Niels Bohr Uses Quantum Mechanics to Explain Atomic Structure
The Bohr model of the atom was a major leap forward in our understanding of atomic structure and quantum mechanics, but it was not the final answer.
Early Theories of Atomic Structure
In 1897, J. J. Thomson announced his discovery of the electron and the fact that atoms must have some structure. Thomson proposed the plum pudding model of the atom. He postulated that the negatively charged electrons were scattered throughout a cloud of positive charge.
In 1909 Ernest Rutherford’s nuclear scattering experiments showed that atoms consist of a positively charged nucleus surrounded by negatively charged electrons. He envisioned the electrons orbiting the nucleus in a matter analogous to the planets orbiting the Sun.
Rutherford’s solar system analogy for atomic structure however had serious problems. Bohr’s model of the atom used principles from quantum mechanics to better explain atomic structure.
Bohr Model of the Atom
In 1913 Niels Bohr proposed his model of the atom. Applying newly discovered principles of quantization to the atom, Bohr was able to solve the problems with Rutherford’s atomic model and to advance quantum mechanics.
Just as Einstein postulated that light is quantized to explain the photoelectric effect, Bohr postulated that electron orbits about the nucleus are quantized.
When electrons orbit the nucleus in more distant orbits, they have more total energy. Bohr called possible electron orbits energy levels or stationary states. In Bohr’s model energy levels are quantized. Only specific discrete energy levels are possible. The lowest energy level is the ground state, and higher energy levels are the first, second, etc. excited states. As long as the electrons are in one of these quantized energy levels or stationary states they orbit the nucleus and remain stable without emitting electromagnetic radiation and losing energy.
Electrons can jump up or down between energy levels, but cannot have energy values between the allowed energy levels. There are no fractional energy levels.
Explaining Spectral Lines
Electrons jumping between energy levels explains spectral lines.
When an electron jumps to a higher energy level, it must get the extra energy from somewhere. One way is by colliding with another atom, leading to collisional excitation. Another way is by absorbing a photon of light. Both photons and energy levels are quantized. When a photon’s energy equals the energy difference between energy levels and it strikes an electron, the electron absorbs the photon and jumps to the higher level. This process causes absorption line spectra.
When a photon is in an excited state, either from collisional excitation or from having previously absorbed a photon, it can jump to a lower level. The electron rids itself of the extra energy by emitting a photon having the right amount of energy. An emission line spectrum results.
Each type of atom has its own unique set of spectral lines because it has its own unique set of energy levels.
Wave Particle Duality and the Bohr Hydrogen Atom
What determines the values of the allowed energy levels? De Broglie’s wave particle duality provided the answer for the Bohr hydrogen atom. Particles, such as electrons, also have wave properties. The possible energy levels for the electron in hydrogen atoms are determined by the electron wavelength. Bohr required an integer multiple of the electron’s wavelength to equal the circumference of the allowed orbits.
This scheme works very well for hydrogen, the simplest atom. One of the limitations of the Bohr model is that it cannot predict the possible energy levels for more complex atoms. More complex quantum mechanical modifications to Bohr’s model came later.
The Bohr model however provides a very good description of atomic structure and represents an important leap forward in our understanding of atoms and quantum mechanics.