Organic Chemistry · Conjugated Compounds and Ultraviolet Spectroscopy
Conjugation, Color, and the Chemistry of Vision
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In 30 seconds
When a molecule absorbs light, an electron jumps from the highest occupied molecular orbital (HOMO) to the lowest unoccupied molecular orbital (LUMO). For an isolated C=C double bond that π → π* transition costs roughly 7 eV — far more energy than a visible photon can supply, which is why ethene and most simple alkenes are colorless. conjugation Alternating single and double bonds that let π electrons delocalize across several atoms Full entry → changes this. In alternating single–double bond systems, the π electrons delocalize across the chain: bonding orbitals drop in energy while antibonding orbitals rise, so the HOMO–LUMO gap shrinks. Once the gap matches a wavelength in the visible range (about 380–750 nm), the molecule absorbs visible light and appears colored.
This topic connects three ideas: why extended conjugation shifts absorption to longer wavelengths, how that shift makes invisible molecules into colored dyes, and how biology uses the same chemistry in the retina, where 11-cis-retinal Conjugated polyene aldehyde derived from vitamin A that captures light Full entry → captures a photon and starts the process we experience as sight.
Why this matters
Color is conjugated-molecule photochemistry in action. β-Carotene (11 conjugated double bonds) absorbs blue light near 453 nm and therefore looks orange — the reason carrots, pumpkins, and egg yolks are colored. Lycopene, with a still longer conjugated chain, makes tomatoes red. Synthetic dyes and textile colorants work on the same principle: tune the conjugated system until its absorption lands in the visible window.
The same physics runs human vision. Photoreceptor cells contain rhodopsin Opsin protein covalently bound to 11-cis-retinal Full entry →, a protein holding 11-cis-retinal — a vitamin A derivative with a conjugated polyene chain — inside a binding pocket. A photon isomerizes retinal from 11-cis to all-trans, the protein changes shape, and a signaling cascade sends a nerve impulse to the brain. Because retinal is built from vitamin A, dietary deficiency causes night blindness — a direct link between nutrition and molecular photochemistry. Understanding conjugation is also the key to interpreting UV–Vis spectra (topics 7–8) and to designing sunscreens and photodynamic therapy agents.
The college version
Core Concepts
The HOMO–LUMO gap shrinks as conjugation grows
Each π bond contributes one bonding and one antibonding orbital; in a conjugated system these combine into delocalized molecular orbitals. The lowest-energy absorption is the HOMO → LUMO transition, with energy:
ΔE = ELUMO - EHOMO
For ethene the gap is large; for 1,3-butadiene smaller; for hexatriene smaller still. Each additional conjugated double bond narrows the gap and moves the absorption maximum (λmax) to longer wavelength.
From ultraviolet to visible: when a molecule becomes colored
The wavelength absorbed is fixed by the photon-energy relation:
ΔE = hcλ
where h is Planck's constant (6.626 × 10-34 J · s), c is the speed of light (2.998 × 108 m/s), and λ is the wavelength. Ethene absorbs near 170 nm, butadiene near 217 nm, hexatriene near 258 nm — all ultraviolet, so colorless. A molecule looks colored only when its λmax falls inside the visible window; it then removes that color from the light reaching your eye, and you perceive the complementary color. β-Carotene absorbs blue (~453 nm) and appears orange; a compound absorbing green appears red or magenta. A shift of λmax to longer wavelength (a bathochromic shift Movement of λmax to longer wavelength Full entry →) is the signature of extended conjugation.
A simple picture: electrons in a box
A useful model treats the delocalized π electrons as particles moving freely along a one-dimensional "box" whose length L is roughly the conjugated chain length. The particle-in-a-box gap scales as:
ΔE ∝ h28mL2
with m the electron mass: a longer box gives a smaller gap and longer absorption wavelength. The model ignores electron–electron repulsion, but captures the central trend: conjugation length controls color.
