Every first-year chemistry class examines atomic theory and the periodic table, following a typical path for each. To motivate atomic theory, perhaps a demo or experiment showcasing conservation of mass is performed, after which the postulates of atomic theory are presented to be learned. However, many of our curricula tend to minimize the contributions of Dalton relative to later development of the periodic table and modern atomic structure. Often, many assessment questions on “classic” atomic theory are little more than identifying postulates of Dalton’s theory that could be answered simply through strong work in reading comprehension.
If you think about it, many properties of importance in modern chemistry and physics arise simply from the particular nature of matter. Indeed, Richard Feynman is famous for stating:
"If, in some cataclysm, all of scientific knowledge were to be destroyed, and only one sentence passed on to the next generations of creatures, what statement would contain the most information in the fewest words? I believe it is the atomic hypothesis (or the atomic fact, or whatever you wish to call it) that all things are made of atoms—little particles that move around in perpetual motion, attracting each other when they are a little distance apart, but repelling upon being squeezed into one another. In that one sentence, you will see, there is an enormous amount of information about the world, if just a little imagination and thinking are applied.1"
Of course, Dalton did not have the benefit of a deep understanding of electrostatic forces that is implied in Feynman’s remarks. However, he specifically tells us where to look when trying to motivate and understand his atomic theory deeply:
“In all chemical investigations, it has justly been considered an important objective to ascertain the relative weights of the simples which constitute a compound. But unfortunately, the enquiry has terminated here; whereas from the relative weights in the mass, the relative weights of the ultimate particles or atoms of the bodies might have been inferred, from which their number and weight in various other compounds would appear, in order to assist and to guide future investigations, and to correct their results. Now it is one great object of this work, to shew the importance and advantage of ascertaining the relative weights of the ultimate particles, both of simple and compound bodies, the number of simple elementary particles which constitute one compound particle, and the number of less compound particles which enter into the formation of one more compound particle.”2
Dalton’s quote gives us a starting point to bring quantitative data to allow our students to develop and model atomic theory as he intended so that we can practice scientific methods as opposed to only reading about history. The following is the first of two articles that provide an overview of a way to make atomic theory more quantitative in your classes. This article provides an introduction and a method to introduce the masses to the periodic table that is both interesting and that can also build confidence and familiarity with simple mass calculations long before moles are studied. If a deeper dive interests you, the second article provides details on a lab activity that has been created to motivate all the postulates of atomic theory from a data-driven perspective.
I really enjoy bringing a historical perspective to the classroom. Many students can’t identify with chemistry early on, and the more opportunities we provide to have them engage with data they believe is easy to collect in the lab, the better. Like many of us, I perform demonstrations of conservation of mass and assign simple homework questions to motivate what kind of data Dalton and his contemporaries could realistically collect at the turn of the 18th century. When ready to take a deeper dive, I tell student that they are assistants in Lavoisier’s lab, collecting data on the combination of carbon and oxygen via combustion. I invite students to offer how much carbon each wants to burn in a combustion chamber. For example, a student may say “5.00 g of C”, after which I tell them that this experiment produces 18.3 g of product. We proceed to another student who may say 18.00 g of C, after which I will tell them 42.00 g of product is made. In this way, we generate a table like the one shown below for us to discuss together:3
This information is all I write on the board, but I then start to ask the class what information we can glean from the data obtained? Is it possible to get the mass of oxygen in the compound? Recalling that we have just discussed conservation of mass, the class goes about calculating the amount of oxygen in each of the samples. Note that we also pay attention to precision in the data through our calculations.
Although the class feels comfortable with the calculation, they don’t necessarily see where we are going with this line of reasoning. I then ask, 'How many unique compounds do you think we have made? ' As students have little experience with synthesis, I remind them that we could have made five different compounds, the same compound five times, or something in between. Does the data allow us to attempt to answer this question?
