Modern · Experiment

The Bohr Model of the Atom

An electron may orbit the nucleus only in certain fixed shells. When it jumps between them it emits or absorbs a single photon — and the colours of those photons are the atom's fingerprint spectrum.

Click an orbit or energy level — or use the buttons — to send the electron there. Jumping down emits a photon; jumping up absorbs one.

Controls

Current level n1
Energy Eₙ−13.60 eV
Orbit radius rₙ0.053 nm
Last transition—
Photon energy ΔE—
Wavelength λ—
Series—
Electron in the ground state (n = 1).
About this experiment ›

What you are looking at

A single electron circles a positively charged nucleus, but — unlike a planet, which could orbit at any distance — it is allowed only a discrete set of shells labelled n = 1, 2, 3, … On the right is the matching energy-level diagram, and along the bottom is the emission spectrum: the specific colours this atom can give off. Send the electron to a new level and watch a photon fly out (jumping down) or get absorbed (jumping up).

Bohr's quantum leap

In 1913 Niels Bohr proposed three rules that broke with classical physics. First, the electron occupies only certain stationary states with quantized angular momentum, and — crucially — does not radiate while in them (classically an orbiting charge should spiral inward and the atom collapse). Second, each state has a fixed energy:
Eₙ = −13.6 · Z² / n² eV
negative because the electron is bound, approaching 0 as n → ∞ (free). The orbit radius grows as rₙ = n²a₀/Z, with a₀ = 0.053 nm the Bohr radius. Third, the atom changes energy only by jumping between levels, emitting or absorbing one photon whose energy exactly bridges the gap:
ΔE = Eᵢ − E_f = h f = h c / λ

Why atoms have line spectra

Because only certain energy gaps exist, an atom emits only certain wavelengths — sharp spectral lines, not a continuous rainbow. For hydrogen these fall into series named by the final level: Lyman (down to n = 1, ultraviolet), Balmer (to n = 2, visible — the four coloured lines on the strip), and Paschen (to n = 3, infrared). The wavelengths obey the Rydberg formula 1/λ = R(1/n_f² − 1/n_i²). These fingerprints are how we identify elements in stars.

Things to try

Excite the electron to a high level and let it cascade down, watching each photon's colour. Find the Balmer lines (any jump landing on n = 2) and see them line up with the coloured spectrum below. Increase the nuclear charge Z for hydrogen-like ions (He⁺, Li²⁺): every energy deepens by Z² and the orbits shrink, pushing the spectrum toward the ultraviolet. Bohr's model works beautifully for one electron but fails for many-electron atoms — the door it opened led to full quantum mechanics.