The chemistry of vision: retinal and rhodopsin
Vertebrate rod cells pack rhodopsin, an opsin protein covalently bound to 11-cis-retinal through a protonated imine (Schiff base) linkage. The retinal chain is a conjugated polyene with six C=C bonds held nearly planar in the binding pocket. Free retinal absorbs near 380 nm (ultraviolet), but in rhodopsin the absorption shifts to about 500 nm (blue-green): charged residues and the protonated Schiff base stabilize the excited state, lowering the transition energy. This "opsin shift Change in retinal's absorption caused by the protein environment Full entry →" moves absorption into the visible range so rod cells can catch dim light.
Absorbing a photon is a photochemical event: the energy rotates the C11=C12 double bond, converting the bent 11-cis isomer into the extended all-trans isomer within a few hundred femtoseconds — fast, stereospecific, and driven by light, not heat. The shape change forces opsin to rearrange, activating a G-protein cascade that produces a nerve signal. Color vision uses the same chemistry with three cone opsins tuned to ~420, ~530, and ~560 nm.
How It Works / Step-by-Step Process
To predict whether a conjugated molecule is colored:
- Count the conjugated double bonds (or π electrons) in the chain.
- Estimate the HOMO–LUMO gap: more conjugation → smaller gap.
- Convert the gap to a wavelength with ΔE = hc/λ.
- Compare with the visible window (380–750 nm): if λmax falls inside, the molecule is colored and appears as the complement of the absorbed color.
- For vision problems: identify the chromophore (11-cis-retinal), recognize that a photon drives cis → trans isomerization, and trace the shape change to the nerve signal.
Common Confusions
| Do not confuse | With | Difference |
|---|---|---|
| The color you see | The color the molecule absorbs | A tomato looks red because it absorbs green-blue light; you see the complementary light that is not absorbed |
| "More double bonds" | "Always colored" | Only when λmax reaches the visible range; butadiene and hexatriene absorb in the UV and are colorless |
| Conjugation length | Molecular size or molar mass | The delocalization length, not the whole molecule, sets the gap; a long alkyl tail adds no color |
| λmax | The perceived color | λmax is what is absorbed; perceived color is the complement |
| The photochemical cis → trans change in vision | A thermal (heat-driven) isomerization | Vision's isomerization is photon-driven, fast, and stereospecific; it happens without heat |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Imagine electrons as marbles in a groove. A short groove has steep ends, so a marble needs a hard push to jump the bump — a big energy gap and invisible ultraviolet light. A longer groove is shallower, so a gentler push works — a smaller gap and visible colored light. Carrots look orange because their long-groove molecule (β-carotene) catches blue light. Your eyes work the same way: a long-groove molecule called retinal catches a light particle, twists, and sends a "I saw light!" message to your brain.
Worked example
Example 1: How much energy does β-carotene absorb at 453 nm?
β-Carotene's strongest absorption is at λ= 453 nm. Start with the photon-energy formula:
E = hcλ
Substitute h = 6.626 × 10-34 J · s, c = 2.998 × 108 m/s, and λ= 453 × 10-9 m:
E = (6.626 × 10-34 J · s)(2.998 × 108 m/s)453 × 10-9 m = 4.39 × 10-19 J
Dimensional analysis: J · s × m/s ÷ m = J ✓. Convert to a per-mole basis with Avogadro's number:
Emol = (4.39 × 10-19 J)(6.022 × 1023 mol-1) = 2.64 × 105 J/mol = 264 kJ/mol
A blue photon delivers ~264 kJ/mol — enough to promote an electron but far below bond-dissociation energies, so absorption does not break the molecule apart.
Example 2: Why does binding to opsin shift retinal's absorption?
