Absorption & Emission of Light—Comprehensive Study Notes

Bohr Model Foundations

• Electrons occupy stable, discrete energy levels (orbits)
  • Energy levels are quantized; no in-between energies allowed.
• Photon absorption promotes an electron from a lower to a higher orbit if, and only if, the photon carries exactly the required energy difference.
  • Energy condition: $E_{photon}=h f=\Delta E$ where $h$ is Planck’s constant and $f$ is photon frequency.
  • If E_{photon}<\Delta E, the transition is forbidden; the electron stays put.
• Photon emission occurs when an electron falls from a higher to a lower orbit.
  • Emitted photon energy: $E_{emitted}=\Delta E$ (same magnitude as the absorption gap).
• Connection to Ch. 1 review: understanding single-electron behavior is the conceptual springboard for multi-electron, molecular, and spectroscopic phenomena tested on the MCAT.

Atomic & Molecular Absorption/Emission Beyond Hydrogen

• Real-world chemistry deals with poly-electron atoms and molecules whose electron distributions form molecular orbitals (MOs).
  • Transitions can involve bonding, antibonding, or non-bonding MOs, giving richer spectra than the simple Bohr picture.

Infrared (IR) Spectroscopy

• Used extensively in organic chemistry to identify functional groups.
• Principle: Different bond vibrations (stretching, bending, twisting) absorb characteristic IR frequencies.
• Absorption → bond’s vibrational mode is excited; output is an IR spectrum (peaks represent vibrational transitions).

UV–Vis Spectroscopy

• Extends concept to ultraviolet and visible wavelengths.
• Monitors electronic (rather than vibrational) transitions between MOs.
• Especially sensitive to:
  • Conjugated π systems
  • Aromatic rings
  • Transition-metal complexes (d–d or charge-transfer transitions)

How Absorption Spectra Are Displayed

• Two common conventions:
  1. Color bar: continuous rainbow with black lines where light is absorbed (missing colors).
  2. Graph: absorption (y-axis) vs. wavelength (x-axis).
  • Peaks → wavelengths where absorbance is greatest.
• Figure (mentioned) shows Earth’s atmosphere absorption across entire EM spectrum—illustrates selective transparency/opacity zones (e.g., ozone absorbing UV-B).

Structural Changes ⇒ Spectral Shifts

• Small molecular modifications (e.g., protonation state) can dramatically shift λ_max (wavelength of maximum absorbance).
• Example: Indicator “phenylphenolonin” (transcript typo; context implies phenolphthalein):
  • Acidic (protonated) form: colorless → absorbs no visible light.
  • Basic (deprotonated) form: bright pink → absorbs all wavelengths except long-red region.
  • Color observed is the complement of absorbed colors (we see what is NOT absorbed).
• Underlying structural reason: protonation alters conjugation length & electron distribution, shifting energy gap between ground & excited MOs.
• General rule: More conjugation → smaller ΔE → absorption shifts to longer λ (lower energy, visible range).

Fluorescence

• Fluorescence = photoluminescence involving multi-step relaxation.
• Process steps:
  1. Excitation: Absorb a high-energy photon (often UV); electron jumps to an excited electronic state.
  2. Internal conversion/vibrational relaxation: Part of energy lost non-radiatively as heat or coupling to lattice.
  3. Emission: Electron returns to ground (or lower excited) state in two or more steps → emits photons of lower frequency, longer wavelength than original UV photon.
• If emitted wavelength lies in visible region (≈400–700 nm), the material glows in a visible color.
• Common fluorescent substances: rubies, emeralds, phosphors in fluorescent bulbs, neon-sign gases.
• Wide array of colors in fluorescent & neon lighting comes from distinct, element- or compound-specific multi-step emission spectra.

Practical & MCAT-Relevant Takeaways

• Memorize the quantitative relationship $E=h f$ and know how to convert between energy, frequency, wavelength ($\lambda=c/f$).
• Recognize spectra as diagnostic fingerprints:
  • IR → functional groups.
  • UV–Vis → conjugation, transition-metal ions, indicators, biological pigments (e.g., heme, chlorophyll).
• Concept check: If a substance appears green, which colors are being absorbed? (Answer: mainly red and violet/blue; green is transmitted/reflected.)
• Ethical/practical note: Atmospheric absorption spectra dictate issues like ozone depletion (UV shielding) and greenhouse gas trapping (IR absorption).

Key Equations & Definitions Recap

• Photon energy: $E= h f =\dfrac{h c}{\lambda}$
• Energy difference between two Bohr levels: $\Delta E=E<em>{n</em>f}-E<em>{n</em>i}= -\dfrac{13.6\,\text{eV}}{n<em>f^2}+\dfrac{13.6\,\text{eV}}{n</em>i^2}$ (hydrogenic atoms).
• Absorption spectrum: set of wavelengths where sample absorbs EM radiation.
• Emission spectrum: set of wavelengths emitted when excited sample relaxes.
• Fluorescence: rapid (10^{-9}–10^{-7} s) emission following UV excitation; obeys Stokes shift (emission λ>absorption λ).

Study Tips

• Draw parallel arrows between electronic transitions (Bohr), vibrational transitions (IR), and color changes (visible) to solidify vertical integration.
• Practice with real spectra: annotate an IR or UV–Vis printout marking λ_max, intensity, and relate features to molecular structure.
• Use indicator color charts to visually connect pH ↔ protonation state ↔ conjugation ↔ absorption ↔ perceived color.