When white light is passed through a glass prism it is split up into its constituent colours and we observe a continuous spectrum of colours from the long wavelength red light to the much shorter wavelength violet light. Isaac Newton performed this very experiment in 1666. You may have carried out a similar experiment in your science lessons using a ray-box and a glass prism; this is shown opposite.
The spectrum of colours that we see when white light is split up into its constituent colours is just a much smaller part of the electromagnetic spectrum. Visible light consists of a range of wavelengths of electromagnetic radiation that we can detect with our eyes these colours are red, orange, yellow, green, blue, indigo and violet (ROYGBIV). The colours of the visible spectrum are shown in the diagram below.
The continuous spectrum of visible light ranges from the low energy long wavelength red light to the higher energy shorter wavelength violet light. Perhaps the most obvious example of the continuous spectrum is the spread of colours in a rainbow.
However the visible spectrum is only a small part of the electromagnetic spectrum which is shown below. You will no doubt remember this from your gcse science lessons!
All electromagnetic waves carry energy that is often referred to as radiant energy; whether they are radio waves, microwaves or gamma waves they all have certain common properties. They all travel at 300 million metres per second (3 x 108 m/s) in a vacuum. As shown in the diagram above they have wavelengths that cover a vast range; from kilometres in the case of radio waves down to 10-13m for gamma rays. Since all electromagnetic wave move at the same speed the frequency and wavelength are both related through the formula:
Speed = frequency x wavelength
c= ν x λ
Perhaps a more obvious question we could ask is what has the electromagnetic spectrum
got to with the structure of atoms?
Surprisingly enough quite a lot really! By analysing the light emitted when a high voltage is passed through gases under pressure it is possible to deduce information about the internal structure of atoms. When atoms are excited by
supplying them with energy of the correct frequency this can excite
the electrons within the atoms to jump to higher
energy level or shells. When these excited
atoms lose this extra
energy the electron drop back
down into
lower energy levels or shells and electromagnetic radiation
is given off at very precise
wavelengths and frequencies.
The type of electromagnetic radiation emitted when an electron returns back down to a lower energy level could be
visible light, ultraviolet or even x-rays as the frequency of the electromagnetic waves emitted depends on
which shells the electron drop down to. The shells or energy levels which
are further from the nucleus contain electrons with lots of
energy so when an electron drop back down to lower energy levels
it will release high energy electromagnetic waves such as x-rays. However if
the electron only fall from say the third energy level
to the second energy level then electromagnetic waves of lower
energy
will be released this is summarised in the diagram below:
When the electron absorbs energy of the correct frequency it becomes excited and jumps to a higher energy level or shell inside the atom. In the diagram below the electron jumps from the ground state or first energy level up to the third energy level. In order to make this jump from the first energy level to the third energy level the supplied energy must match exactly the differences in energy between the two energy levels.
When the electron loses the energy it gained it will drop back down to the first energy level or ground state. When the electron returns back to the ground state the energy it gained is emitted as electromagnetic waves or energy.
This idea that electrons can absorb energy
when they are excited by heating or applying a voltage across
atoms and jump to a higher electron shell or
energy level then simply release this same "packet" of energy when
it drops back down to a lower energy electron shell may seem fairly
simple and obvious to you but it is in fact
one of the foundations of quantum theory; a theory which completely revolutionised our understanding of how the particles found inside atoms actually behave.
The wavelengths and frequencies
of the electromagnetic waves released by electrons when they drop down energy levels can be
seen as a series of coloured lines produced on photographic film. The emitted electromagnetic waves
are passed through a device called a spectroscope which can separate the various wavelengths
of electromagnetic waves and an emission spectrum
is produced; some examples of these are shown below. By making some quite simple calculations from these
emission spectra it is possible to gain an insight into the
internal structure of atoms.
So for example the red line in the hydrogen spectrum at 656nm (nanometres) is due to the electron losing energy and dropping down from the third energy level 3 to energy level 2. If we substitute these values into the Balmer-Rydberg equation we should be able to show that the red light emitted does indeed have a wavelength of 656nm; the calculation to show how to calculate the wavelength of the blue light emitted is also shown below:
You can check the other values for the wavelengths produced in the spectrum by simply substituting the other values for the electron transitions into the Balmer-Rydberg equation The line spectra for hydrogen also has other lines present in the ultraviolet region (called the Lyman series) and the infrared region (called the Paschen series); both these series were named after the people who discovered them. Ultraviolet radiation being of shorter wavelength and higher frequency than either visible or infrared has a much higher energy than both of these electromagnetic waves. The transitions for each of these series are shown in the diagram below:
You can of course use the Rydberg equation to calculate the wavelength of all the electromagnetic waves in each of the series in the image above. When Ernest Rutherford suggested a new nuclear model of the atom following his famous gold foil scattering experiment where the internal structure of the atom was often described as a "mini solar system" with the nucleus at the centre of the atom and the electrons like planets were imagined to orbit the nucleus. However this model presented several problems; one of which was the fact that some physicist argued that a negatively charged electron orbiting around a positively charged nucleus should emit electromagnetic radiation as it loses energy and it should simply crash into the nucleus.
The Nobel Prize winning physicist Niels Bohr however proposed a new theory; he suggested that the electron would remain in orbit around the nuclei in certain predefined orbits; he called these predefined orbits "allowed energy states" and an electron in these allowed energy states or levels would not lose energy and crash into the nucleus. Bohr was able to calculate the energy of these allowed energy levels and also the differences in energy between one level and the next. Bohr used his new model to explain the line spectrum of the hydrogen atom. He also suggested that an electron could ONLY move from one allowed energy level to another by absorbing or emitting energy only of the correct frequency or energy. The electron was permitted to move from one energy level to another but they were not allowed to move in between these allowed energy levels. This means that electrons can only absorb or emit electromagnetic radiation with energies which are the difference between two energy levels. That is Bohr had quantised energy. That is the electron cannot absorb any amount of energy, it can only absorb or emit energy in small packets which correspond exactly to the differences in the energy levels within atoms.