Day 18

Context

We discussed the raisin pudding model and set up the gold foil experiment. We now look at what happened when they performed the gold foil experiment and see how it forced us to change the model of an atom. We then briefly discuss the planetary model and move on to the current model of the atom.

Explanation

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The account of the progression of models for the atom is a good example of how things happen in science. We create models and use them until new data forces us to tweak them or discard them completely for a better model. So, today we go from the raisin pudding model to the planetary model and then to the quantum mechanical model. Quantum mechanics gives us the current model for the atom.

After this class you should be able to:

Gold Foil Experiment

In the gold foil experiment Rutherford, who was the project manager, was assuming that the atoms were as described in the raisin pudding model. With that assumption he predicted that the alpha-particles, which are much larger than the electrons, would go straight though the atoms and hit the detector directly across from the "gun" or source of alpha-particles. The electrons aren't big enough to provide a surface for the alpha-particles to bounce off of and there is not any other object in the atom, so the alpha-particles should go straight through and hit the detector on the other side.

The only problem with Rutherford's prediction is that, while most of the alpha-particles went straight through, some of the alpha-particles were significantly deflected and some of them bounced back toward the source! The raisin pudding model has no way of explaining this. The raisin pudding model no longer agreed with observations and so it needed to be modified or discarded. There is no way to modify the raisin pudding model to account for these observations and so it had to be discarded.

The observations led them to believe that there must be a hard, dense part of the atom that the alpha-particles could bounce off of. They knew that the electrons were much too small and so they decided that the positive part of the atom must be hard and dense. They envisioned a large positive "nucleus" with the electrons outside of this nucleus. A physical model that they were familiar with that has a large object at the center and smaller objects going around it is the solar system. This is where the solar system model of the atom comes from. They thought that the nucleus was like the sun and that the electrons would orbit around the nucleus like planets around the sun. It was about 1911 when this theory was proposed. I have collected a few pages from http://particleadventure.org that describes this and has some pictures to help see what happened. You may want to view those pages.

A few years later the neutron was discovered and we include them in the nucleus with the protons. You are probably familiar with this solar system model of the atom, but it only lasted about 12 years. New experiments that couldn't be explained with the solar model quickly came along and this model also had to be discarded.

Review

Before we move on, lets summarize what we have learned.

Elements are the basic building blocks of all matter. Everything on earth is made from these elements. Atoms are the smallest piece of an element that retains the characteristics of that element. Atoms are composed of protons, neutrons and electrons. The protons and neutrons are almost 2000 times larger than the electrons and reside in the massive nucleus. The electrons are outside of the nucleus, but do NOT orbit around the nucleus like planets around the sun. If the nucleus of a hydrogen atom was the size of the earth, the electron could likely be half way to the sun. There is a lot of space in the atom.

Each element has a distinctive kind of atom. The number of protons in the nucleus determines which element it is. The periodic table allows one to easily determine how many protons are in the atoms of each element by looking at the number that runs consecutively across the table. That number not only numbers the elements, but also designates how many protons there are in the atom of that element. Neutral atoms have the same number of electrons as protons, since the electron has a charge of -1 and the proton has a charge of +1. Note that the number of electrons in an atom can change and that the number of electrons does not identify the element (unless the atom is neutral).

The atomic theory described above has been developed over many years. At one point in history the notion of matter being composed of discrete particles was a radical idea. But more and more evidence came forward to substantiate this view. One theory that was popular for a while was the raisin pudding model. According to this model the atom was composed of a spread-out positively charged cloud with electrons scattered throughout, like raisins (electrons) in pudding (positive cloud). This theory had to be discarded when Rutherford found that alpha particles were scattered by gold atoms. The original Bohr model, which is a planetary model, was popular for a few years (not more than 14 years), but it too was shown to be wrong through experimentation.

Light

Around the same time that Rutherford was doing the Gold Foil experiment another great scientist, Neils Bohr, was conducting experiments with atoms that looked at the different colors of light emitted by the atoms. An interesting aspect of his observations was that each element had a unique set of colors associated with it. This is used today to figure out what elements are in stars and other celestial bodies that we can't get direct samples from.

