0.1 Atomic molecular theory  (Page 2/6)

Determining whether a pure substance is an element or a compound is a difficult and time-consuming process of experimentation. We will assume for our study that the elements have all been identified.

Since matter is anything that can have mass, we will spend much of our time in this study analyzing mass. Without proving it, we will assume the validity of the “Law of Conservation of Mass,” an experimental result that simply says, “The total mass of all products of a chemical reaction is equal to the total mass of all reactants of the reaction.” In other words, matter cannot be created or destroyed by chemical or physical processes. This law makes it possible for us to measure masses of materials during reactions knowing that these masses aren’t variable or unpredictable.

With these assumptions in mind, we can proceed directly to experiments which led to the development of the Atomic Molecular Theory.

Observation 1: mass relationships during chemical reactions

Since matter is anything that has mass, then the Law of Conservation of Mass suggests that matter is also conserved during chemical reactions: whatever we start with, we wind up with, at least in total. However, this does not mean that matter must be made up of atoms. It simply says that matter is reorganized in some way to produce new substances with new properties when a reaction takes place.

Since we know that all substances are made of elements, then we can analyze the masses of the elements that participate in a chemical reaction. Most importantly, we can take a compound, break it into the elements that it is made of, and then find the masses of those elements. From the Law of Conservation of Mass, the total mass of the elements that make up the compound must equal the mass of the starting compound. For example, a particular compound called copper carbonate is composed of the elements copper, carbon, and oxygen. If we take a 100.0 g sample of copper carbonate, we find that it contains 51.5 g of copper, 38.8 g of oxygen, and 9.7 g of carbon. The total of these three masses is 51.5 g+38.8 g+9.7 g = 100.0g and is the same as the mass of the copper carbonate.

This turns out to always be true. It does not matter what sample of copper carbonate we analyze, where it came from, who gave it to us, or how we made it. We always get these same masses of the element.

What if we take 200.0 g of copper carbonate instead? Experiments show us that we get 103.0 g of copper, 77.6 g of oxygen, and 19.4 g of carbon. The total of these three masses is 103.0 g+77.6 g+19.4 g = 200.0g, so again the total mass is conserved. Even more importantly, there is something striking about these numbers when compared to the 100.0 g sample. When we double the mass of the copper carbonate, we also double the amount of copper it contains: 200.0 g of copper carbonate contains 103.0 g of copper, and 100.0 g of copper carbonate contains 51.5 g. The same is true for the amounts of oxygen and carbon.

One way to look at this is that the fraction of the copper carbonate which is copper is the same in both samples: 51.5 g/100.0 g is equal to 103.0 g/200.0 g, which is equal to 51.5%. This is a very important result. After looking at data from many experiments, we find that regardless of the specific sample of copper carbonate, regardless of the mass of that sample, and regardless of where that sample came from, the fraction of the mass of the sample which is copper is always the same, 51.5%. We get similar results for the oxygen and carbon. The fractions of the mass of every sample of copper carbonate are always 51.5% copper, 38.8% oxygen, and 9.7% carbon.