History of the Atom

History of the Atom

“Properties are a function of structure”. In particular the properties of matter are determined by the architecture of atoms. So we devise various “models” or pictures of atoms to help us better understand how nature works. All of these models are human inventions and all are useful in certain circumstances. The true nature of the atom may never be completely understood. The word atom comes from the Greek word “atomos” which means indivisible and is attributed to Democritus 2500 years ago. Today, using a series of computer animation we are going to examine the experiments and reasoning that scientists used to develop the modern atomic theory. We will be examining the following five atomic models.

Billiard Ball Model – Average Kinetic Energy

At the relatively low energies found on the surface of planet Earth, we can envisionatoms as being just like tiny billiard balls. In a gas or liquid they fly around and bounce into each other. Collisions are perfectly elastic – e.g. no loss of energy. In a solid they just vibrate. In either case the temperature of the matter is directly proportional to the average kinetic energy of the atoms. In fact, this is the physical definition of temperature.

Move the temperature bar back and forth. Observe what happens to the velocity of the helium atoms as you change the temperature. Describe how velocity changes with temperature.

Now click the bar that says “enable tracking”. This will enable you to track a single atom. On the graph at the right the red tick marks the speed of this one atom. Now watch this single helium atom. The line graph on the right represents velocity. The single black line represents average velocity. Notice how the single atom speeds up and slows down but on average stays near the average velocity. Keep this window open.

Einstein’s ‘Experimental Proof’ of the Billiard Ball Model

Even though the Billiard Ball Model accurately predicted the Kinetic Theory demonstrated by the above animation, many scientists did not believe that atoms existed because there was no observable experimental evidence. In 1905 Albert Einstein suggested an experiment that would provide observable scientific evidence based on Brownian Motion. Click on animation below. The random movement of the tiny drops can only be explained if there are small objects (atoms) bombarding the droplets. Einstein predicted the expected random drift and this was later corroborated by experiments. (Einstein was a theoretician. But Einstein's predictions were finally confirmed in a series of experiments carried out by Chaidesaigues in 1908 and Perrin in 1909. In fact, since the above animation is on based Kinetic Theory equations, we can actually prove this random drift experiment in cyber space using the previous animation. Introduce a neon atom and notice how this moves similar to the particles in Brownian motion animation.

Thompson Model (Plum Pudding Model)

In 1898 John Thompson discovered negatively charged particles which he called electrons. He reasoned that an atom has a neutral charge and he reasoned that there must be positive charges to neutralize the negatively charged particles. He imagined an atom made up of a distributed positive charge embedded with electrons. This was called the plum pudding model.

Rutherford’s Planetary Model

In 1899 Ernest Rutherford and two associates performed a series of experiments that convinced them that Thompson’s modelcould not be correct. They fired Alpha particles (high energy helium nuclei) at a thin gold foil and observed how much they were deflected. Click below to see the experiment.

Did Rutherford expect large deflections? ___ Why not?______. When he did observe large deflection, what did that tell him about the positive charges?______

______

Limitations of Rutherford’s Planetary Model:

Rutherford reasoned that if there was a dense positive ‘nucleus’ surrounded by much less dense negatively charged electrons, that the electrons must orbit the nucleus, otherwise they would just collide. The planetary model proposed that the electrons orbited the positively charged nucleus under the influence of an electric force. (Positive charges attract negative charges.) This is analogous to planets orbiting the Sun under the influence of a gravitational force.

Demonstration: Foam cut revolving around student with marks on the string.

It was understood that the further the electron was from the nucleus the more energy it would have. (This is also true of planetary bodies orbiting other bodies; the Sun and the Earth and Earth and the Moon.) It takes energy to move an electron into a larger orbit further from the nucleus. This energy can come from radiant energy – light for example. Also when it moves into a lower orbit it would release radiant energy. If it could occupy any orbit and the radiant energy could be at any wavelength and the energy produced would be a continuous spectrum as shown below in graph a. below. That is not what is observed.

Graph b shows what is observed when we heat hydrogen gas to several thousand degrees Celsius.

