orbital Theory (MOT)

According to Molecular Orbital Theory, atomic orbitals combine to form molecular orbitals. The electrons of a molecule occupy these molecular orbitals rather than belonging exclusively to individual atoms.

Molecular orbitals extend over the molecule and describe the distribution of electrons. Their energy and electron occupancy help us predict whether a molecule is stable, how strong its bond is, and whether it is attracted to a magnetic field.

Types of Molecular Orbitals

1. Bonding Molecular Orbitals

Bonding molecular orbitals generally have lower energy than the atomic orbitals from which they form. Electrons occupying bonding orbitals help stabilise the molecule.

Examples include:

• σ1s
• σ2s
• π2p

2. Antibonding Molecular Orbitals

Antibonding molecular orbitals generally have higher energy and weaken the bond when occupied. They are represented by an asterisk (*).

Examples include:

• σ*1s
• σ*2s
• π*2p

Remember: Bonding electrons increase the net bond order, while antibonding electrons decrease it.

Rules for Filling Molecular Orbitals

Electrons are filled into molecular orbitals according to these three principles:

• Aufbau principle: Electrons occupy available orbitals in increasing order of energy.
• Pauli exclusion principle: An orbital can contain a maximum of two electrons with opposite spins.
• Hund's rule: Degenerate orbitals are occupied singly with parallel spins before pairing occurs.

Hund's rule is especially important for understanding the magnetic nature of oxygen and its ions. If electrons remain unpaired in molecular orbitals, the species is paramagnetic.

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2. Bond Order Formula

Bond order is a measure of the net bonding between two atoms in Molecular Orbital Theory. It helps us compare the strength and length of bonds between comparable species.

Bond Order Formula

Bond Order = (Nᵦ − Nₐ)/2

Where:

• Nᵦ = Total number of electrons in bonding molecular orbitals.
• Nₐ = Total number of electrons in antibonding molecular orbitals.

The formula subtracts the antibonding electrons from the bonding electrons and divides the difference by two.

Example: Bond Order of H₂

The H₂ molecule has two electrons. Both electrons occupy the bonding σ1s molecular orbital, while the antibonding σ*1s orbital remains empty.

• Bonding electrons: Nᵦ = 2
• Antibonding electrons: Nₐ = 0

Using the formula:

Bond Order = (Nᵦ − Nₐ)/2

Bond Order = (2 − 0)/2 = 1

Answer: Bond order of H₂ = 1

A bond order of 1 corresponds to a single bond in this example.

What Does Bond Order Tell Us?

• Bond order greater than zero: Generally indicates net bonding in the basic MOT model.
• Bond order equal to zero: Indicates no net bond in the basic MOT model.
• Higher bond order: Generally indicates a stronger and shorter bond when comparable species are considered.

JEE Main Trick: Never forget to subtract antibonding electrons. Adding an electron to an antibonding orbital usually decreases the bond order by 0.5.

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3. Magnetic Nature: Paramagnetic vs Diamagnetic

The magnetic nature of a molecule depends on the arrangement of its electrons in molecular orbitals. To determine whether a species is paramagnetic or diamagnetic, check whether it contains any unpaired electrons.

Paramagnetic Species

A species is paramagnetic if it has at least one unpaired electron. Such a species is attracted to an external magnetic field.

Examples include:

• O₂
• O₂⁺
• O₂⁻

Diamagnetic Species

A species is diamagnetic if all its electrons are paired. Diamagnetic substances are weakly repelled by an external magnetic field.

Examples include:

• N₂
• O₂²⁻

The Golden Rule

At least one unpaired electron → Paramagnetic

All electrons paired → Diamagnetic

Do not determine magnetic nature by looking only at the total number of electrons. You must examine how the electrons are distributed among the molecular orbitals.

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4. Bond Order and Magnetic Nature of O₂

Oxygen is one of the most important examples in Molecular Orbital Theory. According to the molecular orbital configuration, O₂ contains two unpaired electrons, which explains why it is paramagnetic.

Step 1: Count the Electrons

Each oxygen atom has eight electrons. Therefore, a neutral O₂ molecule has:

Total electrons = 8 + 8 = 16

Step 2: Write the Valence Molecular Orbital Configuration

For O₂, the valence molecular orbital configuration is:

(σ2s)² (σ*2s)² (σ2pᶻ)² (π2pˣ)² (π2pʸ)² (π*2pˣ)¹ (π*2pʸ)¹

The two degenerate antibonding π* orbitals each contain one electron. This arrangement follows Hund's rule.

Step 3: Count Bonding Electrons

The bonding orbitals contain:

• (σ2s)² = 2 electrons
• (σ2pᶻ)² = 2 electrons
• (π2pˣ)² = 2 electrons
• (π2pʸ)² = 2 electrons

Total bonding electrons, Nᵦ = 8

Step 4: Count Antibonding Electrons

The antibonding orbitals contain:

• (σ*2s)² = 2 electrons
• (π*2pˣ)¹ = 1 electron
• (π*2pʸ)¹ = 1 electron

Total antibonding electrons, Nₐ = 4

Step 5: Calculate Bond Order

Bond Order = (Nᵦ − Nₐ)/2

Bond Order = (8 − 4)/2

Bond Order = 2

Step 6: Determine Magnetic Nature

O₂ has one unpaired electron in each of its two antibonding π* orbitals. Therefore, it has two unpaired electrons.

