In this post, I will discuss some of the deductions we can make from Dalton’s atomic theory, which you learned about a few posts back in our study of chemistry.
When I discussed Dalton’s atomic theory, I made sure to differentiate in each postulate between terms that are important to know. Among these terms are “elements” and “compounds.” And if you have been an especially attentive student, then you might have been able to come to the conclusion that Dalton’s atomic theory has been able to explain some of the concepts we previously discussed, such as the law of constant composition.
There are a few other topics we’ll have to consider in this post that are explained by Dalton’s atomic theory, but some of them I have either not gone over or would like to reiterate. Let’s begin by making sure that you have a clear understanding a few different, but related concepts:
The Law of Conservation of Mass: This law dictates that during a chemical change, as long as the system is completely closed, the mass (or quantity) of matter will remain constant throughout (Ebbing and Gammon 2009). This is further explained by the principle that mass is neither created nor destroyed, but can be transformed. For instance, if you set a stack of newspapers on fire (for whatever reason), the resulting burnt paper and ash would retain the same quantity of matter as before the chemical change.
The Law of Constant Composition: This law dictates that the proportion of elements within a given compound is always the same for that compound, such as water; that is, all samples of water will contain hydrogen and oxygen in the same composition by mass (which is 8 g of oxygen and 1 g of hydrogen), according to the law of constant composition (Ebbing and Gammon 2009).
The Law of Multiple Proportions: This law is a little bit more confusing, so I’ll explain it a little differently than I did the others. We know from our previous discussions that a pair of elements can combine to form a compound. I also made the point that there is not just one simple combination to two elements forming a compound. I gave the example of NO (one atom each of nitrogen and oxygen) and N2O (with two atoms of nitrogen and one atom of oxygen). What this law states is that the mass of any given element in these compounds (just pick a single element) is proportional to the mass of the element in the other compound, for a fixed mass. As another example, think of carbon and oxygen and how they might interact. In one case, you might get an atom of the element carbon combining with an atom of the element oxygen, to form CO (carbon monoxide), in which there is 1.000 gram of carbon, and 1.3321 grams of oxygen. In another case, an atom of the element carbon might combine with two atoms of the element oxygen to give CO2 (carbon dioxide). In this case, there is still 1.000 gram of carbon, but now there are 2.6642 grams of oxygen for just that single gram of carbon. What do you notice? In CO2, the mass of oxygen is proportional to that of CO, at a ratio of 2:1 (Ebbing and Gammon 2009). (Phew! This concept seems like one of those things that is difficult to explain, but which we can all demonstrate ourselves, like tying our shoes.)
Now that you’ve learned these three concepts, understand them and know their differences. While it might sometimes seem like they are saying very similar things, they have important and clear-cut differences. Happy studying!
A prospective medical student, looking to help others succeed.