Dalton's Atomic Theory

In the early nineteenth century, chemistry lacked a unified framework to explain why substances always combined in fixed proportions or why the mass of reactants always equalled the mass of products. John Dalton, an English chemist and schoolteacher, addressed these puzzles by proposing a systematic atomic theory between 1803 and 1808. His theory gave physical meaning to the ancient Greek idea of the atom and transformed chemistry from a collection of empirical recipes into a rigorous quantitative science. Dalton's Atomic Theory provided the microscopic explanation for the macroscopic Laws of Chemical Combination — specifically the Law of Conservation of Mass (Lavoisier, 1789), the Law of Definite Proportions (Proust, 1799), and the Law of Multiple Proportions (Dalton himself, 1803). By postulating that matter consists of indivisible atoms with fixed masses, Dalton could derive all three laws as logical consequences of atomic behaviour rather than coincidental observations. Although later discoveries — electrons (1897), radioactivity (1896), and nuclear structure (1911) — showed that atoms are neither indivisible nor indestructible in every circumstance, Dalton's core ideas remain the foundation of chemical stoichiometry. Every mole calculation, every balanced equation, and every concept of molecular formula taught in chemistry today rests directly on the postulates Dalton published in his 'New System of Chemical Philosophy' (1808).

Key takeaways
  • Postulate 1 — Matter is Made of Atoms
  • Postulate 2 — Atoms of the Same Element Are Identical
  • Postulate 3 — Atoms Cannot Be Created or Destroyed
  • Postulate 4 — Compounds Form by Fixed Atom Ratios
  • Postulate 5 — Multiple Compounds from the Same Elements

Postulate 1 — Matter is Made of Atoms

All matter is composed of extremely small, indivisible particles called atoms. The word 'atom' comes from the Greek 'atomos' meaning uncuttable. This postulate established that there is a fundamental, discrete unit of matter — not a continuous substance. Modern revision: atoms can be split by nuclear reactions, but chemical reactions do not split atoms.

Postulate 2 — Atoms of the Same Element Are Identical

All atoms of a given element are identical in mass and chemical properties. Atoms of different elements have different masses and properties. This is why every atom of carbon-12 has a mass of exactly 12 atomic mass units (u) and why carbon always behaves as carbon in a reaction. Modern revision: isotopes exist (e.g., ¹²C, ¹³C, ¹⁴C), so atoms of the same element can differ in mass — but average atomic masses account for isotope abundances.

Postulate 3 — Atoms Cannot Be Created or Destroyed

Atoms are indestructible in chemical reactions. They can neither be created from nothing nor destroyed into nothing. This postulate directly explains the Law of Conservation of Mass: if no atoms appear or disappear, the total mass before and after a reaction must be equal. For example, in \ce{2H_2 + O_2 -> 2H_2O}, four hydrogen atoms and two oxygen atoms appear on both sides of the equation.

Postulate 4 — Compounds Form by Fixed Atom Ratios

Atoms of different elements combine in simple, whole-number ratios to form compounds. A given compound always contains the same types of atoms in the same numerical ratio. Water is always H₂O (2:1 ratio of H to O), never H₁.₉O or H₂.₁O. This explains the Law of Definite Proportions: because the atom ratio is fixed, the mass ratio is also fixed. Mass of H in water = 2 × 1.008 u = 2.016 u; mass of O = 15.999 u; ratio H:O by mass = 2.016 : 15.999 ≈ 1 : 7.94, always.

Postulate 5 — Multiple Compounds from the Same Elements

Elements can combine in more than one ratio to form different compounds, but those ratios are always simple whole-number multiples of each other. This explains Dalton's own Law of Multiple Proportions. Carbon and oxygen form two compounds: CO (carbon monoxide, 1:1 ratio) and CO₂ (carbon dioxide, 1:2 ratio). If the mass of carbon is held constant, the masses of oxygen in CO and CO₂ stand in the ratio 1:2 — a small whole number, not a random fraction.

