The idea that matter is made of indivisible particles is ancient. Democritus proposed it in the fifth century BC, Lucretius wrote a poem about it, and for the next two thousand years it remained exactly what it had always been: a philosophical position, argued about but never tested.
What John Dalton did in 1808 was different in one decisive respect. He gave atoms numbers. By assigning each element a relative weight, he turned atomism from a claim about the nature of reality into a tool that chemists could use to predict the outcome of reactions. That is the moment the atom entered science.
A Weather Diary in Manchester
Dalton was born in 1766 into a Quaker family in Cumberland, in the north of England. As a dissenter he was barred from the English universities, so his education came from within the Quaker community and from his own reading. He began teaching at twelve and spent most of his working life in Manchester, giving private lessons in mathematics and natural philosophy.
His first scientific passion was not chemistry but the weather. He kept a meteorological journal from 1787 until the day before his death in 1844, recording well over two hundred thousand observations of temperature, pressure, humidity, and rainfall.
That work led him to a question about the atmosphere that turned out to matter enormously. Air is a mixture of gases with different densities. Why does it not separate into layers, with the heaviest at the bottom?
Pursuing this, he arrived in 1801 at the law of partial pressures: in a mixture of gases, each gas exerts the pressure it would exert if it occupied the container alone, and the total is the sum. Each gas behaves as if the others are not there. Dalton concluded that gases must consist of separate particles, and that particles of different gases must somehow ignore one another.
The Chemistry Behind the Theory
Two established laws provided the framework Dalton needed.
The law of conservation of mass, established by Lavoisier in the 1780s, states that the total mass of the products of a reaction equals the total mass of the reactants. Matter is rearranged, never created or destroyed.
The law of definite proportions, defended by Joseph Proust against considerable opposition, states that a given compound always contains the same elements in the same proportions by mass. Water from any source is always eight parts oxygen to one part hydrogen.
Both laws are simple consequences of an atomic picture. If compounds are formed by combining whole atoms in fixed numbers, and if atoms are neither created nor destroyed, then both results follow immediately.
The Law of Multiple Proportions
Dalton’s own contribution to the evidence was sharper. Some pairs of elements form more than one compound, and he examined the relationship between them.
Carbon and oxygen form two: the gas we now call carbon monoxide and the gas we call carbon dioxide. Take a fixed quantity of carbon in each. The mass of oxygen combined with it in the second compound is exactly twice the mass in the first. Not 1.9 times, not 2.3 times. Twice.
The same pattern appears with the oxides of nitrogen, where the oxygen quantities stand in ratios of small whole numbers. Dalton called this the law of multiple proportions, and it is the strongest argument for atoms available before the twentieth century.
Continuous matter offers no reason for such ratios. If substances combined in arbitrary amounts, the proportions would vary smoothly. Whole number ratios are what you expect if combination means attaching one atom, or two, or three, and nothing in between is possible.
A New System of Chemical Philosophy
Dalton set out his theory in A New System of Chemical Philosophy, published in 1808. Its postulates can be stated compactly:
- Every element consists of atoms, which cannot be created, destroyed, or divided
- All atoms of a given element are identical, and in particular have identical weight
- Atoms of different elements have different weights
- Compounds form when atoms of different elements combine in fixed, simple whole number ratios
- A chemical reaction rearranges atoms without changing them
The second and third points are what made the theory productive. Dalton could not weigh an individual atom, and knew he never would. What he could do was work out relative weights from the proportions in which elements combine. Setting hydrogen at 1, he produced a table of comparative weights for the elements he had data for.
This is the practical core of the whole enterprise. A chemist who knows the relative weights can calculate how much of one substance is needed to react completely with another, without any theory of what atoms are made of.
Where He Went Wrong
Dalton’s weights were, in many cases, incorrect, and the reason is instructive.
To get from measured mass proportions to atomic weights you need to know how many atoms of each element are in the compound, and Dalton had no way to determine this. He adopted what he called a rule of greatest simplicity: if two elements form only one known compound, assume it contains one atom of each.
Applied to water, this gives the formula HO rather than H₂O, and therefore an atomic weight of 8 for oxygen instead of 16. Similar assumptions distorted much of his table.
The solution existed within three years. In 1811 Amedeo Avogadro proposed that equal volumes of gases at the same temperature and pressure contain equal numbers of particles, which allows the formulas to be deduced from volume measurements. His hypothesis was almost entirely ignored for half a century, partly because it required accepting that a molecule of oxygen contains two oxygen atoms, which struck contemporaries as absurd.
Chemistry consequently spent decades with inconsistent atomic weights and competing formulas for the same substances. The confusion was only resolved at the Karlsruhe Congress of 1860, where Stanislao Cannizzaro presented Avogadro’s argument with enough force to convince the assembled chemists. Among those present was Dmitri Mendeleev, who would need a reliable set of atomic weights before the periodic table could be constructed at all.
The Symbols
Dalton also proposed a notation: circles containing dots, lines, or letters, one design per element, which could be grouped to represent compounds. The system is attractive on the page and quickly became unusable as the number of known elements grew.
Jöns Jacob Berzelius replaced it in 1813 with letters taken from the Latin names, giving O for oxygen, Fe for iron, Na for sodium. That is the notation chemistry still uses. Dalton objected to it for the rest of his life.
Atoms as a Convenient Fiction
One point about the reception of the theory is frequently misunderstood. Dalton’s work was accepted quickly and used universally, and yet the physical existence of atoms remained genuinely controversial for another century.
Many chemists treated atomic weights as bookkeeping devices: numbers that make the arithmetic of reactions come out right, without any commitment to invisible particles actually existing. Prominent physicists including Ernst Mach and Wilhelm Ostwald held this position into the 1900s and argued it forcefully. Ludwig Boltzmann spent much of his career defending atoms against exactly this objection.
The dispute was settled by physics, not chemistry. Einstein’s 1905 analysis of Brownian motion predicted precisely how far a suspended particle should wander if it is being bombarded by molecules, and Jean Perrin’s measurements confirmed it. A century after Dalton assigned them weights, atoms were finally shown to be real.
What Came Next
Dalton’s atom was indivisible by definition. That assumption survived until the very end of the nineteenth century, when the discovery of the electron demonstrated that atoms contain smaller components, and the investigation of radioactivity showed that some atoms come apart on their own. Within a decade the atom had a nucleus and the indivisible particle of 1803 was gone for good.
Everything that followed, from atomic structure to quantum theory, depended on first having a quantitative chemistry of elements and compounds to explain. Our edition of Max Planck’s three key publications collects the works where that line of investigation reached the point of forcing a new physics.
Dalton himself remained in Manchester, teaching, measuring the weather, and refusing most honours. When he died in 1844 he was given a civic funeral and roughly forty thousand people filed past his coffin. It was an unusual send off for a schoolteacher who had spent his life weighing gases.
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