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. The colours of the visible spectrum are shown below:
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 and are often referred to as radiant energy, whether they are radiowaves, microwaves or gamma waves 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 radiowaves 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 is what has the electromagnetic spectrum
got to with the structure of atoms?
Surprisingly enough quite a lot really! When atoms are excited by
supplying them with energy this can excite
the electrons within the atoms to move to higher
energy level or shells. When these excited
atoms lose this extra
energy when the electron drop back
lower energy shells electromagnetic radiation
is given off at very precise
wavelengths. The type of electromagnetic radiation emitted could for example be
visible light, ultraviolet or even x-rays, it all depends on
which shells the electrons drop down from. The shells or energy levels which
are further from the nucleus contain electrons with lots of
so when these electrons drop back down to lower energy levels
they will release electromagnetic radiation such as x-rays which contain lots
of energy. However if
the electrons only fall from say the second energy level
or shell to the first shell then electromagnetic radiation of lower
will be released.
The wavelengths and frequencies
of the electromagnetic radiation released by atoms can be
seen as a series of lines produced on film. The emitted electromagnetic radiation
is passed through a spectroscope, this separates the various wavelengths
of electromagnetic radiation and an emission spectra
is produced (some examples 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.
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 matter on the atomic scale behaves.
As an example consider the simplest element, hydrogen. Johann Blamer, a Swiss school teacher noticed in 1885 that the emission spectra of hydrogen contained 4 lines in the visible part of the spectrum. These 4 lines with their corresponding wavelengths (λ) are shown below:
The 4 lines in this part of the visible spectrum
are produced when the electrons in an excited state
(high energy level or electron shell) fall back down to the second shell or energy level
The diagram shows 4 possible transitions from the higher electron levels
hydrogen atom down to the second level or second electron shell. We can
calculate the wavelength of the light released for each of the
transitions shown using the Rydberg-Balmer equation.
Here λ represents the wavelength of the visible light released.
R is a constant called the Rydberg constant, R=1.097 x10-2 nm-1 m and n are simply the electron shell numbers or principal quantum number, m is 2 in the case of the Balmer series, which shows up in the visible spectrum and n is the starting electron shell.
So for example the red line in the hydrogen spectrum at 656nm (nanometres) is due to a transition of the electron from electron shell or energy level 3 down to shell 2. If we substitute these values into the Balmer-Rydberg equation we get:
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 line present in the ultraviolet region (called the Lyman series) and the infrared (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 opposite:
So far we have discussed the movement of electrons between
electrons shells inside atoms. When the electrons
are excited and gain energy they jump to higher
energy levels, and then when the electrons lose this additional
energy it is released again as the electron
falls back down to a lower energy shells.
One of the first scientists to consider this idea was a German physicist called Max Planck. He was trying to solve a completely unrelated problem to do with the intensity and wavelength of electromagnetic radiation released when solids are heated to higher and higher temperatures.
Planck proposed the idea that atoms can only absorb and release energy in certain sized chunks or pieces. He called these small packets of energy "quanta". We call these quanta or packets of energy- photons. Prior to this it was assumed that atoms could absorb and emit a continuous spectrum of energy. However Planck showed that this idea was wrong. Furthermore he showed that the energy of the photons emitted or absorbed by atoms depends on the frequency of the electromagnetic radiation. The simple equation below allows you to calculate the energy of these "quanta" or packets of energy simply from the frequency of the radiation:
Example 1. The line spectra of the metal sodium shows a very characteristic double yellow lines at wavelengths of 589 nm and 589.6nm. Calculate the energy of a photon of yellow light with a wavelength of 589nm. Firstly we need to find the frequency for each of the two yellow lines. Simply rearrange the formula we used above to calculate the frequency of the yellow light:
Speed = frequency x wavelength
c= ν x λ
Simply rearrange the formula to calculate the speed and make frequency the subject of the formula, then it is simply a case of substituting in the values:
So according to Planck's theorem a particle absorbing or emitting radiation with a frequency of 5.09 x 1014 Hz cannot gain or lose energy except in multiples of 3.37 x 10-19J. That is the radiation lost or gained is quantised and the particle cannot gain or lose energy of any other value.
The photoelectric effect is a phenomenon that had confused scientists for a long time. Scientists had observed that when a metal surface has electromagnetic radiation (usually visible light) of the "required" frequency shone on it that a stream of electrons is emitted from the surface of the metal. These emitted electrons, being negatively charged can be attracted to a positively charged anode and an electrical current will flow. This is shown as shown in the diagram below.
Basic photocells work using this effect. The glass tube needs to be evacuated since any air present will prevent the electrons from reaching the anode.
For each metal it was observed that there is a minimum frequency
of light needed before any electrons are emitted,
no matter how intense the beam of light is. If light below this minimum frequency
is shone on the metal then no
Albert Einstein explained the photoelectric effect by suggesting that light did not consist of waves but of tiny particles or packets of energy, called photons. He developed Planck's idea of quantised energy by suggesting that the photons had energies dependant on their frequencies (recall that: E=hv). The higher the frequency the higher the energy of the photons. Einstein suggested that if the frequency of the light is below the minimum value then no electrons are emitted. However if the frequency is increased then the energy of the photons is increased and as long as it is above the minimum energy required to overcome the attractive electrostatic attraction between the negatively charged electrons and the positively charged metal ions in the metallic bond then the electrons will be ejected from the surface of the metal.
