which of the following electronic transitions in the oxygen atom will result in light emission?
Electronic transitions that emit light occur when an electron drops from a higher energy level (n_i > n_f) to a lower one in the oxygen atom, releasing a photon. This fundamental quantum process follows the same rules as in hydrogen-like atoms but accounts for oxygen's multi-electron configuration.
Key Principle
Light emission requires n_initial > n_final because energy is released as the electron relaxes. Absorption (no emission) happens when n_initial < n_final, as energy is absorbed from incoming light.
Common Oxygen Emission Lines
Oxygen atoms produce distinct emission spectra from excited states:
- 777.5 nm and 844.7 nm : From 3p → 3s transitions in neutral oxygen (O I).
- 715 nm and 799 nm : Intra-atomic drops like 3p → 3s in oxygen during interfacial processes.
- These match observed peaks in experiments, confirming downward transitions emit visible/NIR light.
Typical Multiple-Choice Answer
In standard questions (e.g., from forums/quizzes), options like:
- A. n=1 → n=2 (absorption, no emission)
- B. n=1 → n=3 (absorption)
- C. n=2 → n=1 (emission)
- D. n=3 → n=2 (emission)
The transitions resulting in emission are C and/or D (n=3→2 or n=2→1). Upward jumps absorb light instead.
Transition| Direction| Result| Example Wavelength
---|---|---|---
n=3 → n=2| Down| Emission| ~700-800 nm 1
n=2 → n=1| Down| Emission| Shorter UV-visible
n=1 → n=2| Up| Absorption| No light emitted
n=1 → n=3| Up| Absorption| No light emitted 8
TL;DR: Downward transitions (higher n to lower n) emit light in oxygen atoms, like n=3→2. Real spectra confirm this in labs and auroras.
Information from public scientific sources and spectra data.