A chronological list of the moments that actually built the ideas this course teaches — not a complete history of chemistry, but the handful of discoveries a first-semester student keeps running into by name. Each entry links back to the glossary term, module, or lesson where that idea shows up in this course.
Ancient world
c. 440 BCE — Democritus and the first atomic idea
The Greek philosopher Democritus proposed that matter is made of indivisible particles he called atomos — "uncuttable." It was a philosophical guess with no experimental support behind it, and it would sit largely untested for over two thousand years, but the name survived all the way to the modern atom.
The birth of modern chemistry
1661 — Robert Boyle separates chemistry from alchemy
Boyle's The Sceptical Chymist argued that matter is built from combinations of primary particles and pushed for chemistry to define its terms through experiment rather than alchemical tradition — one of the first arguments for what an element actually is, in the modern sense.
1774 — Joseph Priestley isolates oxygen
By heating mercury oxide and collecting the gas it released, Priestley isolated oxygen and noted that it made a candle burn brighter and kept a mouse alive longer than ordinary air — without yet understanding why. Carl Wilhelm Scheele had independently produced the same gas slightly earlier but published later; Priestley usually gets the credit for the first published discovery.
1789 — Antoine Lavoisier and the conservation of mass
Lavoisier's careful before-and-after weighing of reactions in sealed vessels established that mass is neither created nor destroyed in a chemical reaction — the law of conservation of mass that every balanced equation in Module 4 depends on. He also named oxygen (from Priestley's gas) and hydrogen, and helped kill off the old phlogiston theory of combustion in the process. He is often called the father of modern chemistry for exactly this shift toward quantitative, mass-based reasoning.
1803 — John Dalton's atomic theory
Dalton proposed that each element consists of identical, indivisible atoms with a characteristic mass, that atoms combine in fixed whole-number ratios to form compounds, and that chemical reactions simply rearrange atoms rather than creating or destroying them — the direct ancestor of Module 2's atomic model, more than two thousand years after Democritus's guess and finally backed by measurement.
1811 — Avogadro's hypothesis
Amedeo Avogadro proposed that equal volumes of any gas, at the same temperature and pressure, contain equal numbers of particles — regardless of what the gas actually is. The idea was mostly ignored for decades until Stanislao Cannizzaro revived it in the 1860s to help settle a chaotic mess of inconsistent atomic weights across the field. The mole and the Avogadro constant at the center of Module 3 both carry his name.
Ordering the elements
1869 — Mendeleev's periodic table
Dmitri Mendeleev arranged the 63 elements known at the time by atomic weight and recurring chemical behavior, leaving deliberate gaps for elements he predicted must exist but hadn't yet been discovered — and correctly forecast the properties of gallium, scandium, and germanium years before each was found. That predictive power, more than the arrangement itself, is what convinced the field to take his periodic table seriously.
1913 — Henry Moseley reorders the table by atomic number
Mendeleev's table was ordered by atomic weight, which put a few elements in the wrong sequence. Moseley's X-ray experiments showed that atomic number — the actual proton count — was the true organizing property, resolving those inconsistencies and putting the modern periodic table on solid physical footing shortly before he was killed in World War I at age 27.
Inside the atom
1897 — J. J. Thomson discovers the electron
Using a cathode-ray tube, Thomson showed that atoms contain much smaller, negatively charged particles — the electron — proving that Dalton's "indivisible" atom actually has internal structure. His proposed "plum pudding" model (electrons scattered through a positive sphere) would only stand for about a decade.
1898 — Marie and Pierre Curie and radioactivity
Building on Henri Becquerel's discovery that uranium salts emitted a mysterious kind of radiation, Marie Curie coined the term "radioactivity," and she and Pierre Curie isolated two new elements responsible for it — polonium (named for her native Poland) and radium. Marie Curie remains the only person to win Nobel Prizes in two different sciences (Physics, 1903; Chemistry, 1911). Their work is the direct ancestor of the alpha decay and beta decay this course's glossary covers past Module 2's isotope notation.
1909 — Millikan's oil-drop experiment
Robert Millikan suspended charged oil droplets between electric plates and measured the electron's charge with startling precision — a single, fixed value that every electron carries, confirming that charge itself comes in discrete units rather than a continuous range.
