Atomic Structure

Syllabus Unit: Physical Chemistry

Why it matters

Chapter overview

This topic traces the model of the atom from Bohr's quantized planetary orbits to the full quantum-mechanical picture — the four quantum numbers that address every electron, the orbital shapes they describe, and the systematic rules (Aufbau, Pauli, Hund) that determine exactly how electrons fill an atom's orbitals.

Chemistry Syllabus Physical Chemistry ~25% weightage

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n = 1 (Ground State)E = -13.6 eV, closest to the nucleus

n = 2E = -3.4 eV

n = 3E = -1.51 eV

n = 4E = -0.85 eV

1Bohr's Model & Energy Levels

Bohr fixed the electron into quantized circular orbits, explaining the hydrogen spectrum's discrete lines but breaking down for multi-electron atoms.

Learning Objectives

State Bohr's postulates: quantized angular momentum, stationary orbits, and photon emission/absorption on orbit change.
Calculate the energy of an electron in the nth orbit of hydrogen using En = -13.6/n² eV.
Calculate orbit radius using rn = 0.529 n²/Z Å.
Explain the origin of hydrogen's line spectra (Lyman, Balmer, Paschen series) using electron transitions.

Key Terms

QuantizationBohr's postulate restricting angular momentum to integral multiples of h/2π: mvr = nh/2π.

Stationary OrbitAn orbit in which an electron can revolve without radiating energy, contrary to classical electromagnetic theory.

Ground StateThe lowest-energy orbit (n=1) an electron occupies in an unexcited atom.

Balmer SeriesHydrogen spectral lines from electron transitions ending at n=2 — the only series that falls partly in the visible region.

Rutherford's nuclear model couldn't explain why orbiting electrons don't simply spiral into the nucleus while radiating energy. Bohr fixed this by proposing that only certain, quantized orbits are allowed.

Bohr's three postulates

Electrons revolve in fixed circular orbits ('stationary states') without radiating energy.
Angular momentum is quantized: mvr = nh/2π, where n = 1, 2, 3...
Energy is absorbed or emitted only when an electron jumps between orbits: ΔE = hν = E2 − E1.

En = -13.6 Z²/n² eV rn = 0.529 n²/Z Å

The hydrogen spectrum

Each spectral series corresponds to electrons falling to a specific lower orbit: Lyman (falls to n=1, UV), Balmer (falls to n=2, visible), Paschen/Brackett/Pfund (fall to n=3/4/5, infrared).

Exam TipThe Rydberg formula 1/λ = RH(1/n1² − 1/n2²) with RH = 1.097×10⁷ m⁻¹ lets you calculate the wavelength of any emitted photon directly — worth memorising both the formula and RH.

Common MistakeBohr's model gives quantitatively correct results only for hydrogen-like (one-electron) species — H, He⁺, Li²⁺ — not for any multi-electron atom.

Limitations

Couldn't explain fine spectral splitting seen under magnetic (Zeeman) or electric (Stark) fields.
Assigned the electron a fixed orbit — a definite position and momentum simultaneously — violating the Heisenberg uncertainty principle.
Failed entirely to predict spectra for multi-electron atoms.