The important thing about light in this context is that each shade of light has a different energy. So, each of the colors that Bohr saw represented different energies. In other words each element has a specific set of energies associated with it.

Bohr and Discrete Energy Levels

Bohr knew that if anything moved in the atom it would be the electrons, so he attributed the unique set of colors that he observed to the energy of the electrons changing as they moved from one energy level to another one. Going from a higher energy to a lower energy level would release energy which could be in the form of light and that would be consistent with his observations.

Also, it would be necessary to say that the energy levels were discrete and fixed for each element because only certain colors (energies) were observed and not a continuous range of colors is given off. Since the solar system model was the current model at that time Bohr was trying to explain his observations in the context of electrons orbiting around the nucleus like planets around the sun. To get discrete energy levels he said that each orbit corresponded to a specific energy and that there were only certain orbits available in an atom. The available orbits would be unique for each element and would result in a unique set of colors being given off by each element. This would allow him to explain his observations about light by saying that the electrons were moving from one discrete orbit to another discrete orbit and giving off light in the process.

His prediction that electrons reside in discrete energy states and that the electrons can move between these energy states is a major contribution to our understanding of atomic theory and is part of the current model. Electrons can move to higher energy states when excited. They can be excited by light, heat, electrical energy, etc. Electrons, like everything else, want to be in the lowest energy state. So, after they are excited by the light or other energy, they "fall" back down to the lower energy state. Conservation of energy requires that energy be given off when going from a high energy to a low energy. This energy is often given off as light. This accounts for Bohr's observations. Notice, however, that it isn't required that the electrons go around in orbits, they just need to have different discrete energy states available to them. Another way to say that atoms have discrete energy levels is to say that they are quantized. To be quantized means to have discrete levels.

To bolster his theory, Bohr used the solar system model to mathematically model a hydrogen atom. He used hydrogen because it only has one proton and one electron in the neutral atom. His mathematical model perfectly matched his observations and it predicted other energies of light that were later found by experiment. On a roll, he tried to do the equations using the solar system model for lithium which has two protons and two electrons. Unfortunately he couldn't get it to work. Others tried to modify the equations to account for elliptical orbits (more like planets), but to no avail. In fact using the solar system model only works for hydrogen. This and other experiments within about 12 years forced them to conclude that the solar system model couldn't be correct and that it must be discarded just like the raisin pudding model before it. The current model retains some of the ideas and terminology of the original Bohr model, but it is probabilistic in nature and uses quantum mechanics to describe the atom.

Line Spectra

Again, it is a curious finding that each element gives off characteristic energies (colors) of light and that no two elements have the same set of colors. And this is explained by electrons getting excited and then falling back down to the lower energy state! It is also found that each element only absorbs certain characteristic energies. Furthermore, the energies absorbed and the energies emitted are the same! This again leads to the conclusion that the energy levels are discrete or quantized, like steps on a ladder. The electrons can't be found between steps (energy levels). Due to this quantization the energies observed form a line spectrum. It is distinguished from a continuous spectrum (which has all possible colors) in that only certain colors are observed, with dark spaces in between. Why only certain colors? Because there are only a few energies available for the electron to reside in. Remember, every color of light has a different energy. In this case the energy the light has would correspond to the difference in energy levels within the atom. Each element has its own unique set of energy levels and so each element has a unique line spectrum.

Why Is This Important?

These unique line spectra enable us to determine what elements are in the stars without having a sample of the star's material available to us. Neon signs, night sticks, glow in the dark toys, and fireworks give off their characteristic colors due to this quantum effect.

The selective absorption of light by atoms and molecules gives rise to much of the color around us. The leaves are green because chlorophyll only absorbs the complement of green (when the complement of green is absorbed we see green). Tomatoes are red for a similar reason. The world would certainly be a boring place without quantum effects!

Homework

The homework associated with Day 18 is on Canvas.