Relative Atomic Weights
Let's look at some of the things you have already learned in this course. You know that elements can combine with one another to form compounds. For example, hydrogen and oxygen combine to form water. Magnesium and oxygen combine with one another to form magnesium oxide. Other elements combine with one another to form a wide variety of compounds. You also know that the weight ratio of the elements in each compound is fixed. That is why we came up with the Law of Constant Composition. You calculated the weight ratios for quite a number of compounds in Lesson 2. The results of those calculations are listed in Example 3 for this lesson. A few of those are shown here.
Consider the first two compounds listed in this table. Let's compare the amount of oxygen that combines with 1 gram of hydrogen in water to the amount of carbon that combines with 1 gram of hydrogen in methane. Dividing 7.94 g of oxygen by 2.98 g of carbon we get a 2.66:1 ratio of oxygen to carbon. Note how that compares very nicely with the weight ratio of oxygen to carbon in the compound we call carbon dioxide. Coincidence? Not at all. It is the basis of the idea of combining weights of the elements. If we take 1 g of hydrogen as its combining weight and use it as a starting point, then 7.94 g is the combining weight of oxygen and 2.98 g is the combining weight of carbon. Not only will those weights of those elements combine with 1 g of hydrogen, they will combine with one another.
With a more complete list of weight ratios for compounds, we would see many more of these interrelationships. Chemists, including Dalton, did compile a more complete list, did see many more interrelationships and did summarize those relationships by making a list of combining weights. Dalton took this a step further and compiled a list of relative atomic weights. Because of limited precision, he presumed that the weights were integer values. The atomic weight of hydrogen was 1 and the atomic weight of oxygen was 8. These weights were relative atomic weights because the actual size and weight of the atoms was not known. But still they were fairly real and useful in explaining the composition of compounds.
Many elements seemed to have more than one combining weight. But they were generally multiples of one another. Dalton theorized that the smaller combining weight was the actual atomic weight and the larger value represented more than one atom of the element. This use of his theory prompted the formulation of the Law of Simple Multiple Proportions.
There were at least two significant problems with values that Dalton used for his atomic weights. One is the presumption that they were integer values. That was taken care of by more precise and accurate measurements. The other was that some of the atomic weights were off by a factor of two or three. Oxygen is a notable example. This was taken care of by the application of Avogadro's Hypothesis to the determination of correct molecular formulas and molecular weights. Dalton determined that the atomic weight of oxygen was 8 based on the presumption that the formula of water was HO. When the formula of water was determined to be H2O, that showed that the correct atomic weight of oxygen was 16.
E-mail instructor: Sue Eggling
Clackamas Community College