Free 11-cis-retinal absorbs near 380 nm; in rhodopsin it absorbs near 500 nm. Compute both photon energies with the same formula:
E = hcλ
At 380 nm:
E380 = (6.626 × 10-34)(2.998 × 108)380 × 10-9 = 5.23 × 10-19 J
At 500 nm:
E500 = (6.626 × 10-34)(2.998 × 108)500 × 10-9 = 3.97 × 10-19 J
The gap shrinks by:
Δ(ΔE) = 5.23 × 10-19 - 3.97 × 10-19 = 1.26 × 10-19 J ≈ 76 kJ/mol
The protein environment stabilizes the excited state more than the ground state, lowering the transition energy by roughly 76 kJ/mol and moving the absorption into the visible blue-green — the range where rod cells need maximum sensitivity.
Key takeaways
- Conjugation delocalizes π electrons and shrinks the HOMO–LUMO gap; longer chains → longer λmax (bathochromic shift).
- A molecule appears colored only when it absorbs in the visible range (~380–750 nm); perceived color is complementary to absorbed color.
- Photon energy and wavelength: ΔE = hc/λ — longer wavelength = less energy.
- β-Carotene: 11 conjugated double bonds, λmax ≈ 453 nm — absorbs blue, appears orange.
- Vision: 11-cis-retinal (vitamin A derivative, conjugated polyene, bound as a protonated Schiff base) absorbs a photon and photoisomerizes to all-trans-retinal, triggering the visual cascade.
- The protein environment shifts retinal's absorption from ~380 nm (free) to ~500 nm (in rhodopsin) — the "opsin shift".
- Dyes, sunscreens, and photodynamic therapy agents all exploit conjugation-tuned absorption.
Check yourself
6 review questions from the chapter. Try each one, then open the answer.
Why does extending a conjugated chain shift the absorption maximum to longer wavelengths?
Show answer
Conjugation delocalizes the π electrons, raising the HOMO and lowering the LUMO. The smaller gap means a lower-energy photon (longer wavelength) drives the transition.
β-Carotene absorbs near 453 nm. Is that in the visible range, and what color does β-carotene appear?
Show answer
Yes — 453 nm lies in the visible window (380–750 nm). β-Carotene absorbs blue light, so it transmits the complementary light and appears orange.
Compute the photon energy (in J and in kJ/mol) for light of 453 nm. Use h = 6.626 × 10-34 J · s, c = 2.998 × 108 m/s, NA = 6.022 × 1023 mol-1.
Show answer
E = hc/λ= (6.626 × 10-34)(2.998 × 108)/(453 × 10-9) = 4.39 × 10-19 J; per mole, 4.39 × 10-19 × 6.022 × 1023 = 2.64 × 105 J/mol = 264 kJ/mol.
What structural change does an absorbed photon cause in 11-cis-retinal, and why does that change matter for vision?
Show answer
The photon drives photoisomerization of the C11=C12 double bond from 11-cis to all-trans. The shape change forces opsin to rearrange, triggering the cascade that produces a nerve impulse — the first chemical step of vision.
A dye absorbs strongly at 620 nm (orange-red). What color will it appear?
Show answer
It absorbs orange-red light, so it appears as the complement — blue-green (cyan).
Why is vitamin A deficiency associated with night blindness?
Show answer
Retinal is synthesized from vitamin A (retinol). Without enough dietary vitamin A, the retina cannot make enough rhodopsin, so dim-light (night) vision fails.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- conjugation
- Alternating single and double bonds that let π electrons delocalize across several atoms
- HOMO / LUMO
- Highest occupied / lowest unoccupied molecular orbital
- π → π* transition
- Promotion of a π electron into an antibonding orbital
- λmax
- Wavelength of maximum absorption
- bathochromic shift
- Movement of λmax to longer wavelength
- 11-cis-retinal
- Conjugated polyene aldehyde derived from vitamin A that captures light
- rhodopsin
- Opsin protein covalently bound to 11-cis-retinal
- opsin shift
- Change in retinal's absorption caused by the protein environment
- π → π^ transition
- Promotion of a π electron into an antibonding orbital
Sources & references
This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.
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