Much of what we know about the atom, has been determined by the way matter interacts with light. Therefore, in order for us to gain some insight into the nature of the atom, we need to have a working understanding of light.

For years, there were two competing models of light. One model, put forth by Isaac Newton, stated that light consists of tiny particles, called corpuscles. The other model stated that light was a wave. In 1800, Thomas Young proved that light was a wave by showing that light could diffract and produce a predicted diffraction pattern after passing through a double slit.

Because light is a wave it has a wavelength, l and a frequency, f.

All waves follow the universal wave equation:  speed (c) = l x f

Light is a very small part of the electromagnetic spectrum. Electromagnetic waves are different from sound waves and other waves because no material is required to transmit electromagnetic radiation. Most of the electromagnetic spectrum is invisible to the human eye.

We can look at an overview of electromagnetic spectrum in the diagram below.

The Photoelectric Effect

It seemed that Thomas Young’s diffraction experiment was the last word on the nature of light. Because light can reflect, refract and diffract, it seemed that the particle model was completely disproved. In 1900 an experiment called the photoelectric effect changed our whole understanding of what light was.

The photoelectric effect occurs when light is shined onto a negatively charged metal plate. The prediction before the experiment was that if the light was bright enough, the energy from the light would knock the electrons off of the metal plate. The idea was that light being a wave carries energy just like a water wave. Small waves have little energy but a tidal wave has lots of energy. Dim light would be like a small wave and would not have enough energy to eject the electrons from the metal. Increasing the brightness should increase the energy to the point that eventually there will be enough energy in the light to knock the electron free. While this all seems quite reasonable, the light, no matter how bright, did not eject any electrons.

The next step was to see if the color of light had any effect. White light was broken up into different colors. The lower frequency red and yellow light had no ability to eject electrons no matter how bright. At higher frequencies like blue and violet, the electrons were ejected from the metal. Interestingly enough, even dim blue and violet light seemed to be able to eject electrons from the metal.

It took Albert Einstein nearly 10 years to develop an explanation for what was actually happening in this experiment. The explanation changed our whole way of looking at light. Light does have a particle nature as well as a wave nature. These particles are called photons and have an energy that is directly related to their frequency.

E_photon = hf

E_photon = energy per photon

h = Planck’s constant = 6.63 x 10^-34 J*s

f = frequency

Because f = c/ l , energy can also be calculated in terms of wavelength

E_photon = h c/ l

While this may seem quite abstract, this model helps explain how matter interacts with light.

Absorption and Emission of Light by Matter

When a substance has light passing through it, some energies of light are absorbed and others are allowed to pass through. The result is an absorption spectrum.

When a substance is heated, it will emit only certain frequencies (colors) of light.. This is an emission spectrum.

The presence of absorption and emission lines tell us that atoms can only absorb and emit specific energies of light. Different elements have different absorption and emission spectra.

The explanation for why atoms emit or absorb specific energies is that the electrons are only allowed to be in certain energy levels. This is the main idea of quantum theory. Electrons can only have specific amounts of energy. When electrons absorb the right amount of energy, they can move from a lower to a higher energy level. When they move from a higher to a lower state, they emit a photon with a specific amount of energy.

This information eventually led to a modification of Rutherford’s model of the atom. Neils Bohr proposed a model that looks similar to Rutherford’s model except for the discrete energy levels of the atom.

Third Quantum Number - Orbitals ( m = -l, ....0, ....+l )

How many orbitals are there in a sublevel?

The Fourth Quantum Number - Spin (s)

  Within an orbital how many electrons can there be? 2

How do the electrons fill an orbital?

One with clockwise spin (+1/2)

and one with counterclockwise spin (-1/2).

What does this arrangement look like?

Electrons fill orbitals from low energy to high energy. When a sublevel is filled, the electrons fill the orbital one electron at a time and pair up only when all the orbitals in that sublevel are filled.

Remember, the electrons are actually in an electron cloud. The s cloud is spherical, the p orbital clouds look like dumb bells oriented along the x, y and z axes. The drawings above help us visualize the energy of the electron levels and sublevels.

The electron configuration can help us explain why an element reacts the way it does, why it has a certain oxidation state or several oxidation states.

Orbitals: http://wulff.mit.edu/orbs/

Electron Dot Diagrams

Electron dot diagrams are useful because they allow us to "see" the outer electron configuration of an atom. This helps us guess how the element will react and what the shape of the resulting molecule may be. When we draw the electron dot diagram for an element, we include only the outer electrons. This includes electrons from the s and p sublevels (orbitals) with the highest quantum number.

How do we draw the electron dot diagram for oxygen?

1) Consult the periodic table and determine the atomic number for oxygen.  

   8

2) Draw out the electron configuration of the element.    

1s2 2s2 2p4

3) Identify the sublevels with the largest principle quantum number and draw the orbital configuration for the electrons in these sublevels.     

4) Write the symbol. Place the dots at 12, 3, 6and 9 O'clock, corresponding to the outer s, px, py and pz orbitals.