Bohr atom light lab

When atoms absorb energy, electrons jump to higher shells; when they fall back, they emit light of specific colors. This simulation lets you excite atoms and watch the photons fly, connecting energy levels to the bright lines of atomic spectra. Learning goals: describe how electron transitions produce emission spectra; calculate photon energies from shell transitions in hydrogen; explain why each element's spectrum is unique.

Atomic number: 6

Shell occupancy: 2 · 4

Simplified Bohr model — real electron arrangements follow subshell (s, p, d, f) filling rules.

Controls

Element selector: Chooses the element shown (hydrogen through calcium).
Excite electron: Adds a photon of chosen energy to excite the electron.
Show energy levels: Toggles labels showing shell numbers and energies.
Spectrum view: Displays emitted photons as colored spectral lines.

The chemistry

E_n = -\frac{13.6\ \text{eV}}{n^2}

Bohr energy levels for hydrogen

\Delta E = 13.6\left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right)\ \text{eV}

Photon energy from a transition

E = \frac{hc}{\lambda}

Photon energy and wavelength

Assumptions

The Bohr model is a simplified teaching model, not the full quantum picture.
Exact Bohr energy levels apply strictly to one-electron species like hydrogen.
Electrons do not literally orbit the nucleus like planets.

Guided activity

Select hydrogen and excite the electron to n = 3, then watch it fall back.
Count the emitted photons and match each to a spectral line color.
Compare the emission spectrum of hydrogen with that of helium.
Excite an electron with too little energy and observe that nothing happens.

Challenge questions

Predict the color of light emitted in the n = 3 to n = 2 transition before running it.
Explain why each element has a unique set of spectral lines.
Describe one observation the Bohr model explains well and one it cannot.