• Spectroscopy
• The Colors of Light
• The Bohr model of the atom
• Quantum Mechanics
  • Determinism
  • The Uncertainty Principle
  • Quantum numbers in atoms
  • Energy Levels in atoms

The following had been known during the 19th century:

• accelerated charges produce light
• and hence emit energy

If we picture an electron as in orbit around the nucleus, it should radiate light

• changing direction is acceleration! (a force is required from something to change direction)

If the electron radiated due to its motion around the nucleus, it would lose energy and soon spiral into the nucleus.

The world should collapse instantly!

• Fortunately it doesn't.

So what is wrong with this picture?

Enter Quantum Mechanics (QM)

Quantum Mechanics was developed during the revolution which occurred in physics from 1900 - 1930.

Both Special Relativity and General Relativity were developed during this time.

• The discrete (quantum) nature of the energy "levels" of the electron gives QM its name.
• QM describes the microscopic world.
  • Physics of the small
• There are some very non-intuitive things associated with it.

• Classical physics is deterministic.
• That is, a given cause always leads to the same result.
• Even chaotic behavior is deterministic.
• However, in quantum mechanics this is not the case!

• The uncertainty principle of QM states that we cannot know both where something is and how fast it is moving.
• Thus we cannot predict exactly what will happen in a given experiment.
• We can only give the probability of an outcome.
• The more accurately you measure the position the less accurately you know the velocity and vice versa.

• Atomic particles (electrons, protons, etc.) sometimes behave like particles and sometimes like waves.
• Photons can do this too!

Example:

An example of particles behaving like waves is interference. You have probably seen the wakes from two different boats "interfere" on the water. In some cases they enhance one another (constructive interference) while at other times they cancel one another (destructive interference). Photons do this but so can particles!

The microscopic world of atoms and electrons

leads us to ...

the super-macroscopic world of stars and galaxies.

An electron can reside in one of many orbits.

• n = number of orbits = Principal Q.N.
• n = 1 is the lowest energy state.
• If n > 1, then the atom is "excited."

For convenience, instead of drawing circular orbits in which the energy is high for each successive orbit, we draw an "Energy Level Diagram" as shown on the right below. The energy levels of an atom are now represented by horizontal lines. Energy increases as you move upward in the diagram.

An electronic transition occurs when a electron moves between two orbits.

When absorption of a photon occurs, an electron goes up, e.g. from n=1 to n=2.

Emission of a photon occurs when an electron moves down, e.g. from n=2 to n=1.

An analogy to the energy levels in the atom is the "Quantum Stepladder" where the rungs on the ladder correspond to energy levels in the atom.

A spectrum is the intensity of light seen from an object at different wavelengths.

• e.g. done by spectroscopy with a prism.
  

Individual atoms, like H, show spectral lines, i.e. For H, these are the Balmer lines in the visible.

For a "conversational" discussion on spectral lines can be found at the following link.