In 1913, Niels Bohr proposed a model for the Hydrogen atom to explain its line spectrum. He introduced the idea that electrons move in fixed, non-radiating orbits.
This experiment provided the first direct evidence for quantized energy levels in atoms. By bombarding Mercury vapor with electrons, they observed that energy was absorbed only in specific discrete amounts.
The periodic drops in current at specific voltages (4.9V for Mercury) confirmed that electrons lose energy only when it matches a specific transition energy of the atom.
Arnold Sommerfeld extended Bohr's model by introducing elliptical orbits. This helped explain the "fine structure" of spectral lines.
The modern description of the atom uses four quantum numbers to define the state of an electron:
| Quantum Number | Symbol | Defines |
|---|---|---|
| Principal | n | Main shell/Energy level (1, 2, 3...) |
| Azimuthal (Orbital) | l | Subshell shape (0 to n-1) |
| Magnetic | ml | Orientation in space (-l to +l) |
| Spin | ms | Direction of spin (+1/2 or -1/2) |
This experiment proved the existence of Electron Spin and Space Quantization. By passing a beam of Silver atoms through a non-uniform magnetic field, the beam split into two distinct spots.
Conclusion: The splitting confirmed that electrons have an intrinsic magnetic moment due to spin, which can only have two orientations (up or down).
When high-speed electrons strike a target, they produce two types of X-rays:
Henry Moseley discovered that the frequency of characteristic X-rays is related to the atomic number (Z).
This law proved that Z (Atomic Number) is more fundamental than atomic weight for organizing the Periodic Table.
Tip: For the Frank-Hertz experiment, remember that the mean free path of the electron must be large enough to reach the excitation energy, which is why the experiment is conducted at low pressures.