Atomic Structure and Spectra · AtomLab

The alpha-scattering experiment

Thomson pictured the atom as a uniform ball of positive charge with electrons embedded in it like raisins (the “plum pudding” model). If that were so, alpha particles passing through gold foil would be deflected only very slightly.

Geiger and Marsden fired alpha particles from a radioactive source at very thin gold foil and counted the flashes on a zinc sulfide screen through a microscope: almost all alpha particles went nearly straight on, a few were deflected through large angles, and a very few (about one in eight thousand) turned through more than 90° — some almost straight back. From this Rutherford proposed the nuclear model of the atom.

Turn the screen and count for one minute at several angles; then dive to the atomic scale and wait for a head-on collision. How electrons are deflected in electric fields is shown in FieldLab.

Hydrogen spectrum

Hydrogen in a discharge tube glows a rosy violet. Through a grating its light is not a continuous rainbow but a few separate bright lines — a line spectrum. Every element has its own lines, like a fingerprint.

The grating has 600 lines per millimetre, and first-order lines satisfy d sin θ = λ. Turn the telescope until a line sits on the cross-hairs in the eyepiece, record the angle, and the wavelength follows. Balmer found that hydrogen's four visible lines obey 1/λ = R(1/2² − 1/n²); R is called the Rydberg constant.

How a grating works is shown in OpticsLab; the energy hν of a single photon is in PhotonLab.

Bohr's model and transitions

Bohr proposed that electrons can move only in certain orbits, and that an atom in one of these stationary states does not radiate. The energy of the hydrogen atom is Eₙ = −13.6 eV / n², and the orbit radius is rₙ = n²r₁.

When an atom jumps from a higher level m to a lower level n it emits one photon, hν = Eₘ − Eₙ; conversely, only a photon whose energy exactly equals a level difference can be absorbed. The jumps that end on n = 2 are precisely the Balmer lines of lesson 2. With more than 13.6 eV the electron is knocked right out of the atom (ionisation).

Click a level on the energy diagram to fire exactly the right photon. Bohr's model cannot handle atoms with many electrons; what electrons really look like is shown by the electron clouds in OrbitalLab.

The Franck–Hertz experiment

Electrons from the hot filament (the cathode) are accelerated by the accelerating voltage U; after passing the grid they must also overcome a small retarding voltage to reach the plate. The tube contains mercury vapour (or neon).

While an electron's kinetic energy is less than mercury's first excitation energy, 4.9 eV, its collisions with mercury atoms are elastic and it loses almost no energy. Once it reaches 4.9 eV it hands that whole amount to a mercury atom, almost stops, fails to reach the plate, and the current drops. So the current drops about every 4.9 V. The excited mercury atoms give off 253.7 nm ultraviolet light (photons of 4.9 eV).

With a neon tube the excitation energy is about 18.7 eV and excited neon atoms give off orange light, so you can see bands of light in the tube directly. Photon energies are covered in PhotonLab.

In-class testing

Look back over the four topics and answer the multiple-choice questions below.