Caption: An ABSTRACT diagram of a hydrogen atom changing energy levels and emitting a photon. In general, a change in energy levels is called a atomic transition.

Features:

1. Atoms don't really look like this.

   Atoms can be imaged directly in various ways. But they do NOT have sharp edges, and so they look like fuzzy balls. In general, a real image of an atom does NOT give a clear picture of its structure in any direct way.

2. The ABSTRACT diagram derives from the Bohr atom which was a model of the hydrogenic atom (i.e., a hydrogen-like atom) introduced in 1913 by Niels Bohr (1885--1962). The Bohr atom is historically important in the development of quantum mechanics, but it is a wrong model.

   But it is right in that there are NOT a continuum of energy states allowed for the electrons that surround the atomic nucleus which for hydrogen (the simplest of all atomic nuclei) is a single proton represented by the central dot in the diagram. A hydrogen atom also has only one electron represented by a small dot in the diagram.

   Instead of a continuum of energy states, there is only a discrete set with discrete energies. This is main reason why we call quantum mechanics quantum mechanics: the energy states are QUANTIZED.

3. For atoms, we call the energy states energy levels.

   The quantized energy levels in the ABSTRACT diagram are represented by a quantized set of circular orbits.

   The larger the circular orbit, the higher the energy of the energy level.

   The smallest circle represents the ground state, the lowest energy energy level allowed by quantum mechanics.

   The Bohr atom did posit actual circular orbits, but that turned out to be WRONG.

4. To emphasize, the circular orbits are NOT the real appearance of the energy levels.

   In fact, each energy level is a spread out density distribution for an electron. The electron exists in a continuum superposition of positions with the amount of it at any point being determined by the density distribution.

   The electron is usually only one energy level at at time---but it can be in a superposition of energy levels.

5. When the electron changes energy levels (i.e., makes an atomic transition), energy is emitted or absorbed.

6. A common way for this to happen is by the emission or absorption of a photon: the quantum particle of electromagnetic radiation.

   The energy of the emitted or absorbed photon is determined by the conservation of energy.

   An absorption process requires an incident photon.

7. An emission can be a spontaneous emission or can be stimulated emission caused by an incident photon.

8. There is NO spontaneous emission NOR stimulated emission from the lowest energy level the ground state.

   In the ground state, the atom simply has no REMOVABLE energy.

   It does have an IRREMOVABLE zero-point energy dictated by quantum mechanics.

9. Since the energy levels are quantized, only certain energies are allowed for the emitted or absorbed photons.

   This means only certain frequencies and wavelengths are allowed for the emitted or absorbed photons.

10. If ΔE is the change in energy of the atom for an emission and also thereby the energy of the emitted photon, the frequency ν of the emitted photon is determined by the de Broglie relation ΔE=hν, where the h is the Planck constant h = 6.626070040(81)*10**(-34) J-s which is a fundamental constant of nature.

    Wavelength is given by de Broglie relation ΔE=hc/λ, where λ is wavelength.

11. Every atom and molecule has a unique set of quantized energy levels, and so every atom and molecule has a unique discrete spectrum of frequencies/wavelengths of emission/absorption. These spectra are called line spectra. The study/analysis of line spectra is called spectroscopy.