The Key Equation

\ceC+O2−>CO2\ce{C + O2 -> CO2}
  • \ce{C} — Carbon atom (12.011 u): Each carbon atom contributes exactly 12.011 atomic mass units. The fixed mass of the carbon atom, as stated in Postulate 2, ensures a predictable mass contribution to any compound it forms.
  • \ce{O2} — Oxygen molecule (31.998 u): Molecular oxygen consists of two identical oxygen atoms bonded together. In the reaction, both oxygen atoms are incorporated into CO₂, illustrating that atoms recombine but are not created or destroyed (Postulate 3).
  • \ce{CO2} — Carbon dioxide (44.009 u): The product compound always contains carbon and oxygen in the whole-number ratio 1:2 (Postulate 4). Mass % of C = (12.011/44.009) × 100 = 27.29%; mass % of O = 72.71% — values that are constant for any pure sample of CO₂.

Worked Example

Problem. Two oxides of nitrogen are analysed. Oxide A contains 30.43% nitrogen by mass; Oxide B contains 46.67% nitrogen by mass. Show that these data are consistent with Dalton's Law of Multiple Proportions, and identify the likely atom ratios (N:O) in each oxide.

mO(in \ceNO2)mO(in \ceNO)=2×15.9991×15.999=21\frac{m_{O}(\text{in } \ce{NO2})}{m_{O}(\text{in } \ce{NO})} = \frac{2 \times 15.999}{1 \times 15.999} = \frac{2}{1}

Answer. The ratio of oxygen masses combining with equal masses of nitrogen is 2:1 (a simple whole-number ratio), confirming the Law of Multiple Proportions. Oxide A is NO₂ and Oxide B is NO.

Figure
Diagram
Dalton's Atomic Theory: Multiple Proportions VisualisedCO vs CO₂ — same elements, different whole-number ratiosCarbon Monoxide (CO)C : O atom ratio = 1 : 1C12.011 uO15.999 uTotal mass = 28.010 uCarbon Dioxide (CO₂)C : O atom ratio = 1 : 2OCOTotal mass = 44.009 uLaw of Multiple Proportions — Mass of O per fixed mass of CIn CO:15.999 g Oper 12.011 g C→Ratio = 1 : 2(simple whole number!)In CO₂:31.998 g Oper 12.011 g C
Origins · 1 minJohn Dalton and how this was discovered
JD
John Dalton
1766–1844 · British (English)

Dalton proposed the first modern quantitative atomic theory (1803–1808), explaining the Laws of Conservation of Mass, Definite Proportions, and Multiple Proportions through the behaviour of discrete, mass-specific atoms. He also produced the first table of atomic weights and pioneered the use of chemical symbols to represent atoms in combinations. His 'New System of Chemical Philosophy' (1808) remains one of the most influential texts in the history of science. Dalton suffered from red-green colour blindness — a condition now sometimes called 'Daltonism' in his honour. He studied his own condition scientifically and in 1794 wrote the first detailed scientific paper ever published on colour blindness. He even requested that his eyes be examined after his death to investigate the physical cause, making him a pioneer in both atomic theory and vision science.

  • Law of Conservation of Mass — Dalton's Postulate 3 (atoms cannot be created or destroyed) is the atomic-level explanation for Lavoisier's Law of Conservation of Mass. Understanding both together links the macroscopic observation to the microscopic model.
  • Law of Definite Proportions — Proust's law states that a compound always contains elements in fixed mass ratios. Dalton's Postulate 4 explains this: because atom ratios are fixed whole numbers and atomic masses are constant, mass ratios must also be constant.
  • Law of Multiple Proportions — Dalton himself discovered this law and used it as evidence for his theory. It is a direct prediction of Postulate 5 and the worked example in this article demonstrates how to apply it quantitatively.
  • Atomic Structure — Later discoveries of protons, neutrons, and electrons refined Dalton's model. Studying atomic structure shows which postulates remain valid (fixed mass ratios) and which required revision (atoms are divisible; isotopes exist).
  • Mole Concept and Avogadro's Number — The mole is the chemist's counting unit for atoms and molecules. It quantifies Dalton's whole-number atom ratios at a laboratory scale, connecting atomic theory directly to measurable masses of substances.