Electron Arrangement

Electron Arrangement

Atomic Emission Spectra

• When an element is excited (energy is added to it), it will emit light of a characteristic colour. This colour is a mixture of light energy waves that have different frequencies. Since each frequency represents a different colour of light, the colour that the atom emits is a mixture colours. This emitted light be broken into its various component frequencies using a prism. A line spectrum is produced comprised of only those colours of light corresponding to the frequencies that made up the originally emitted light.

• White light, like that from the sun, is comprised of ALL frequencies of visible light. Passing white light through a prism will produce a continuous spectrum - like a rainbow - as opposed to a line spectrum.

[pic 1]

[pic 2]

• When an atom’s electrons absorb energy they may jump to a higher energy level. They then lose the absorbed energy, ultimately falling back to ground state, and emit the energy as light. The colour of light depends on the amount of energy emitted, which determines the frequency of the wave.

Convergence

The energy levels are not evenly spaced like the rungs on a ladder. The higher the energy, the smaller the difference in energy between successive energy levels. This means that at higher energies the lines in a spectrum will converge (get closer together) as frequency increases.

If an electron gains more energy than the convergence limit for that atom, it is no longer attracted by the nucleus and is lost…the atom becomes a cation.

The Hydrogen Spectrum – Lyman, Balmer, Paschen

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The two lines furthest to the left have wavelengths less than 400nm and therefore are considered to be in the ultraviolet spectrum (>10nm – 400nm). The “Balmer” series contains but is not limited to visible spectrum (>400 – 700nm) wavelengths.[pic 4]

The spectrum is divided into a number of distinct series named after the people that discovered them. Each series corresponds to transitions in which the electron falls to a particular energy level. They occur in different spectral regions because as the energy levels increase, they converge at [pic 5]. This means that all transitions to the n=1 level include the emission of a large amount of energy between n=2 and n=1 → a high energy, UV emission.