Chapter 4: AtomsClass 12 Physics — summary, notes, extra questions & MCQ quiz

Summary

Rutherford's alpha-scattering experiment showed that an atom's positive charge and nearly all its mass are concentrated in a tiny dense nucleus, with electrons orbiting in the surrounding space. However, classical physics predicted that orbiting electrons should radiate energy and spiral inward, making such an atom unstable—and it could not explain the sharp line spectra of gases. Bohr resolved this for hydrogen with three postulates: electrons occupy stationary orbits without radiating; angular momentum is quantised as \(mvr=n\dfrac{h}{2\pi}\); and radiation of frequency \(\nu=(E_i-E_f)/h\) is emitted or absorbed only during transitions between orbits. These yield quantised orbit radii and energies \(E_n=-\dfrac{13.6}{n^2}\) eV, explaining the hydrogen spectral series—Lyman, Balmer, Paschen and others—through the Rydberg formula. The energy-level diagram and the concepts of ground and excited states follow naturally. Though the Bohr model succeeds beautifully for hydrogen and hydrogen-like ions, it fails for multi-electron atoms and cannot explain spectral fine structure, foreshadowing the need for full quantum mechanics, yet it remains a milestone in understanding atomic structure.

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