Before Dalton wrote A New System of Chemical Philosophy, two of the three fundamental laws of chemistry were known. We have already used conservation of mass, but the second, the law of definite proportions4, provides us with an answer to the question posed above. The statement of the law, that the ratios of the masses of elements in a compound are always fixed, can be illustrated by calculating the ratio of the mass of oxygen to the mass of carbon in each sample:
Note that once the ratio is created, patterns immediately appear in the data, and the class agrees that likely we have created three unique compounds, and I label these A, B, and C respectively. What can we learn about these compounds?
One of the biggest challenges of 19th century chemical analysis was to determine the molecular formula of compounds, particularly organic compounds. Modern methods of structural determination were more than a century away, and thus analysis based on mass was the one of the only effective methods. As we show our students, percent composition data will easily yield the empirical formula of the compound. But, without knowledge of the molar mass, the molecular formula is elusive. This problem will be addressed in an upcoming article.
Even if we do not have formulas, some interesting fundamental calculations can be performed with the data. For example, consider compound A. Let’s say you have a sample of this compound but instead of 5.00 g of C in the compound, you have 13.0 g of C. Because the ratio of the mass of O to the mass of C must be fixed by definite proportions, we can easily calculate both the mass of O and the total mass of the compound.
These calculations can be done for any compound and provide an early gateway to what will be done later in stoichiometry. What I want students to focus on is the fact that the ratio of 13.0 to 5.00 is the same as 2.66 to 1.00, and that ratio is always fixed for A by the Law of Definite Proportions.
These calculations become much more enlightening if we do know the formula of one of the compounds. To illustrate this, consider compound B, which I tell the class has the formula CO (obtained through some means we don’t know yet)
Note that for B, the ratio of O to C is 24.00 to 18.00 or 1.33 to 1.00. So, let’s consider three calculations.
What mass of O is in a sample of B that contains:
- 67.8 g of C
- 0.723 g of C
- 12.0 g of C
In each case, the calculation follows an identical path:
Look at the last answer in more detail. I ask students to compare the last calculation to the information in the periodic table, and they observe:
This is typically very intriguing and perhaps confusing for the class. Why are the numbers for our final calculation the same as the masses in the periodic table? Since we know the formula of the compound is CO, that means that there is always one oxygen in the compound for each carbon. Thus, the number 1.33/1.00 is not just the ratio of the mass of all the O to the mass of all the C in B, it must also be the mass of one oxygen atom to one carbon atom, or that oxygen weighs 1.33 times more than carbon! Note that the masses in the table are given no units – C is just 12.0 (to three significant figures). This is the relative mass of C, and by itself has no meaning. But, if C is defined as 12.0, then O must be 16.0 because an oxygen atom always weighs 1.33 times more than a carbon atom from composition data. Thus, Dalton was one of the first to actually “derive” masses for the elements. For any one element pick a number and call that its relative mass. However, to correctly derive the masses of all the other elements, both mass and formula data for a compound containing the two elements must be known. Obtaining the masses of the elements is usually straightforward via experiment, but obtaining the molecular formula of the compound is a much stickier proposition.
For many classes, making this connection may be as far as you want to go, and the attached handout shows some potential questions you can ask based on this approach. However, the observation about relative masses opens the door to a completely data driven approach to atomic theory. This will be discussed in the next article.
- Feynman, Richard P., Robert B. Leighton, and Matthew Sands. The Feynman Lectures on Physics. Vol. 1, Addison-Wesley, 1963, Ch. 1, “Atoms in Motion.”
- Dalton, John. A New System of Chemical Philosophy. Part I, Chapter III, “On Chemical Synthesis,” 1808. Quoted in “John Dalton,” Dalton Web Zone, Le Moyne College, web.lemoyne.edu/giunta/dalton.html. Accessed 01 Sept. 2025.
- I do this in class in real time with a calculator, deciding which compounds I will make on the spot. This could of course also be pre-made, but having students be an active part of the data generation is key to the modeling process
- The law of definite proportions is ascribed to Joseph Proust, an important chemist for his early quantitative work on compound composition. See Ashworth, William B., Jr. “Scientist of the Day – Joseph Louis Proust.” Linda Hall Library, 26 Sept. 2018, www.lindahall.org/about/news/scientist-of-the-day/joseph-proust/. Accessed 01 Sept. 2025.
