Electronic Spectra: 20 flashcards
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1. What is an electronic spectrum?
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The pattern of light absorbed or emitted when electrons jump between quantised energy levels in atoms or molecules.
2. Which equation relates photon energy to the energy gap?
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$h\nu = \Delta E$
3. In which region of the electromagnetic spectrum do electronic transitions typically occur?
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Ultraviolet (UV) and visible (and sometimes near-infrared).
4. What is the difference in appearance between absorption and emission spectra?
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Absorption: dark lines/bands on a bright background. Emission: bright lines/bands on a dark background.
5. Why do atomic spectra consist of sharp lines?
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Because atomic energy levels are discrete (quantised), so transitions occur at specific energies.
6. Why do molecular electronic spectra appear as broad bands?
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Each electronic level is split into many vibrational and rotational sub-levels, so many closely spaced lines merge into a band.
7. State the Franck–Condon principle.
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The most intense vibronic transition is the one whose upper-state vibrational wavefunction overlaps most with the lower-state one, usually a vertical transition because nuclei move slowly during the electronic jump.
8. What is the spin selection rule for electronic transitions?
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$\Delta S = 0$ — the total spin must not change.
9. State the Laporte rule.
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In centrosymmetric molecules, $g \leftrightarrow g$ and $u \leftrightarrow u$ transitions are forbidden; only $g \leftrightarrow u$ is allowed.
10. Are forbidden transitions impossible? Explain.
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No, they are weak, not impossible; they can appear through mechanisms like spin–orbit coupling.
11. Write the Beer–Lambert law.
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$A = \varepsilon c l$
12. What is a chromophore?
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The structural unit of a molecule responsible for its colour (e.g., a conjugated $\pi$ system).
13. How does extending conjugation affect absorption?
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It lowers the HOMO–LUMO gap, causing a red shift (absorption at longer wavelengths).
14. How is the observed colour of a substance related to the wavelengths it absorbs?
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The observed colour is complementary to the absorbed wavelengths.
15. Name the hydrogen series that fall in the UV, visible, and IR regions.
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Lyman (UV), Balmer (visible), Paschen (IR).
16. What does the Rydberg formula predict?
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The positions of hydrogen spectral lines: $1/\lambda = R(1/n_1^2 - 1/n_2^2)$.
17. What are the main components of a UV–visible spectrophotometer?
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Light source, monochromator, sample, detector.
18. What do strong bands ($\varepsilon \sim 10^4$) usually indicate?
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Allowed $\pi \to \pi^*$ transitions.
19. What do weak bands ($\varepsilon \sim 10$) often indicate?
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Forbidden $n \to \pi^*$ or spin-forbidden transitions.
20. Give two real-world applications of electronic spectra.
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Ozone UV shielding and photosynthesis (chlorophyll absorption); also forensics, astronomy, device design.
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