Topic 1 - Atomic Structure
The word atom comes from the Greek for "indivisible". An atom is the smallest particle of a chemical element that retains the chemical properties of the element. Atoms are composed of subatomic particles, and to understand the behaviour of an atom, we must first understand its constituent particles.
Electrons are tiny electrically charged particles. They have a negative charge, very little mass and they exist in the empty space surrounding the nucleus of the atom, which contains all the other particles. In their elemental states, atoms are not charged, and will have the same number of electrons as they have protons. Electrons can behave as particles and also as waves; this is known as the wave-particle duality of matter. It is only significant for things which are of similar size to atomic particles. Electrons exist in different energy levels or orbitals, filling the lowest energy levels first.
Protons are much larger than electrons. They are about 1800 times heavier than an electron and they have a positive charge equal and opposite to that of an electron.
Neutrons were the last of the nuclear particles to be discovered, as they have no charge so they are not deflected by a magnetic field. They are almost the same weight as protons, and normally there are roughly equal numbers of protons and neutrons in the nucleus.
At the dawn of the 20th century, scientists knew that negatively charged particles existed. Because most atoms have a neutral charge, scientists thought that positive particles might exist to balance the negative particles. They were also curious about how many charged particles were in the atom and how the particles were arranged. Thomson proposed an atomic model in 1904 in response to these curiosities.
In Thomson’s "Plum Pudding Model" each atom was a sphere filled with a positively charged fluid. The fluid was called the "pudding." Scattered in this fluid were electrons known as the "plums." The radius of the model was 10-10 metres. Thomson suggested that the positive fluid held the negative charges, the electrons, in the atom because of electrical forces. However, this was only a very vague explanation and failed to provide any definite answers.
The Millikan Oil Drop Experiment was designed to obtain the charge of an electron which would then allow for the calculation of the electron’s mass. J.J. Thomson had been using an almost identical experiment to that of the Millikan one except he had been using a vapor cloud for the suspension measurement instead of an oil droplet. The water vapor made the experiment too difficult because it evaporated quickly. So Robert A. Millikan of the University of Chicago used droplets of oil for the suspension measurement.
The way this experiment works is that a droplet of oil will come out of an atomizer and go through a tiny slit in an electrode. From the electrode the droplet will pass into a chamber with an electrode parallel from the electrode it had just passed. In this chamber Millikan was able to balance or suspend the droplet by the charge the droplet had picked up when passing through the air. Usining the amount of voltage needed to suspend the droplet he could then calculate the charge the the droplet.
Millikan calculated the force gravity would have on the droplet and then equaled that to the observed amount of charge that was required to suspend the droplet to determine the absolute charge of a particular droplet. When Millikan calculated the force gravity had on the droplets he very soon realized that the amount of charge on each minutely affected the velocity of the fall. This was proof of a very small electron mass. Millikan also noticed that the absolute charge of each droplet was a multiple of a smallest quanta of charge, 1.6 X 10-19 coulombs. He assumed this to be the charge of a single electron. Millikan could now calculate the mass of an electron with Thomson’s ratio of charge to mass (qe/m = 1.76 X 1011 coul/kg)
m = 1.6 X 10-19 coul
1.76 X 1011 coul/kg
m = .091 X 10-30 kg
The electron is apprently has a very small mass. The mass of the electron is 1836 times smaller than that of the hydrogen ion. If you multiply 1836 (.091 X 10-30 kg) you get 1.66 X 10-7 kg which is approximately the value of one atomic mass unit.
Ernest Rutherford was observing the effects of shooting a narrow beam of small alpha particles at a thin gold foil. Rutherford noticed when the alpha particles struck the thin metal, some of them scattered instead of continuing straight through.
The discovery occurred when Hans Geiger, one of Rutherford’s assistants noticed that the number of alpha particles scattered by an angle greater than 10 degrees was much more than predicted.
| Rutherford wrote: | ||
| "I had observed the
scattering of alpha particles, and... it was as if you had fired a
15-inch naval shell at a piece of tissue paper and the shell came right
back and hit you... it was then that I had an idea of an atom with a
minute massive centre, carrying a charge."
Source: Rutherford et al., Project Physics Unit 5 Text, 1971 |
||
From Rutherford’s discovery came the realization of the idea of the nucleus – a small, dense concentration of charge and mass. Rutherford explained that this was possible because any particle that ran into the nucleus would rebound and get deflected, which would not happen if a particle had simply traveled through an atom "cloud." Dr. Geiger led an experiment to verify these conclusions and proved Rutherford correct.
Rutherford’s discovery set the stage for Bohr’s orbital model of the atom, which utilised the idea of a central nucleus. View how the Rutherford experiment worked here.
Atoms are normally described in terms of two key numbers, their atomic number and their mass number.
The number of protons in the nucleus is the most important aspect of an atom. This number determines which element an atom belongs to. The atomic number of an atom can tell you:
- The number of protons in the nucleus of the atom
- The number of electrons in the atom when it is neutral
- The atom's position in the periodic table
Mass Number (A) - Nearly all of an atom's mass comes from the nucleus. Since we know that the mass of a proton is almost equal to that of a neutron, we can measure the mass of an atom in terms of the number of particles in its nucleus. The mass number can tell you:
- The total number of particles in the nucleus
- The number of neutrons in the nucleus (remember to subtract the atomic number)
- The relative atomic mass of an atom
Isotopes are Atoms with the same atomic number but different mass number. Isotopes are shown like this: AZXY
(Where A is the mass number, Z is the atomic number, Y is the charge on
the atom and X is the symbol for that element.) For example, hydrogen
has three isotopes.
Isotope Calculations
It is fairly simple to work out the number of each sub atomic particle from an isotope's symbol. For a simple example, lets look at a calcium ion: 4020Ca2+ Using these rules, you should be able to work out that a calcium-40 ion has 20 protons, 20 neutrons and 18 electrons.
· The number of protons is equal to Z.
· The number of neutrons is equal to A - Z.
· The number of electrons in a neutral atom is equal to Z.
· The number of electrons in a positive ion is equal to Z - Y.
. The number of electrons in a negative ion is equal to Z + Y.