Final Answer: O₂ has bond order 2 and is paramagnetic.

JEE Main Trick: O₂ is paramagnetic because it contains two unpaired electrons in its antibonding π* orbitals. This is a classic application of Molecular Orbital Theory.

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5. Bond Order Comparison: O₂⁺, O₂, O₂⁻ and O₂²⁻

Comparing oxygen and its ions is a frequently useful exercise for JEE Main. The key is to understand which orbital gains or loses electrons when the charge changes.

Comparison Table

Species	Bond Order	Unpaired Electrons	Magnetic Nature
O₂⁺	2.5	1	Paramagnetic
O₂	2	2	Paramagnetic
O₂⁻	1.5	1	Paramagnetic
O₂²⁻	1	0	Diamagnetic

Why Does O₂⁺ Have Bond Order 2.5?

When one electron is removed from O₂, it is removed from an antibonding π* orbital. The number of antibonding electrons decreases by one.

Bond order = (8 − 3)/2 = 2.5

One unpaired electron remains, so O₂⁺ is paramagnetic.

Why Does O₂⁻ Have Bond Order 1.5?

When one electron is added to O₂, it enters an antibonding π* orbital. The number of antibonding electrons increases by one.

Bond order = (8 − 5)/2 = 1.5

One unpaired electron remains, so O₂⁻ is paramagnetic.

Why Does O₂²⁻ Have Bond Order 1?

Two electrons are added to O₂ to form O₂²⁻. Both π* orbitals become fully occupied, and all electrons are paired.

Bond order = (8 − 6)/2 = 1

Since all electrons are paired, O₂²⁻ is diamagnetic.

Shortcut for Oxygen Species

• Remove one electron from O₂: bond order increases by 0.5.
• Add one electron to O₂: bond order decreases by 0.5.
• Add two electrons to O₂: bond order decreases by 1 overall.

Bond order sequence:

O₂⁺ > O₂ > O₂⁻ > O₂²⁻

For these comparable oxygen species, higher bond order generally corresponds to a stronger and shorter bond.

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6. JEE Main Solved Question: O₂⁻

Question

Determine the bond order and magnetic nature of the superoxide ion, O₂⁻.

Step 1: Start with Neutral O₂

For neutral O₂:

• Bonding electrons = 8
• Antibonding electrons = 4
• Bond order = (8 − 4)/2 = 2

Step 2: Add One Electron

O₂⁻ contains one extra electron compared with neutral O₂. This electron enters an antibonding π* molecular orbital.

Therefore:

• Bonding electrons = 8
• Antibonding electrons = 5

Step 3: Calculate Bond Order

Bond Order = (Nᵦ − Nₐ)/2

Bond Order = (8 − 5)/2

Bond Order = 3/2 = 1.5

Step 4: Determine Magnetic Nature

After adding one electron, one of the π* orbitals contains a pair of electrons, while the other contains one unpaired electron.

Since at least one unpaired electron is present, O₂⁻ is paramagnetic.

Final Answer: Bond order = 1.5; magnetic nature = paramagnetic.

Exam Shortcut: For O₂⁻, adding one electron to an antibonding orbital reduces the bond order of O₂ from 2 to 1.5.

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7. Common Mistakes and JEE Main Exam Tricks

Mistake 1: Adding Antibonding Electrons Instead of Subtracting Them

The correct formula is:

Bond Order = (Nᵦ − Nₐ)/2

Always subtract the antibonding electron count from the bonding electron count.

Mistake 2: Forgetting the Effect of Electron Addition

If an electron enters an antibonding orbital, the bond order decreases by 0.5. If an electron is removed from an antibonding orbital, the bond order increases by 0.5.

Mistake 3: Calling O₂ Diamagnetic

O₂ is paramagnetic because its molecular orbital configuration contains two unpaired electrons in the π* antibonding orbitals.

Mistake 4: Confusing Bond Order with Magnetic Nature

Bond order and magnetic nature are related to electron configuration, but they answer different questions. Bond order measures net bonding, while magnetic nature depends on whether electrons are unpaired.

Mistake 5: Forgetting Hund's Rule

The two degenerate π* orbitals in O₂ are occupied singly before pairing. This leaves two unpaired electrons and explains the paramagnetic nature of O₂.

Mistake 6: Assuming Every Species with a High Bond Order Is Diamagnetic

A species can have a relatively high bond order and still be paramagnetic. For example, O₂⁺ has bond order 2.5 and is paramagnetic.

