The familiar dispersion of white light is illustrated below:
Solids, liquids and dense gases glow at high temperatures. The emitted light, examined using a spectroscope, consists of a continuous band of colours as in a rainbow. A continuous spectrum is observed. This is typical of matter in which the atoms are packed closely together. Gases at low pressure behave quite differently.
The spectral lines in the visible region of the atomic emission spectrum of barium are shown below.
It was necessary to explain how electrons are situated in atoms and why atoms are stable. Much of the following discussion refers to hydrogen atoms as these contain only one proton and one electron making them convenient to study.
In 1913, it was Neils Bohr who solved many of the problems at the time by proposing that the electron revolves around the nucleus of the atom with a definite fixed energy in a fixed path, without emitting or absorbing energy. The electron in the hydrogen atom exists only in certain definite energy levels. These energy levels are called Principal Quantum Levels, denoted by the Principal Quantum Number, n. Principal Quantum Level n = 1 is closest to the nucleus of the atom and of lowest energy. When the electron occupies the energy level of lowest energy the atom is said to be in its ground state. An atom can have only one ground state. If the electron occupies one of the higher energy levels then the atom is in an excited state. An atom has many excited states.
When a gaseous hydrogen atom in its ground state is excited by an input of energy, its electron is 'promoted' from the lowest energy level to one of higher energy. The atom does not remain excited but re-emits energy as electromagnetic radiation. This is as a result of an electron 'falling' from a higher energy level to one of lower energy. This electron transition results in the release of a photon from the atom of an amount of energy (E = hn) equal to the difference in energy of the electronic energy levels involved in the transition. In a sample of gaseous hydrogen where there are many trillions of atoms all of the possible electron transitions from higher to lower energy levels will take place many times. A prism can now be used to separate the emitted electromagnetic radiation into its component frequencies (wavelengths or energies). These are then represented as spectral lines along an increasing frequency scale to form an atomic emission spectrum.
Principal Quantum Levels (n)
for the hydrogen atom.
Comment:
A hydrogen atom in its Ground State.
The electron occupies the lowest possible energy level which in the case of hydrogen is the Principal Quantum Level n = 1.
The Bohr theory was a marvellous success in explaining the spectrum of the hydrogen atom. His calculated wavelengths agreed perfectly with the experimentally measured wavelengths of the spectral lines. Bohr knew that he was on to something; matching theory with experimental data is successful science. More recent theories about the electronic structure of atoms have refined these ideas, but Bohr's 'model' is still very helpful to us.
For clarity, it is normal to consider electron transitions from higher energy levels to the same Principal Quantum Level. The diagram below illustrates the formation of a series of spectral lines in the visible region of the spectrum of electromagnetic radiation for hydrogen, called the Balmer Series.
As referred to above for hydrogen atoms, electron transitions form higher energy levels all to the n = 2 level produce a series of lines in the visible region of the electromagnetic spectrum, called the Balmer Series. The series of lines in the ultra-violet region, called the Lyman Series, are due to electron transitions from higher energy levels all to the n = 1 level, and these were discovered after Bohr predicted their existence.
Within each series, the spectral lines get closer together with increasing frequency. This suggests that the electronic energy levels get closer the more distant they become from the nucleus of the atom.
No two elements have the same atomic emission spectrum; the atomic emission spectrum of an element is like a fingerprint.
The diagram to the right illustrates the formation of three series of spectral lines in the atomic emission spectrum of hydrogen.
