Atomic Emission Lab
Explore how atoms emit light at specific wavelengths. Use the Rydberg equation to predict hydrogen spectral lines across the Lyman, Balmer, and Paschen series. Observe flame test colors for common elements and connect quantum energy levels to visible light.
Guided Experiment: Hydrogen Balmer Series Investigation
If electrons transition from higher energy levels (n=3 to n=7) down to n=2 in hydrogen, what wavelengths of light do you predict will be emitted? Will the lines be evenly spaced in the spectrum?
Write your hypothesis in the Lab Report panel, then click Next.
Controls
Emission Results
Energy Level Diagram
Horizontal lines show hydrogen energy levels. Blue arrows show electron transitions for the selected series. Numbers are wavelengths in nm.
Data Table
(0 rows)| # | Trial | Element | n (upper) | n (lower) | Wavelength(nm) | Energy(eV) | Color |
|---|
Reference Guide
Atomic Emission
When electrons in an atom fall from a higher energy level to a lower one, the energy difference is released as a photon of light. Each element has a unique set of energy levels, producing a characteristic emission spectrum — a fingerprint of bright lines.
For hydrogen, energy levels are given by:
where n = 1, 2, 3, ... is the principal quantum number. The ground state (n=1) has the lowest energy at -13.6 eV.
Bohr Model
The Bohr model describes hydrogen as an electron orbiting a proton in fixed circular orbits. Each orbit corresponds to a quantum number n. Electrons absorb energy to jump to higher levels and emit photons when falling back.
The photon energy equals the energy level difference:
where h = 6.626 × 10-34 J·s and c = 3 × 108 m/s.
Rydberg Equation
The Rydberg equation predicts the wavelengths of spectral lines in hydrogen:
where RH = 1.097 × 107 m-1, n1 is the lower level, and n2 > n1 is the upper level.
- Lyman series: n1 = 1 (UV)
- Balmer series: n1 = 2 (visible)
- Paschen series: n1 = 3 (IR)
Flame Tests
Flame tests identify elements by the color of light they emit when heated. Metal ions absorb heat energy, exciting electrons to higher levels. As electrons fall back, they emit photons at characteristic wavelengths.