Instead of light at all wavelengths we see only certain wavelengths. A more serious problem with the planetary model is that the electron can be expected to continuously radiate energy and as they give up energy they would rapidly spiral into the nucleus. Clearly this does not happen. Atoms are quite stable.

The Bohr Model

To overcome these limitations, Niels Bohr proposed the following modifications to Rutherford’s model. Bohr suggested that only certain orbits are allowed and there was a minimum energy level below which the electron could not descend below. The theory that explains this is called quantum physics since the energy of the electrons can only exist at specific levels or quanta.

Some things are continuous and some things are quantized.

At a Thanksgiving day dinner the sweet potatoes are quantized, the gravy is continuous. When an electron moved from one orbit to another it absorbs or radiates energy only at those levels corresponding to the difference in energy between orbits.

Clink on the link below.

www.mhhe.com/physsci/astronomy/applets/Bohr/applet_files/Bohr.html

Have the electron jump from the higher orbit to the lower one. Then have it jump back to the higher orbit. When it jumps to a lower orbit it emits a photon at the energy correspoinding to the difference in energy between the two orbits. When it jumps from the lower orbit to the higher orbit it absorbs a photon at exactly the same wavelengh. This is called a dark spectrum.

Using the default orbits, have it jump from 4 to 2 and 3 to 2. These create the two lines in the figure above – called the Balmer series for the hydrogen atom. These lines are in the visible spectra and were the first to be discovered. There are other series jumping to the first orbit and to the third orbit but these are in the ultraviolet spectrum and infrared spectrum and were not discovered until much later. Every substance has a unique spectrum and the science of spectroscopy can be used in forensics to indentfy minute quantities of substances. And scientists discovered helium in the dark spectrum of sunlight years before they had discovered the element on earth – hence its name which comes from helios which is Greek for the Sun.

Adjust the orbits to the setting to correspond to the table below.

For each of the settings in table below, indentify the metal using the color chart.

Metal
Set orbit 3 to 3.34 Jump between 3 and 4
Set orbit 3 to 3.97 Jump between 3 and 4
Set orbit 3 to 4.15 Jump between 3 and 4
Set orbit 3 to 4.24 Jump between 3 and 4
Set orbit 3 to 4.33 Jump between 3 and 4
Set orbit 3 to 4.42 Jump between 3 and 4
Set orbit 4 to 8.02 ; set orbit 2 to 2.71 Jump between 2 and 4

The colors in fireworks come from these and other metals. See if you can identify them in the fireworks bellow. List three metals you find in the fireworks display.

www.maylin.net/Fireworks.html

Color of fireworks / Metal that corresponds to that color

Schrödinger Model

The Bohr model predicts the hydrogen spectrum with great accuracy. In fact it was this excellent agreement between the prediction and the observed results that caused scientists to accept the Bohr model. However, the Bohr model was not able to explain the spectrum of atoms with multiple electrons. To accomplish this, physicists needed a model based on Schrödinger’s wave mechanics -- which is an extension of Bohr’s quantum physics. So first we have to imagine that the electron behaves like a wave. Don’t feel bad if you find this strange. Even physicists find quantum physics strange. However, it predicts the behavior of matter at the atomic level with great precision. Click on the wave model it the animation below.

Here is what the orbitals look like. There are exactly n2 orbitals for each energy level and each orbital can contain only two electrons

Analysis and Conclusions: (Answer the following questions.)

1. Match the model to the following descriptions

Atoms are made of hard indivisible balls______

The positive charge is uniformly distributed within the atom______

The electrons orbit the positively charges nucleus at arbitrary distances______

The electrons can only occupy certain fixed orbits and photons are absorbed or

emitted at energies corresponding to the differences in energy between orbits

______

The exact position of the electron cannot be predicted only the probability of

finding at some position in space.

1. Explain how Rutherford’s experiment led him to believe that positive charges where at concentrated at fixed locations rather than distributed throughout the atom.

1. What were the problems with Rutherford’s planetary model?

1. What were the strengths of the Bohr model? What were the limitations?