Electronic transitions

 In spectroscopy, electronic transitions refer to the movement of an electron from one energy level to another within an atom or molecule. When a molecule absorbs energy (typically in the form of electromagnetic radiation), its electrons can be excited to higher energy levels. There are several types of electronic transitions, depending on the molecular orbitals involved. These include

1. σ → σ* (Sigma to Sigma Anti-bonding)

• Description: Involves the transition from a bonding σ orbital (lowest energy bonding orbital) to an anti-bonding σ* orbital.
• The energy required: Very high.
• Common in: Simple molecules with single bonds (e.g., alkanes).
• Example: Ethane (C–C single bond) requires UV light for this transition.
• Application: Generally observed in the vacuum UV region (below 200 nm).

2. n → σ* (Non-bonding to Sigma Anti-bonding)

• Description: An electron from a non-bonding orbital (lone pair) is excited to a σ* anti-bonding orbital.
• Energy required: Moderate.
• Common in: Molecules with atoms like oxygen, nitrogen, or halogens (having lone pairs).
• Example: Water (H₂O), where the oxygen's lone pair can transition to the σ* orbital.
• Application: UV-visible region.

3. π → π* (Pi to Pi Anti-bonding)

• Description: An electron from a bonding π orbital (from a double bond) is excited to a π* anti-bonding orbital.
• Energy required: Moderate.
• Common in: Molecules with conjugated double bonds (e.g., alkenes, aromatic compounds).
• Example: Ethylene (C=C double bond).
• Application: Typically observed in the UV and visible range, particularly in organic compounds with conjugation.

4. n → π* (Non-bonding to Pi Anti-bonding)

• Description: A non-bonding electron (from a lone pair) is excited to a π* anti-bonding orbital.
• Energy required: Relatively lower energy compared to other transitions.
• Common in: Carbonyl compounds, amines, or other groups with lone pairs adjacent to unsaturated bonds.
• Example: Formaldehyde (H₂C=O), where oxygen's lone pair transitions to a π* orbital.
• Application: Occurs in the UV-visible region, often leading to significant peaks in absorption spectra.

5. Charge Transfer (CT) Transitions

• Description: Involves the transfer of an electron between two species, usually from a donor to an acceptor, within the same molecule or between molecules.
• The energy required: Varies.
• Common in: Metal complexes or conjugated organic systems.
• Example: In transition metal complexes, where the electron moves from a ligand to the metal or vice versa.
• Application: Often results in intense absorptions in the UV-visible spectrum.

6. d → d Transitions

• Description: Involves the promotion of an electron from one d-orbital to another in transition metal complexes.
• The energy required: Lower than charge transfer transitions.
• Common in: Transition metal ions where the d orbitals are split by ligand field effects (e.g., octahedral or tetrahedral complexes).
• Example: The color of many transition metal complexes is due to d → d transitions.
• Application: Visible range; explains the color of many transition metal complexes.

7. f → f Transitions

• Description: These occur in the f-orbitals of lanthanides and actinides.
• The energy required: Relatively small energy differences.
• Common in: Rare earth and actinide elements.
• Example: Europium (Eu³⁺) and Terbium (Tb³⁺) compounds exhibit f → f transitions.
• Application: Luminescent properties, often used in lighting and display technologies.

Summary of Energy Hierarchy:

• σ → σ* > n → σ* > π → π* > n → π* > d → d > f → f

This hierarchy reflects the typical energy required for these transitions. The highest energy transitions (σ → σ*) occur in the far UV, while others can occur in the visible or near UV region. Each transition gives rise to different absorption bands, which are useful in techniques like UV-Vis spectroscopy for identifying molecular structures.