Chances are, you’re not like me: you cringe when someone says the word “chemistry.” Your memory of it vaguely harkens back to an undergrad class whose knowledge you flushed once you finished the final. That’s okay — my goal for this series is to change that paradigm, make it understandable to all, and provide some entertainment.
With phlogiston and the four classical elements relegated to mythology, the stage was set for the creation and advancement of knowledge in chemistry. I realize that “creation of knowledge” is a phrase that most don’t encounter in their daily lives, and it begs the question, what does that even mean? For my purposes, I’ll define it as the discovery of time-tested, critical concepts, theories, ideas, etc. still in use today.
The Law of Definite Proportions
One of the first concepts that is a critical foundation for chemistry that we (“we” being chemistry instructors) typically gloss right over is the law of definite proportions. This law states very simply that given a pure substance, its constituent components will combine in definite, constant proportions no matter the source or method of preparation. For example, the law of definite proportion states that carbon dioxide will always derive 27% of its mass from carbon and 73% of its mass from oxygen no matter where it came from. This concept was later rebranded into atomic terms (i.e., instead we say carbon dioxide will always contain one atom of carbon and two atoms of oxygen), and that’s essentially how we teach it now. It seems like common sense, but when Joseph Proust originally proposed it in 1799, it was very much under serious debate, specifically with more famous chemist, Claude Berthollet — a disciple of Lavoisier and co-author of the new chemical naming system.
Proust, while not nobility, was still pretty well off as the son of a pharmacist. He was educated in Paris and spent his life in chemistry despite a small interruption from Napoleon. His investigations on definite proportions started chemists down the path of discovering the atom. To be more explicit, he defined the differences between compounds (substances containing more than one element that follow the law of definite proportions), mixtures (combined substances — but not chemically — and, therefore, don’t follow the law of definite proportions), and solutions (a mixture that involves a substance and a liquid, and again, doesn’t follow the law of definite proportions).
So yeah, we’re heading toward the discovery of the atom, which is crazy to think about. When I think of the discovery of something, I think of that discovery being something tangible, something I can see or touch. Well, an atom wasn’t visualized until 1970 and wasn’t traditionally photographed until 2018, but back in the early 1800s, chemists were starting to accept its existence because these small, indivisible particles called atoms that make up an element were helping to explain a lot of the empirical results.
The Law of Multiple Proportions
Another scientist published something called the law of multiple proportions in 1803 just as Proust published his law. This law was essentially an extension of the law of definite proportions: it said that for different compounds containing the same elements, those ratios of the elements would be whole numbers. Take carbon monoxide (one proportion of carbon and one proportion of oxygen) and carbon dioxide (one proportion of carbon and two proportions of oxygen) for example: the law of multiple proportions says that the ratios of oxygens in those two compounds must be a whole numbers.
It’s best to think of these two laws as a framework for chemical combination, rules for how to think about compounds (or “molecules”). Again, it seems too simplistic, borderline silly, to give these laws any second thought nowadays because the outcome of these laws is so ingrained in us, it’s just second nature. In the early 1800s, though, chemists had to define a foundation; otherwise, you get phlogiston. Do you want phlogiston?!? The answer is no.
By the way, the chemist that proposed the law of multiple proportions? That was John Dalton (1766–1844). Dalton was the first to break the mold of famous chemists with rich parents — he grew up dirt poor and struggled to find formal education. He was smart as a whip, though, studying Newton in his teenage years and teaching at age 12. I don’t know about you, but 16-year-old-me is not spending his free time doing geometry. Another interesting tidbit: while Dalton lived a considerably long time, he suffered two strokes, one of which occurred while arguing about chemistry! What an absolute beast.
Dalton remains one of my favorite examples of an underdog making an world-changing contribution to science. He’s also a prime example of “academic grit,” a truly undervalued characteristic of “smart” people (I’ll talk about grit another time).
The Proposal of the Atom
In formally defining the atom, Dalton had to first give these particles a mass, which was clearly dependent on the element. He started with water, whose proportions were 12.6% hydrogen and 87.4% oxygen by mass. If hydrogen and oxygen form water in a 1:1 ratio (in reality, it’s two hydrogen atoms that combine with one oxygen atom, but Dalton didn’t know this at the time), then that ratio of hydrogen to oxygen must also apply to the individual weights of the atoms that make up hydrogen and oxygen. Since hydrogen was the lightest, it made sense to set its weight to one and compare all the other heavier elements to it. This resulted in a relative “atomic weight” of one for hydrogen and of seven for oxygen.
Quick aside: the atomic weight of hydrogen is actually one (and change), and it was pretty lucky for Dalton to decide to assign it as one. As for oxygen, its atomic mass is 16, and we’ll discuss what accounts for the difference in the next article.
If this seems arbitrary (why does hydrogen get to be the standard?), it’s because it is, a fact that Dalton recognized as well. He postulated that water could be any number of ratios of hydrogen to oxygen in composition (2:1 or 1:2 for example), but avoided that rabbit hole with his premise of simplicity. However, this is a problem that chemists would struggle with for the next half century.
Dalton’s motivation for forming these relative weights began with his work with atmospheric gases, specifically water vapor. From there, he came up with the law for partial pressures and became aware of some interesting solubility phenomena. Working with William Henry (of Henry’s Law fame) on the solubility of gases, Dalton noted that lighter, simpler gases (e.g., hydrogen) were much less soluble than heavier, more complex gases (e.g., carbon dioxide). He needed a way to dig in deeper to this trend, so he invented the relative atomic weight system.
Dalton’s system would be refined for next 50 years as chemists tried to nail down absolute atomic masses rather than relative ones. Nonetheless, he proposed a novel idea that fed into the next wave of research that would become the basis for modern chemistry.
Brock, William H. The Chemical Tree: A History of Chemistry. New York: Norton and Co, 2000.
Ihde, Aaron J. The Development of Modern Chemistry. New York: Dover Publications, 1984.