John Dalton, a British meteorologist and chemist working in the early 1800s, is recognized for proposing the first comprehensive atomic theory of matter. Dalton’s work was a powerful synthesis of existing chemical observations, transforming the ancient, philosophical concept of the atom into a modern, testable scientific theory. His genius lay in interpreting the precise, quantitative data gathered by his contemporaries to build a physical model of the atom.
The Quantitative Foundation of Chemical Combination
Dalton built his theory upon two distinct laws governing chemical reactions that were already well-established through careful, quantitative laboratory measurements. The first was the Law of Conservation of Mass, formulated by Antoine Lavoisier, which demonstrated that matter is neither created nor destroyed during a chemical change. Lavoisier showed that the total mass of the reactants before a reaction precisely equaled the total mass of the products afterward.
The second foundational principle was the Law of Definite Proportions, championed by Joseph Proust. This law stated that a pure chemical compound always contains the same elements combined together in the same proportion by mass. For example, any sample of water always consists of hydrogen and oxygen in a mass ratio of 1:8. These two laws provided Dalton with a compelling puzzle: why did chemical reactions always involve such specific and unchanging mass relationships?
These quantitative regularities suggested that matter was composed of discrete, unchanging units rather than a continuous substance. Dalton hypothesized that these units, which he called atoms, must possess a characteristic mass for each element. He proposed that chemical reactions involved the combination or separation of these entire, indivisible particles.
Proving the Law of Multiple Proportions
The most specific experimental evidence Dalton provided to support his atomic hypothesis was his work on the Law of Multiple Proportions. This law addresses situations where two elements can combine in more than one way to form two or more different compounds. Dalton observed that in these cases, the masses of one element that combine with a fixed mass of the other element are related by ratios of small whole numbers.
A clear example involves carbon and oxygen, which can combine to form carbon monoxide (\(\text{CO}\)) and carbon dioxide (\(\text{CO}_2\)). In carbon monoxide, approximately 1.33 grams of oxygen combine with a fixed mass of 1.00 gram of carbon. For carbon dioxide, 2.66 grams of oxygen combine with that same 1.00 gram of carbon. The ratio of the oxygen masses combining with the fixed mass of carbon is 2.66 to 1.33, which simplifies exactly to a 2:1 ratio.
This simple, whole-number ratio is explained only if the elements are composed of atoms that combine as whole units. In carbon monoxide, one carbon atom combines with one oxygen atom, while in carbon dioxide, one carbon atom combines with two oxygen atoms. Dalton performed similar quantitative analyses on other pairs of elements, consistently finding these simple mass ratios.
The Postulates Derived from Experimental Evidence
Dalton’s atomic theory took the empirical evidence from these three laws and translated them into a set of theoretical postulates about the nature of matter. The Law of Definite Proportions led to the conclusion that atoms of a given element are identical in mass and properties.
The Law of Conservation of Mass was explained by the postulate that atoms cannot be created, destroyed, or subdivided during a chemical reaction. They are simply rearranged to form new combinations.
The conclusion derived from the Law of Multiple Proportions established that atoms combine in simple, whole-number ratios to form compounds. Dalton’s ability to use quantitative experimental data to support a corpuscular theory of matter allowed him to calculate the first table of relative atomic weights for the elements known at the time. This achievement marked the beginning of modern chemistry, establishing a theoretical framework that could predict and explain the results of chemical experiments.