Important Exam Tricks

• Bond order = (bonding electrons − antibonding electrons)/2.
• Bonding orbitals are generally lower in energy than their corresponding antibonding orbitals.
• Antibonding orbitals are marked with an asterisk (*).
• One unpaired electron is sufficient for paramagnetism.
• All electrons paired means the species is diamagnetic.
• O₂ has two unpaired electrons and bond order 2.
• O₂²⁻ has no unpaired electrons and bond order 1.
• For comparable species, higher bond order generally means a shorter, stronger bond.

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8. Quick Revision Formula Sheet

Concept	Important Result
Bond order formula	(Nᵦ − Nₐ)/2
Bonding molecular orbital	Generally lower energy; stabilises the molecule
Antibonding molecular orbital	Generally higher energy; weakens the bond when occupied
Paramagnetic	At least one unpaired electron
Diamagnetic	All electrons paired
O₂⁺	Bond order 2.5; paramagnetic
O₂	Bond order 2; paramagnetic
O₂⁻	Bond order 1.5; paramagnetic
O₂²⁻	Bond order 1; diamagnetic
Effect of removing one antibonding electron	Bond order increases by 0.5
Effect of adding one antibonding electron	Bond order decreases by 0.5

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9. One-Minute Revision

• MOT: Atomic orbitals combine to form molecular orbitals.
• Bonding MO: Lower energy and stabilising.
• Antibonding MO: Higher energy and bond weakening when occupied.
• Bond order: (Nᵦ − Nₐ)/2.
• Paramagnetic: At least one unpaired electron.
• Diamagnetic: All electrons paired.
• O₂: Bond order 2; two unpaired electrons; paramagnetic.
• O₂⁺: Bond order 2.5; paramagnetic.
• O₂⁻: Bond order 1.5; paramagnetic.
• O₂²⁻: Bond order 1; diamagnetic.
• Bond strength: Higher bond order generally means stronger bonding for comparable species.

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10. Practice Questions for JEE Main

Question 1

What is the bond order of H₂ according to Molecular Orbital Theory?

1. 0
2. 0.5
3. 1
4. 2

Answer: (c) 1

Calculation: Bond order = (2 − 0)/2 = 1.

Question 2

Which of the following species is diamagnetic?

1. O₂
2. O₂⁺
3. O₂⁻
4. O₂²⁻

Answer: (d) O₂²⁻

Reason: All electrons are paired in O₂²⁻.

Question 3

What is the bond order of O₂⁺?

1. 1
2. 1.5
3. 2
4. 2.5

Answer: (d) 2.5

Reason: Removing one electron from an antibonding orbital increases the bond order by 0.5.

Question 4

Why is neutral O₂ paramagnetic?

1. It has no electrons.
2. It has two unpaired electrons in π* orbitals.
3. All its electrons are paired.
4. Its bond order is zero.

Answer: (b) It has two unpaired electrons in π* orbitals.

Question 5

What is the bond order of O₂⁻?

1. 0.5
2. 1
3. 1.5
4. 2.5

Answer: (c) 1.5

Calculation: Bond order = (8 − 5)/2 = 1.5.

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11. Chemical Bonding Revision PDF

Embed your Bond Order and Magnetic Nature notes or revision PDF below. Replace the placeholder with the direct URL of your uploaded PDF.

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12. Frequently Asked Questions (FAQs)

Q1. What is bond order in Molecular Orbital Theory?

Bond order is calculated as half the difference between the number of bonding and antibonding electrons: (Nᵦ − Nₐ)/2.

Q2. What does a positive bond order indicate?

A positive bond order generally indicates net bonding in the basic Molecular Orbital Theory model.

Q3. Why is O₂ paramagnetic?

O₂ has two unpaired electrons in its antibonding π* molecular orbitals, making it paramagnetic.

Q4. What is the bond order of O₂?

The bond order of O₂ is 2.

Q5. What is the bond order of O₂⁺?

The bond order of O₂⁺ is 2.5. It is paramagnetic because one unpaired electron remains.

Q6. What is the bond order of O₂⁻?

The bond order of O₂⁻ is 1.5, and it is paramagnetic.

Q7. Why is O₂²⁻ diamagnetic?

Its electrons are paired in the molecular orbitals, so it has no unpaired electrons and is diamagnetic.

Q8. How does adding an electron to an antibonding orbital affect bond order?

Adding one electron to an antibonding orbital decreases bond order by 0.5.

Q9. Which oxygen species has the highest bond order among O₂⁺, O₂, O₂⁻ and O₂²⁻?

O₂⁺ has the highest bond order, equal to 2.5.

Q10. Which rule helps explain the unpaired electrons in O₂?

Hund's rule explains why electrons occupy the two degenerate π* orbitals singly before pairing.

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Final Thoughts

Bond Order and Magnetic Nature become easier once you understand molecular orbital filling. Focus on the bond order formula, the difference between bonding and antibonding orbitals, and the presence or absence of unpaired electrons.

For JEE Main 2027, practise comparing O₂⁺, O₂, O₂⁻ and O₂²⁻ until you can quickly determine both their bond order and magnetic nature.

Final JEE Main Tip: First determine the molecular orbital electron configuration. Then calculate bond order and count the unpaired electrons. This two-step method helps you answer both types of questions accurately.