The Early History of Chemistry

Fundamental Chemical Laws

The development of chemistry as a science was guided by several fundamental laws, established through careful experimentation and observation. These laws laid the foundation for our understanding of chemical processes and the nature of matter.

• Law of Conservation of Mass: In a chemical process, mass is neither created nor destroyed. Antoine Lavoisier (1743–1794) demonstrated this law in 1774, which was later published in his textbook Elementary Treatise on Chemistry (1789). Example: When burning wood, the total mass of the products (ash, gases) equals the mass of the original wood and oxygen used.

• Law of Definite Proportion: A given compound always contains exactly the same proportion of elements by mass. Joseph Proust (1754–1826) proposed this law in 1794, accepted in 1812. Example: Water (H2O) always contains hydrogen and oxygen in a mass ratio of approximately 1:8.

• Law of Multiple Proportions: When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first element can be reduced to small whole numbers. John Dalton (1766–1844) formulated this law in the early 1800s. Example: Carbon and oxygen form CO and CO2; the ratio of oxygen masses that combine with a fixed mass of carbon is a simple whole number (2:1).

Dalton's Atomic Theory

In 1808, John Dalton presented his atomic theory, which provided a conceptual framework for understanding chemical reactions and the composition of matter.

• Postulate 1: Each element is made up of tiny particles called atoms. (The term 'atom' comes from ancient Greek philosophy—see Leucippus and Democritus.)

• Postulate 2: The atoms of a given element are identical; atoms of different elements are different in some fundamental way or ways.

• Postulate 3: Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms.

• Postulate 4: Chemical reactions involve reorganization of the atoms—changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

Dalton also assumed that nature was as simple as possible, which led him to prepare the first table of atomic masses (weights).

Early Gas Laws and Avogadro's Hypothesis

In 1809, Joseph Gay-Lussac discovered that at the same temperature and pressure, gases react in simple volume ratios. For example, 2 volumes of hydrogen gas react with 1 volume of oxygen gas to produce water.

Amedeo Avogadro (1776–1856) proposed in 1811 that equal volumes of gases, at the same temperature and pressure, contain equal numbers of molecules. This led to the following relations:

• 2 volumes of hydrogen + 1 volume of oxygen → 2 volumes of water

• 2 molecules of hydrogen + 1 molecule of oxygen → 2 molecules of water

Early Experiments to Characterize the Atom

Building on atomic theory, scientists sought to understand the structure of the atom through experimentation.

J.J. Thomson and the Discovery of the Electron

J.J. Thomson (1856–1940) performed experiments with cathode ray tubes (1898–1903). These tubes are partially evacuated glass tubes through which electrical currents are passed.

• High voltage produced a cathode ray from the negative electrode (cathode).

• The ray was repelled by the negative pole of an applied electric field, indicating it was made of negatively charged particles—later called electrons.

Thomson used electric and magnetic fields to measure the charge-to-mass ratio of the electron:

• Charge-to-mass ratio:

He proposed the Plum Pudding Model of the atom, where electrons are embedded in a cloud of positive charge.

Millikan's Oil Drop Experiment

Millikan's experiment allowed for the determination of the magnitude of the electron's charge. Combining this with Thomson's charge-to-mass ratio, the mass of the electron was calculated:

• Mass of electron:

Radioactivity

Radioactivity was first recognized in 1896 by Henri Becquerel and later studied by Marie Curie. Radioactive decay involves the emission of particles from unstable atomic nuclei.

• Alpha particles (α): Particles with a 2+ charge (about 7300 times the electron's mass); now known as helium nuclei.

• Beta particles (β): High-speed electrons.

Rutherford's Gold Foil Experiment and the Nuclear Model

Ernest Rutherford (1871–1937) tested the Plum Pudding Model in 1911 using the gold foil experiment. He observed that some alpha particles were deflected at large angles, which could only be explained by the existence of a dense, positively charged nucleus at the center of the atom.

• Nuclear Model: The atom consists of a small, dense nucleus containing protons (and later discovered, neutrons), with electrons moving around the nucleus at relatively large distances.

Comparison of Atomic Models

Example: Rutherford's experiment showed that most alpha particles passed through the gold foil, but a few were deflected, indicating a small, dense nucleus.

Additional info: The discovery of the neutron by James Chadwick in 1932 further refined the nuclear model, explaining the presence of neutral particles in the nucleus.