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Changes in Atomic Weight StandardsValues for relative atomic weights have not remained the same since Dalton's time. They have changed over the years. This table shows a little bit of the history of atomic weight standards and the changes that have been made over the years. (This is also shown in Example 6 in your workbook.)
In the early 1800s hydrogen was used as the standard with a weight of one unit. Other elements are compared to it. Oxygen had a relative weight of 16, carbon a relative weight of 12, nitrogen a relative weight of 14, and gold a relative weight of 196. By the middle 1800s chemists had done enough research to realize that the relative atomic weights were not really integer values. They were dealing with measured values that were close to integers in many cases but they were not exact whole numbers. (Berzelius published a list of atomic weights in 1828.) With additional precision and care taken in the measurements and still using hydrogen as a standard of exactly one, oxygen was found to have an atomic weight of 15.87, and carbon was found to have a relative weight of 11.92, nitrogen 13.90, and gold 196.41. Sometime in the mid-1800s, but I am not sure exactly when it was, chemists switched to oxygen as the standard. Oxygen became the standard because it was more readily available and it made more combinations with other elements. But rather than set oxygen = 1, they assigned oxygen the value of 16 units so that the relative atomic weights stayed about the same as they were before. They actually turned out to be closer to integer values, although that is not important. Precision and accuracy of atomic weights continued to improve. (In 1916 the American chemist Theodore Richards received the Nobel Prize for his exact determination of many atomic weights.) In the early 1900s, physicists made a shift to a new standard. The reason behind this change was that they learned about isotopes in the early 1900s. They discovered that not all atoms of a particular element had the same weight. Although most oxygen atoms weighed 16 units, some weighed 17 units and some weighed 18 units. Each isotope of an element had its own isotopic weight and the atomic weight for each element was an average value of those isotopic weights. To talk about the average weight became quite a problem in some of the work that they were doing, so they shifted their reference standard to one particular isotope of oxygen called oxygen-16. If you look down those two columns, you can see that the weights are pretty much the same. You have to go out about to the fourth or fifth significant digit before you start getting any discrepancy between the weights that the physicists and the chemists used. (There is a 0.027% difference between the values.) Nevertheless, that difference was there, and it was not a good thing. The last column shows the current standard. Chemists and physicists simply have too much in common to abide by two different sets of atomic weights; and so in 1961, they settled on the isotope carbon-12 as the standard at exactly 12. Settling on a particular isotope rather than an elemental average gave physicists the precision and reliability that they needed. Settling on that particular isotope kept the numbers pretty close to the previous weights that the chemists had been using. So there you have a brief history of the atomic weights standards. You will soon be using atomic weights to do a variety of calculations. A short list of atomic weights is in Example 5 in your workbook. Working with just the seven elements listed in Example 5 will be much too limiting before long. For the atomic weights of the other elements, you can use the longer alphabetical list of elements, symbols, and atomic weights in your textbook or refer to a periodic table.
E-mail instructor: Sue Eggling Clackamas Community College |
