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The roots of atomic theory can be traced back over 2000 years ago to the ancient Greeks. The Greeks were responsible for many advancements in civilization, but one of the things that drove them was the question of origin. Who are we? What are we made of? Where did we come from?
These questions were not limited to humans, but to all matter in the world. Around 400 BC a philosopher named Democritus proposed a theory of composition of matter. Democritus proposed that all matter was composed of tiny indivisible particles.
His reasoning was simple. A piece of wood could be divided into two equal pieces. Those two pieces could then be divided again to form four pieces. Those four pieces could be divided to form eight, and then eight to form sixteen, etc. Eventually, Democritus reasoned, a point would be reached where the wood could not be divided any further. This point represented the building blocks of all matter.
In fact, the word “atom” is based on the Greek word “atomos”:
* “tomos”: to cut
* “a”: not
* “atomos”: uncuttable
This Greek influence lives with us today, even if Democritus’ name remains relatively unknown.
Other more famous philosophers are remembered for their theories of the structure of matter. Perhaps the one that is most famous is Aristotle’s theory, which said that matter was composed of four base elements: fire, water, earth, and air.
Why did Aristotle’s theory take hold, and Democritus was forgotten? Most likely the lack of proof. While Democritus’ theory was fundamentally sound, there could never be proof in ancient Greece of this tiny particle.
How could something exist without being able to see it? Aristotle had visual evidence for his theory, and therefore his theory was believable. In retrospect, Aristotle was wrong, but at the time proper scientific principles prevailed.
The alchemists were an interesting group of people that actually never contributed to the development of the atomic model. However they did contribute to chemistry in other ways.
The main quest of the alchemist was to find a way to turn common metals into the precious metal gold. Although we now know that such a feat is impossible (since that would involve actually changing the number of protons in the nucleus of an atom, and that involves some sort of nuclear reaction), the alchemists did not. Many famous people of the Renaissance period were alchemists, including Leonardo DiVinci and Saint Thomas Aquinas.
Their pursuits, while in vain, did contribute to the world of chemistry. The alchemists were the first scientists to perform systematic experimentation. They recorded every experiment they performed, and shared their knowledge with other alchemists. This expanded the knowledge base of chemistry hundred fold. Such recording had never taken place. The knowledge of chemicals and their reactions if the original data that guides us today.
Out of the expermental data, three theories were formed:
* Law of definite composition – a compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample.
* Law of conservation of mass – the mass of a compound is the sum of the masses of the elements that were reacted to produce that compound
* Law of multiple proportions – (applies to different compounds made of the same elements, like H2O and H2O2) – the mass ratio of one of hte elements that combines with a fixed mass of the other element can be expressed in small whole numbers.
A modern movie actually dealt with the subject of alchemy. Hudson Hawk was made in the late 80’s, and stars Bruce Willis and Andie McDowell. If you get a chance, rent it. I find it quite entertaining, but then again I have a wierd sense of humor…
John Dalton was an English schoolteacher (aren’t teachers great?) in the late eighteenth and early nineteenth centuries. In 1808 Dalton proposed an explanation for the three laws that the early chemists had discovered. Dalton reasoned the only explanation was the existance of atoms. His explanation can be summed up in the following statements:
1. All matter is composed of tiny indivisible particles called atoms that cannot be created or destroyed.
2. Atoms of a given element are identical in their physical and chemical properties.
3. Atoms of different elements combine in simple whole number ratios to form compounds.
4. A chemical reaction occurs when atoms are combined, separated, or rearranged.
In statements 1 and 4, Dalton accounts for the conservation of mass. By relating atoms to the measurable property of mass, Dalton makes Democritus’s idea of the atom into a scientific theory. The difference between the two is the fact that a theory can be tested and either proven or disproven.
In fact, Dalton’s theory is not entirely correct. For example, we now know that the atoms is divisible (into protons, neutrons and electrons, and even those can be divided further). Although not all of Dalton’s theory was correct, it gave a starting point for others to test, disprove, or add to in years to come.
The atomic theory proposed by Dalton in the early 1800’s held true until the end of that century. At that point scientists began finding characteristics about the atom that pointed to a more complicated picture than the one painted by Dalton years ago.
The first discovery that altered our picture of the atom actually came from physics, not chemistry. Physicists were working with cathode ray tubes (see diagram below). A cathode ray tube consists of a glass tube with most of the air taken out. When the tube is hooked up to a voltage source, electricity flows from the cathode to the anode, and the remaining gas in the tube glows.
from Chemistry: Visualizing Matter (c.1996: Holt, Reinhart and Winston)
It was observed that when a small paddle wheel (like the ones in old water powered mills) was placed in the beam going between the anode and cathode, the wheel rotated. Reseachers assumed that the electricity must be composed of tiny particles that would be able to rotate the wheel.
Finally in 1897 J.J. Thompson found that the beam going between the anode and cathode could be deflected by bringing a magnet close to the cathode ray tube. The deflection that Thompson observed showed that the beam must have been made up of negatively charged particles. These particles were called electrons. Later experimentation showed the charge and mass of these tiny particles. The mass was proven to be around 2000 times lighter than the mass of the lightest known atom, hydrogen. This lightness implied that atoms must be composed of different types of particles, electrons being quite small and some other matter being quite heavy.
Since it was known that atoms were electrically neutral, this other matter must contain an equal amount of positive charge as the electrons, despite the fact that the positive matter must have an extremely large mass in comparison. From this information, Thompson developed the first basic model of the atom. His model was called the “lum pudding” model, since it resembled plum pudding, which is a dessert common in England that consists of a ball of breading with pieces of fruit stuck in it (sounds appetizing, doesn’t it?). Thompson’s model featured the negatively charged electrons randomly “stuck” into a ball of postively charged matter. In the diagram below, the orange represents the positive matter (the bread in plum pudding) and the blue dots represent the electrons (the fruit in plum pudding).