Neils Bohr used the emission spectrum
of hydrogen and developed the ideas of quantised energy to
develop a new model of the atom. One, which improved on Rutherford's nuclear model of the atom.
Bohr's model of the atom, shown opposite, had the electron orbiting the nucleus in
shells with each shell representing a
energy level and that these shells had definite energy level, that is they are
quantised. Bohr developed the idea that as electrons are
excited and gain energy, they eventually lose
this energy and drop back down to lower energy levels. Bohr using this
idea was able to calculate the radius of the hydrogen atom and the energy values for each of the
shells within the hydrogen atom. He also suggested that there were only certain allowed energy levels
which we call shells or energy levels, this is the idea that the energy levels are
quantised and only
certain values are allowed.
That is to say the electron can be in the first or second shell/energy
level but not somewhere
inbetween. However Bohr's model and calculations only worked for atoms with 1 electron.
He could not explain the spectrum produced by other atoms with
more than one electron.
The next big step in understanding the internal structure of atoms was put forward by Edwin Schrodinger, one of Neils Bohr students. Schrodinger produced a series of equations based on what we now call the quantum mechanical model of the atom. This model rejects the electron as a particle but instead focuses on the wavelike properties of the electron. Waves are normally described using a series of equations, the solutions to these waves equations are called waves functions or orbitals, they are usually represented by the Greek symbol psi, Ψ. These wave functions provide information about the position, energy and orientation of the electrons inside atoms. Ψ2, the wave function squared is a probability density function which as its name suggests provides a probability of where the electrons can be found inside the atoms. It is important to realise that an orbital is not the same as an orbit. In the model of the atom from GCSE, the Bohr atom, the electron is imagined as a particle orbiting the nucleus. However an orbital is a 3d- area in which there is a 95% chance of the electron being there (remember that this new theory the electron is NOT a particle but it behaves as a wave). An orbital is a volume of space around the nucleus where there is an excellent chance or probability of the electron being found. It is not an orbit! The Schrodinger's wave equations produce 3 quantum numbers that offer a description of the location, energy and orientation in space of the electrons inside the atom.
As above we described that when an electron
moves from one energy level inside an atom to a higher energy level then it will absorb
energy (electromagnetic radiation) of a particular wavelength
or frequency; and when this same
back down to its original energy level then it will emit radiation of the same
wavelength and frequency as
An emission spectra is produced by recording the wavelengths of electromagnetic radiation emitted by the electrons inside an atom when they lose energy and move from higher electron energy levels to energy levels with lower energies. When the emission spectra of atoms were first examined in detail, lines which initially appear as single lines were often in fact two lines very close together. This created a bit of a problem since there were not enough energy transitions inside the atoms to account for all these "extra lines". The image below shows a typical emission spectra for an element, in this case sodium.
The explanation of these double lines was provided by considering the property of electron spin. Now quantum mechanics assumes that the properties of the electron are based on its wavelike behaviour, but to simplify things consider for a second that the electron is a particle. Since the electron has a charge, a spinning charged particle will generate a magnetic field with the orientation of this magnetic field determined by the direction of spin. Quantum theory only allows two values for this spin property of the electron, that is it is quantised. The new property of the electron is assigned to a new quantum number, the electron spin quantum number with allowed values of +1/2 and -1/2. The two directions of spin produce magnetic fields which are directly opposed to each other. These two oppositely directed magnetic fields are responsible for the double lines produced in the line spectra of many elements. This is illustrated in the diagram below:
The presence of this magnetic effect due to the property of electron spin can be easily demonstrated if atoms with unpaired electrons are passed through a magnetic field as shown below. If there was no magnetic effect from the electrons in the atoms then they would simply pass through the poles of the magnet and not be affected by it in any way. However this is not what is observed. In a famous experiment called the Stern-Gerlach experiment the atoms of an element are deflected from their path by the external magnet. In this famous experiment which was carried out in 1921 Stern and Gerlach it was demostarted that a beam of neutral atoms could be split into two groups when passed through a magnetic field. You might be surprised by this result since you might imagine that neutral atoms would simply pass straight through the magnetic field and not be deflected in any way. However Stern and Gerlach observed that the magnetic field produced by electron spin property caused the atoms to be deflected into two separate groups. The fact that the atoms split into two separate groups indicate that there are only two allowed values for this spin property of the electron.
This new electron spin quantum number is important in deciding how the electrons are arranged inside the atom. The Pauli Exclusion Principle, proposed by the physcist Wolfgang Pauli, states that no two electrons in an atom can have the same 4 quantum numbers. This is important when we come to decide how the electrons occupy orbitals, which will be considered in the section on the AUFBAU principle.