Molecular Orbital Diagrams and Electronic Configuration with Examples
Molecular Orbital Theory (MOT) is a key quantum mechanical model that describes chemical bonding by combining atomic orbitals to form new molecular orbitals spread over an entire molecule. Unlike older approaches, such as valence bond theory, MOT explains molecular structure, stability, and electronic properties with much higher accuracy—vital for JEE Main Chemistry success and clear understanding of chemical bonding concepts.
What is Molecular Orbital Theory?
In Molecular Orbital Theory, electrons occupy molecular orbitals that extend across a whole molecule rather than being localised between individual atoms. When two atoms combine, their atomic orbitals (AOs) overlap and create two new molecular orbitals—one lower in energy (bonding) and one higher (antibonding). This fundamental distinction enables MOT to accurately predict features like the magnetism of O2 and even bond order discrepancies for molecules such as N2 and B2.
Key Assumptions and Main Points of MOT
- Atomic orbitals of similar energy and proper symmetry overlap to form molecular orbitals.
- Molecular orbitals extend over the entire molecule and are not assigned to any one atom.
- There are always equal numbers of bonding and antibonding molecular orbitals formed.
- Electrons fill molecular orbitals starting from the lowest energy (Aufbau principle).
- Bond order is calculated as half the difference between bonding and antibonding electrons.
- Follows Hund’s rule and Pauli’s exclusion principle for MO electron filling.
Types and Formation of Molecular Orbitals
When two atomic orbitals (like 1s or 2p) from adjacent atoms overlap constructively (in-phase), they form a bonding molecular orbital (lower energy, symbolised as σ or π). When the overlap is destructive (out-of-phase), an antibonding molecular orbital (higher energy, noted as σ* or π*) forms. The type—σ (sigma) or π (pi)—depends on the axis and symmetry of overlap. For example, end-to-end overlap forms σ/σ*, while sidewise overlap forms π/π* orbitals.
Comparison: Atomic Orbital Theory vs. Molecular Orbital Theory
| Valence Bond Theory | Molecular Orbital Theory |
|---|---|
| Electrons localised between atoms, bonds formed by overlapping AOs. | Electrons occupy orbitals spread over the whole molecule. |
| Cannot explain O2 paramagnetism or bond order exceptions. | Explains paramagnetism of O2 and other anomalies. |
| Limited prediction of magnetic and spectral properties. | Accurately predicts magnetic, spectroscopic, and stability features. |
This difference is especially important for JEE questions comparing molecule-based predictions—like why O2 is paramagnetic while N2 is diamagnetic—where only molecular orbital theory provides a satisfactory answer.
How to Draw Molecular Orbital Diagrams: H2, N2, O2
You can visualise molecular orbital diagrams by arranging energy levels: atomic orbitals combine to form σ and π bonding orbitals and their corresponding antibonding σ* and π* orbitals. The energy ordering changes for lighter molecules (up to N2) and heavier ones (from O2 onwards):
- For H2: Only 1s orbitals combine—forming σ1s (bonding) and σ1s* (antibonding).
- For N2: The ordering is σ1s, σ1s*, σ2s, σ2s*, π2p, σ2p, π2p*, σ2p*.
- For O2: The σ2p orbital drops below π2p in energy.
Drawing these diagrams stepwise helps in electron filling, bond order calculation, and predicting paramagnetism—frequent JEE main questions.
How to Write Molecular Orbital Electronic Configuration and Calculate Bond Order
To find the molecular orbital electronic configuration for species like O2 or N2:
- Count the total number of electrons in the molecule or ion.
- Fill the MOs according to the correct energy order, applying Hund’s rule and Pauli’s exclusion principle.
- Solve for bond order using the formula: Bond Order (B.O.) = (nbonding − nantibonding) / 2.
For example, for O2 (16 electrons): Electronic configuration = (σ1s)2 (σ1s*)2 (σ2s)2 (σ2s*)2 (σ2pz)2 (π2px)2 (π2py)2 (π2px*)1 (π2py*)1. Bond order = (10 − 6) / 2 = 2.
Applications: Predicting Bond Order, Stability, and Magnetism
- MOT explains paramagnetism in O2, as it has two unpaired electrons in its π2p* orbitals.
- Helps differentiate the bond order and strength of N2 (triple bond, B.O. = 3) and O2 (double bond, B.O. = 2).
- Predicts the possibility of molecule existence: e.g., He2 (B.O. = 0) does not exist under normal conditions.
- Explains magnetic properties (dia-/para-magnetism) and supports JEE Main reasoning questions.
- Enables correct ranking of species by bond order, bond length, and reactivity for quick JEE comparisons.
JEE Key Points: Common Pitfalls, Tricks, and Revision Table
| Quick Point | JEE Relevance |
|---|---|
| Sigma (σ) bonding MOs are lower in energy than pi (π) MOs, except for B2, C2, N2. | Correctly draw diagrams for elements Z ≤ 7 vs. Z ≥ 8. |
| Bond order zero means the molecule is not stable. | Account for ions—O22- and O2+ have modified bond orders. |
| Unpaired electrons in MOs predict paramagnetism. | O2 is paramagnetic (2 unpaired electrons in π2p*). |
| Antibonding electrons reduce bond order and molecule stability. | Careful when asked about stability and existence. |
Downloadable Notes and Practice for Final Revision
- Summary PDFs: Quickly review main points, diagrams, and formulae for molecular orbital theory for O2, N2, B2 etc.
