A model of the atom and the electromagnetic spectrum

Vibrating charges produce radiation. This is the source of electromagnetic radiation, e.g. visible light, X-rays and gamma rays.

At the end of the 19th century, physicists knew that atoms contained electrons and the motion of these electrons resulted in electromagnetic radiation. However, a mystery remained. Solving this mystery, and the work of scientists including Fraunhofer, Brewster, Kirchoff, Bunsen and Ǻngstrom, took forward a new area of science – spectroscopy – which revolutionised solar physics and astronomy.

Prior to Isaac Newton’s prism experiment in 1666, the nature of celestial bodies such as the Sun and stars was not well understood. The luminescence from such bodies was attributed to an ‘essence’.

In 1802 William Wallaston observed lines in the spectrum from the Sun. The systematic study of these lines was undertaken by a German optician, Joseph Fraunhofer.
 

© University Corporation for Atmospheric Research

This is a reproduction of Fraunhofer's original 1817 drawing of the solar spectrum. The more prominent dark lines are labelled alphabetically; some of this nomenclature has survived to this day

 

In the 1850s Kirchoff (a German physicist, 1824–1887) and Bunsen (a German chemist, 1811–1899) took another step forward in spectroscopy with the discovery that when various substances were held in the flame of a Bunsen burner and the light produced passed through a prism, distinct spectra were seen, which correspond to distinct elements in the substance. You may have carried out similar experiments in science at school, for example burning a copper salt in a Bunsen flame and observing the distinctive green colour, or sodium (yellow flame), or potassium (lilac flame), or calcium (red flame).

In the 1860s this technique was used by Kirchoff and Bunsen to discover two new elements, caesium and rubidium. Other elements that were discovered using this approach were gallium, helium, argon, neon, krypton and xenon.

We can use a spectroscope to observe various light sources around you, e.g. daylight (do not point the spectroscope directly at the Sun) and the artificial lighting within the classroom If you do not have a spectroscope in school, it is possible to build a simple one that works just as effectively (http://www.exploratorium.edu/spectroscope/).

A spectrophotometer allows you to observe and record data of the spectrum for a range of light sources. You can compare the spectrum of an ordinary filament lamp with time from its first switch-on. Compare this with the spectrum of an ordinary filament lamp with differing voltage across the lamp. Also contrast these with the spectrum of an energy-saving lamp with time.

Max Planck and the quantisation of energy

In 1900 the German physicist Max Planck (1858–1947; Nobel Prize winner 1918) proposed the quantisation of the energy of a vibrating molecule, i.e. it could only take on certain fixed values.

This helped to explain the gaps and lines in the spectra of elements.

 

 He proposed that:

E µ f

with a constant of proportionality 6.626 × 10–34 J s, this came to be known as Planck’s constant (h).

 

E = hf

E = energy in joules (J) f = the frequency in hertz (Hz) h = Planck’s constant = 6.626 × 10–34 J s

This was a revolutionary idea which provided the basis for quantum physics. Einstein built on Planck’s work, proposing that light was also quantised in photons.

 

The Bohr model of the atom

 

In 1911–1913 the Danish physicist Neils Bohr (1885–1962; Nobel Prize winner 1922) collaborated in Cambridge and Manchester with JJ Thomson and Ernst Rutherford.

Bohr knew that a charged body in motion must emit energy as electromagnetic radiation.

This meant that orbiting electrons should emit energy, as a result lose energy and spiral into the nucleus of the atom. This would suggest that the atom cannot be stable.

Bohr struggled to reconcile the accepted model of the atom with this conclusion.
In 1913, Bohr proposed a new model of the atom, which was recognised with the award of the Nobel prize for physics in 1922, and (with some further improvement to take account of Heisenberg’s work in the 1920s), remains the model that accounts for the physical and chemical properties of elements today.

In the Bohr model:

• Electrons exist only in allowed orbits. An electron in an allowed orbit does not radiate energy, i.e. Bohr’s model proposed that the classical electromagnetic theory was incorrect. Electrons can jump between allowed states, and this jump is associated with absorption or emission of electromagnetic radiation.
• Electrons in different orbits have different energies, thus these orbits can be considered energy levels.
• Electrons tend to occupy the lowest available energy level, i.e. the level closest to the nucleus.

The allowed orbits of electrons can be represented in an energy level diagram. The diagram below represents the energy levels within a hydrogen atom.

Electrons can exist either in the ground state E0 or in various excited states E1–E5.

    

An electron in an excited state can make a transition into a lower energy state.

In a hydrogen atom, the possible transitions are:

E1 → E0, E2 → E0, E3 → E0, E4 → E0, E5 → E0
E2 → E1, E3 → E1, E4 → E1, E5 → E1
E3 → E2, E4 → E2, E5 → E2
E4 → E3, E5 → E3
E5 → E4

When an electron makes a transition from a higher energy state to a lower energy state where does the excess energy go? It is emitted as a photon of electromagnetic radiation. In the Bohr model, it follows that the possible excess energies must be quantised, i.e. in specific ‘packets’, since the electron can make transitions only between specified energy levels.

The frequency of the emitted radiation is proportional to the difference in the energy levels (ΔE):

ΔE µ f

It follows that since ΔE is quantised, the photons emitted can only be of specified frequency.

 

The Bohr model and spectra

This theory applies only to free atoms. A free atom is one that is not affected by neighbouring atoms, as in a gas. In sources such as filament lamps, the atoms are not free and their electrons may be shared in bonding with other atoms. This results in an infinite number of possible transitions, giving an infinite number of lines, i.e. a continuous spectrum.