Most of us are familiar with the representation of the atom shown in Figure 3, proposed by Niels Bohr in 1913: a nucleus orbited by electrons. While there have been many major additions to our knowledge about the atom over the past hundred years, the Bohr model still provides a convenient way to look at the atom - even though it was originally missing one major component (the neutron), which we'll get to in a moment.
During the early twentieth century, a number of physical chemists worked on the relationships between elements. Hydrogen was, by then, known to be the lightest of the elements and uranium thought to be the heaviest. In 1907, J.J. Thompson constructed an apparatus to measure the relative weights of the elements - and ran into a problem. The weights of the elements did not correspond to the integers used to classify their position on the table of elements. Moreover, he found some atoms that seemed to have two (or more) different weights according to his "mass spectrometer." Neon, for instance, was found to have two weights - 20 and 22 - when compared with the weight of a hydrogen atom.
In 1932, spurred on by his mentor, Ernest Rutherford, James Chadwick of Cambridge University discovered the existence of the neutron, an atomic particle with the same weight as the proton, but without an electrical charge. It soon became clear that, while elements were identified (and reacted chemically) by their number of protons, many had differing numbers of neutrons. These variations became known as isotopes. [Actually the word isotope had been coined by Nobel laureate Frederick Soddy almost twenty years earlier; he just didn't know why they existed until the discovery of the neutron.]
It is important to note that the atomic number of an element is equal to the number of protons in the nucleus, while the atomic weight is the sum of the protons and neutrons.
Because isotopes have different numbers of neutrons, most atomic weight figures in the periodic table are not integers, but a weighted average of the atomic weights of all the isotopes of that element. We'll be using two conventions to denote the atomic weight of an isotope. For example, a carbon atom with eight neutrons (and its mandatory six protons, if it is to be carbon) will be written simply as carbon 14 or, alternatively, with the atomic weight superscript written before the elemental symbol, as in "14C." (You may find in other literature that the superscript appears after the element symbol.)
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Caption of Figure 3: Representation of Carbon 12 Atom
The carbon atom has 6 protons and 6 electrons giving it an atomic number of 6. Most carbon atoms have 6 neutrons which, when added to the protons, give it an average atomic weight of about 12.
If the atom were drawn to scale with the orbits (or fields) of the electrons being the size of an average bedroom, the electrons would be microscopic, while the nucleus would be about the size of a pin head. It is only in the tiny nucleus that atomic phenomena (such as radioactivity) occur.
On the other hand, chemical properties of an element are related to its atomic number (the number of protons) - although it is actually the electrically counteracting electrons that share orbits (or fields) with other elements to make chemical compounds.
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