…the study of how light and matter interact Spectroscopy …the study of how light and matter interact

Wave model Light is a form of electromagnetic radiation and behaves as a wave. It has a characteristic wavelength (λ) and frequency (ν). Different colours of light have different wavelengths Wave speed, frequency and wavelength can be related by the equation: c = λ v

Particle Model In 1905 Albert Einstein proposed that light was a stream of tiny ‘packets’ of energy called photons. Different coloured light had different amounts of energy in their photons.

Linking the models… E = h v The wave and photon models can be linked using the following equation: E = h v Where E = the energy of a photon h = planck constant v = frequency of the light

Atomic Spectra Atoms become excited by absorbing energy. This energy is then released as they return to their ground state. n=1 n=4 n=5 n=6 n=3 n=2 Energy

Emission Spectra Spectrum consists of a series of coloured lines on a black background. The atoms only emit at precise frequencies and the prism splits the light to show these frequencies as lines of colour within the visible spectrum.

Complete Spectrum of Visible Light

Absorption Spectra Spectrum consists of a series of black lines on a background of the whole visible spectrum. The atoms only absorb at specific frequencies and the prism splits the light to show these frequencies as lines of black as that particular colour has been absorbed. The black lines in the absorption spectra correspond exactly with the coloured lines in the emission spectra.

Atomic Spectroscopy The sequence of lines in an atomic spectrum are characteristic to the atoms of that element. It is like a ‘chemical finger print’ and can be used to identify the element in mixtures of atoms or even when it is part of a compound. The intensities of the lines provide a measure of the abundance of that element. This is how cosmologists determine the composition of stars vast distances away from the Earth.

Hydrogen Spectra

More Spectra…

Bohr’s Explanation Bohr thought that the lines in the spectra were caused by electrons moving between energy levels, or shells. Excited electrons absorb energy to jump to higher energy levels then release this energy again when they drop back down to a lower energy level. Bohr’s explanation was controversial at the time because it relied upon the quantisation of energy.

Key Points in Bohr’s Theory The electron in the H atom is allowed to exist only in defined energy levels A photon of light is emitted or absorbed when the electron moves from one energy level to another. The energy of the photon is equal to the difference between the two energy levels (∆E) The frequency of the emitted or absorbed light is related to ∆E by : ∆E=hv

Lyman Series of the H atom The Lyman series is the series of lines originating from electrons returning to the ground state level, level 1. The Balmer series originates from electrons returning to level 2 from levels 3,4,5…

Ionisation Energy The energy levels become closer together until they converge. At this point the electron is lost from the atom. The energy difference between this point and the ground state is known as the ‘ Ionisation Energy’. It can be represented by: X(g)  X+(g) + e- n=1 n=4 n=5 n=6 n=3 n=2 Energy