While this model was more than 2000 years in the making, it didn’t last very long. Within 15 years more discoveries would occur that would move our model much closer to the one we know today.
The next major step in the development of the atomic model occurred in 1911 by New Zealander Ernest Rutherford, a student of J.J. Thompson. Rutherford and his associates performed one of the most famous experiments in chemistry – the gold foil experiment.
Rutherford’s group took a source of alpha radiation (radiation that is positively charged) and fired it at a thin sheet of gold foil. Gold is a good metal for this use because it can be hammered in to a sheet extremely thin and still hold together and retain all of its properties. The gold foil was placed in a chamber lined with what amounted to x-ray film. This film would show any radiation that either passed through the gold foil or was redirected by the foil.
The group’s hypothesis was that the radiation would pass right through the foil, since the plum pudding model of the atom had mass and charge evenly distributed through out the atom. The majority of the radiation did just that – passed right through. However a measureable amount of radiation was deflected, not only to the right and left of the expected target point, but some radiation was deflected backwards toward the source (see the diagram above – the grey box in the lower left is the radiation source, and the yellow trails are the paths of the radiation).
The only possible solution to these results meant that the Plum Pudding model of the atom was incorrect. Rutherford, along with his top assistant Niels Bohr, developed their own model of the atom. The data suggested that the majority of the mass of an atom must be concentrated in the center of the atom, which Rutherford called the nucleus. This was evidenced by the fact that the majority of the radiation passed through the gold foil. The nucleus must have a positive charge, the scientists figured, since that would be the force that would cause the positive radiation to deflect.
As for the negative charge of the atom (which had to exist to make the atom neutral), Rutherford did not have any experimental evidence to show where or how that was organized. He did hypothesize that the electrons must be extremely tiny in mass, since no attraction occurred between the electrons and the positive radiation (that is, none of the radiation stuck in the gold foil). To account for the electrons, Rutherford simply said they were randomly distributed around the nucleus.
While Rutherford’s atom was a major step forward in the atomic model, it was not entirely correct. Niels Bohr continued to fine tune the model in the area of the electrons, and that is described on the next page.
Niels Bohr was one of the scientists that worked with Rutherford and was one fo the major players in developing Rutherford’s model of the atom (in fact, the model is sometimes called the “Rutherford-Bohr” model).
Rutherford did a wonderful job discovering and proving the existance of the nucleus. However, Rutherford could not explain the negative charges, the electrons. His model left the electrons “orbiting” the nucleus in some randon way, shape, or form.
Bohr continued experimenting with the atom in an attempt to further explore its structure. He found that when atoms were exposed to a source of energy (for example, light or electrical energy) the atoms give off light. Bohr reasoned that the electrons were responsible for the light emission. The electrons absorb the energy that is given to the atom and jump to a different energy level, then settle back to their original position. When the settling occurs, the energy that was absorbed is given away in the form of light. However the light that was emitted was not in a continuous spectrum, but existed in very specific wavelengths. The different wavelengths given off represented the different amounts of energy the electrons absorbed. The emission spectrum for hydrogen is shown below, along with the experimental set-up.
from Chemistry (copyright 1998, McGraw-Hill Companies)
Bohr hypothesized that the reason for this type of light emission was due to the fact that the electrons in the atoms were not randomly thrown around the nucleus as Rutherford guessed, but existed in specific orbits. This was based on the work of Max Planck, who proposed the quantum theory of energy. The quantum theory states that energy does not exist continuously, but instead exists in tiny packets, called photons.
Bohr figured that since the energy existed in specific quantities, the electrons must exist in specific orbits around the nucleus. That is, if an electron absorbs a certain amount of energy, that electron must move a specific distance to a different orbit. This model (the planetary model) is probably familiar to you from ninth grade physical science. An example is shown below, with the nucleus in the center of the atom and the electrons existing in the orbits around the nucleus.
This model of the atom also has numbers of electrons associated with each orbit (2,8,8,18,18,32,32). While these numbers are important, as we will look at a little later, the model is not entirely correct.
The Quantum model
Here is where things get a little fuzzy (literally). The next logical step in the progression of the atomic model is to actually locate the electrons and measure the distance of the orbit they are in. The problem with these little electron thingies, though, is the fact that they are so incredibly small. Actually locating them is quite an adventure.
Not only is the position of the elctron important, but the velocity of it is important as well. By knowing the position and velocity of an electron at one moment, we can make a prediction of the electron’s position and velocity at a later time.
Here is where things get interesting. A German scientist named Werner Heisenberg made an important observation. Hiesenberg proposed that it is impossible to know both the exact position and exact velocity of an electron at the same time. Huh?
Let’s examine this a little closer. When we look at an object like a computer or a car we are seeing the light waves that are reflected from that object. In other words, to see an object a photon of light must hit the object and reflect back to our eyes. These collisions between a large car and a photon of energy has a negligable effect on the car.
When a photon of light collides with an electron results in a large effect on the electron because of its size. If we were able to pinpoint an electron’s position, we cannot know its velocity because the collision with light has changed it (like two billiard balls colliding). If we pinpoint an electron’s velocity we cannot know its position because it changes too quickly. This concept is called Heisenberg’s Uncertainty Principle.
Because of this uncertainty describing the electrons in terms of specific orbits is impossible. Instead we can only give a possible area the electrons could be located in at any specific time. These possible locations are called electron clouds. Instead of orbits, each different energy level can abe described as a different cloud.
We are going to stop our discussion of the development of the atom will stop here. The shapes of the clouds and the quantities of energy associated with each level are not important to us at this level of chemistry.
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