UV-Visible Spectroscopy and Molecular Orbitals

Overview of UV-Vis Spectroscopy

• Purpose: Identify wavelengths absorbed by a compound.
• Instrument: UV-Vis spectrophotometer.
• Wavelength Range: 200 - 800 nanometers.
• Output: Absorption spectrum.

Example: 1,3-Butadiene

• Absorption Peak: ~217 nanometers (lambda max).
• UV Region: No color observed (colorless).

Structural Composition

• Carbon Hybridization: Each of the four carbons is sp2 hybridized.
• Orbital: Presence of four p orbitals.
• Molecular Orbital Formation:
  • 4 atomic orbitals combine into:
    • 2 bonding molecular orbitals
    • 2 anti-bonding molecular orbitals
• Energy Levels:
  • Bonding orbitals lower in energy.
  • Anti-bonding orbitals higher in energy.

Electron Configuration

• Pi Electrons: 4 pi electrons.
• Filling Order: Electrons fill lowest energy orbitals first.
• Ground State: All pi electrons in bonding molecular orbitals.

Excited State

• Light Interaction: Absorption of energy by pi electron.
• Electron Transition: HOMO to LUMO transition.
• Energy and Wavelength Relation:
  • Energy of photon = Planck's constant (h) x frequency (\nu).
  • Frequency (\nu) = speed of light / wavelength (\lambda).
  • Energy (E) = (h \times c / \lambda).
  • Inverse relationship between energy and wavelength.

Molecular Orbital: Ethanol

• Pi Electrons: 2 pi electrons.
• Orbital Transition: Pi to pi star transition.
• Wavelength: Approx. 180 nanometers (often below UV-Vis range).

Non-bonding Electrons

• Orbital: Non-bonding orbitals are higher in energy than bonding orbitals.
• Transition: n to pi star transition.
• Wavelength: Approx. 290 nanometers.

Energy Wavelength Relationship

• Energy Difference: Smaller difference between orbitals leads to longer wavelength absorption.
• Color Implications: Decreasing energy difference increases absorbed wavelength, important for color understanding.

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Further exploration into these concepts will discuss the implications of color and wavelength relationships.