- Practice Molecular Orbital Theory MCQs in Chemical Bonding and Molecular Structure Mock Test and Practice Paper.
- Attempt worked problems in Atomic Structure Mock Test for MO diagrams and electron filling.
- Strengthen fundamental connections with Hybridization and Bond Order articles.
- Refer to Revision Notes for fast, last-minute summarisation.
Further JEE Resources Connecting to Molecular Orbital Theory
- For a deep dive into all bonding models, see Chemical Bonding and Molecular Structure.
- Test your understanding with JEE Chemistry Question Sets covering MOT, JEE Main Paper formats, and Mock Test Series.
- Clarify key MO diagram rules with Sigma and Pi Notation and explore Molecular vs. Hybrid Orbitals.
- Understand atomic orbital energy principles in Atomic Structure.
- See application to periodic trends in Periodicity Practice Paper.
- Advance into topics like p-Block Elements where advanced MO explanations are common in JEE questions.
To master molecular orbital theory for JEE, focus on stepwise MO diagrams, bond order formulas, and recognition of magnetic properties. Vedantu resources and expert-reviewed notes align with NCERT and JEE Main exam standards, supporting crystal-clear revision and success.
FAQs on Molecular Orbital Theory Explained: Concepts, Diagrams & Exam Guide
1. What is molecular orbital theory in chemistry?
Molecular Orbital Theory (MOT) is a quantum mechanical explanation of chemical bonding that describes electrons as being delocalized over entire molecules, not just individual atoms.
Key points include:
- Atomic orbitals combine to form molecular orbitals that extend over the whole molecule.
- These orbitals are classified as bonding and antibonding molecular orbitals.
- MOT explains bond order, magnetic properties, and stability of molecules more accurately than older models.
2. What are the main postulates or points of MOT?
The main points of Molecular Orbital Theory (MOT) are summarized as follows:
- Atomic orbitals combine to form molecular orbitals only if they have similar energy and symmetry.
- There are equal numbers of bonding and antibonding molecular orbitals formed.
- Electrons fill molecular orbitals in order of increasing energy, following the Aufbau principle, Pauli exclusion principle, and Hund's rule.
- Bond order is calculated as: ½ (number of electrons in bonding orbitals – number in antibonding orbitals).
- MOT predicts bond strength, length, magnetic character, and stability of molecules.
3. How is a molecular orbital formed?
Molecular orbitals are formed by the linear combination of atomic orbitals (LCAO) with similar energies and compatible symmetry as two atoms approach each other.
Types include:
- Bonding molecular orbitals (lower energy than atomic orbitals, increased electron density between nuclei)
- Antibonding molecular orbitals (higher energy, node between nuclei)
4. What is the difference between bonding and antibonding molecular orbitals?
Bonding and antibonding molecular orbitals differ mainly by their energy and electron density location:
- Bonding orbitals: Lower energy, electron density between nuclei, stabilizes the molecule.
- Antibonding orbitals: Higher energy, node (zero electron density) between nuclei, destabilizes the molecule.
5. What are the basic assumptions of molecular orbital theory?
The basic assumptions of MOT are:
- Atomic orbitals combine only if their energies are similar and symmetry is compatible.
- The number of molecular orbitals equals the number of combining atomic orbitals.
- Electrons fill molecular orbitals according to standard filling rules.
- Electrons are considered delocalized over the entire molecule, not attached to specific atoms.
6. How does MOT explain the paramagnetism of O2?
MOT explains that O2 is paramagnetic because its molecular orbital diagram shows two unpaired electrons in the π* (pi antibonding) orbitals.
This means:
- Presence of unpaired electrons leads to attraction by a magnetic field (paramagnetism).
- Valence bond theory does not explain this property, but MOT does.
7. What is a molecular orbital in simple terms?
A molecular orbital (MO) is a region in a molecule where there is a high probability of finding an electron, formed by the combination of atomic orbitals from bonded atoms. It is not limited to a single atom, but covers the entire molecule, allowing electron delocalization.
8. How do you calculate bond order using molecular orbital theory?
Bond order in MOT is calculated using the formula:
Bond order = ½ (Number of electrons in bonding MOs – Number in antibonding MOs)
Steps:
- Write the electronic configuration in molecular orbitals.
- Count electrons in bonding and antibonding MOs.
- Apply the formula to get bond order, which predicts bond strength and stability.
9. State the applications of molecular orbital theory.
MOT has several key applications in chemistry:
- Predicting molecular stability and bond order
- Explaining magnetic properties (e.g., paramagnetism and diamagnetism)
- Understanding the electronic configuration of molecules
- Predicting reactivity and bond length
10. What are the types of molecular orbitals?
The main types of molecular orbitals are:
- Sigma (σ) molecular orbitals: Formed by head-to-head overlap of atomic orbitals; symmetrical about the bond axis.
- Pi (π) molecular orbitals: Formed by side-to-side overlap of atomic orbitals; electron density above and below the bond axis.
- Each type can be bonding or antibonding depending on the phase combination.
11. Why does MOT sometimes predict bond order as zero for existing molecules?
MOT predicts bond order as zero when the number of electrons in bonding and antibonding molecular orbitals are equal, indicating no net bonding.
In such cases:
- The molecule is unstable and should not exist (e.g., He2 by MOT predicts bond order = 0).
- Sometimes, in reality, weak bonds or van der Waals interactions might exist, but MOT focuses on covalent bonding.
