Experiment 4: THE ATOMIC THEORY OF MATTER

Introduction

The atomic theory of matter was established originally by experiments involving masses of

elements that combine in chemical reactions. The Law of Conservation of Mass was a

generalization of the observation that when carefully measured, the mass of the products of a

chemical reaction was always equal to the mass of the reactants. No matter is created or

destroyed in a chemical reaction.

In this experiment we study a particular type of compound to show how it illustrates two other

historically important laws, and how the atomic theory could have been deduced from such laws.

In doing this, one attempts to discern the frame of mind of scientists who worked two centuries

ago, before the chemical community finally understood how to determine correct atomic weights

and chemical formulas. You will note that only masses (i.c., not moles) are used in the analysis.

The first of these laws is called the Law of Definite Composition and states that the percent by

mass of the elements in a pure compound is always the same. To put it another way, the relative

amounts of each element in a compound are an essential characteristic of that compound. If the

amounts were different, the compound also would have to be different.

The second of these laws is called the Law of Multiple Proportions and states that if two

elements combine to form more than one compound, the masses of one element that combine with

a fixed mass of the other are in the ratio of a small whole number. This law is best illustrated by an

example. Hydrogen and oxygen combine to form water in which for every 1.000 gram of oxygen

there is always 0.1260 gram of hydrogen. However, hydrogen and oxygen also can combine to

form a different compound, hydrogen peroxide, which has very different properties from water.

In hydrogen peroxide, for every 1.000 gram of oxygen present there is 0.0630 gram of hydrogen.

Now using these data and comparing the amount of hydrogen in each compound that combines

with a fixed mass of oxygen, 1.000 g in this case, one obtains:

0.1260 grams hydrogen / 0.0630 gram hydrogen = 2.00, small whole number.

Mass relationships consistent with this Law can be depicted graphically in separate hypothetical

compounds (represented by cases A and B) made from elements X and Y, or from W and Z. The

amounts of Y combining with X might be represented as in the figure on the left below for

example. Looking at this pair, the amounts of

Y combining with the same amount of X are in

the ratio of 1.6/.8 or 2:1. Elements W and Z

combine as shown by the pair on the right. The

amounts of Z combining with the same amount

of W are in the ratio of 2.7/1.8, which simplifies

to 3:2. In each case the ratios appear to consist

of simple whole numbers.

1gx

1gx

.8gY

A

1.6gY

1gW

1gW

1.8g2

2.7 g 2

B