1911 — Rutherford's gold foil experiment finds the nucleus
Ernest Rutherford's team fired alpha particles at thin gold foil expecting them to pass straight through Thomson's diffuse "plum pudding" atom. Most did — but a small fraction bounced almost straight back, which Rutherford called "as if you fired a fifteen-inch shell at tissue paper and it came back and hit you." The only explanation was a tiny, dense, positively charged nucleus at the atom's center with mostly empty space around it — the model Module 2 still starts from.
1913 — Bohr's model of the hydrogen atom
Niels Bohr proposed that electrons orbit the nucleus only at specific, quantized energy levels, and that light is absorbed or emitted only when an electron jumps between them — correctly predicting hydrogen's line spectrum for the first time. The model's literal orbits didn't survive later quantum mechanics, but the core idea — electrons occupy discrete energy levels, not just any energy — is exactly what Module 10 still teaches.
1932 — Chadwick discovers the neutron
James Chadwick identified a third subatomic particle with no electric charge and roughly a proton's mass, resolving a long-standing puzzle about why an atom's mass didn't match its proton count alone. The neutron completed the proton-neutron-electron picture Module 2 opens with, and made isotopes finally make physical sense.
Bonds, shape, and structure
1916 — G. N. Lewis and the shared electron pair
Gilbert Lewis proposed that a covalent bond is a pair of electrons shared between two atoms, and introduced the dot-structure diagrams — Lewis structures — that Module 12 is built entirely around. It replaced a vaguer, more mysterious idea of bonding with something a student could actually draw and reason from.
1930s — Linus Pauling and modern bonding theory
Across a series of papers (collected in his 1939 book The Nature of the Chemical Bond), Linus Pauling worked out the quantitative electronegativity scale, hybrid-orbital bonding, and resonance — the theoretical backbone of Modules 12–13. Pauling remains the only person to win two unshared Nobel Prizes (Chemistry, 1954; Peace, 1962).
1953 — Watson, Crick, and Franklin determine DNA's structure
James Watson and Francis Crick proposed the double-helix structure of DNA, built directly on Rosalind Franklin's X-ray diffraction data (most famously "Photo 51") along with Maurice Wilkins's work — a structure that immediately explained how genetic information could be copied. It's biology's landmark moment built entirely from chemical structure reasoning: base pairing, hydrogen bonding, and molecular geometry, the same tools Module 12 introduces for far smaller molecules.
Splitting and rearranging the nucleus
1938 — Hahn, Strassmann, and Meitner discover nuclear fission
Otto Hahn and Fritz Strassmann found that bombarding uranium with neutrons produced much lighter elements — a result that didn't make sense until Lise Meitner and Otto Frisch (working from exile after fleeing Nazi Germany) correctly interpreted it as the uranium nucleus splitting in two, releasing enormous energy in the process. Nuclear fission is the reaction this course's glossary connects to both nuclear power and, more grimly, the weapons developed within the following decade.
Chemistry at the molecular-materials frontier
1985 — Kroto, Curl, and Smalley discover buckminsterfullerene
While studying carbon clusters formed by vaporizing graphite with a laser, the team found a remarkably stable 60-carbon molecule shaped like a soccer ball — a third form of pure carbon (after diamond and graphite), named buckminsterfullerene after the architect Buckminster Fuller's geodesic domes. It kicked off an entire field of carbon nanomaterials and won the 1996 Nobel Prize in Chemistry.
2004 — Geim and Novoselov isolate graphene
Using ordinary adhesive tape to repeatedly peel layers off a piece of graphite, Andre Geim and Konstantin Novoselov isolated graphene — a single, atom-thick sheet of carbon — and measured its extraordinary strength and conductivity. The "tape method" is one of modern chemistry's best examples of a Nobel-winning result (2010) built from a strikingly simple technique.
2016 — The periodic table's seventh row is completed
IUPAC officially named the last four elements of period 7 — nihonium (113), moscovium (115), tennessine (117), and oganesson (118) — completing the seventh row of the periodic table exactly 147 years after Mendeleev's first version. Every one of these superheavy elements is synthetic, unstable, and exists for only fractions of a second before decaying, made a few atoms at a time in particle accelerators rather than found in nature.
Why this list stops where it does
Chemistry didn't stop being made in 2016 — it's an active research field with new results published every week — but the entries above are the ones a first-semester student runs into by name often enough that it helps to know where they came from. For the frontier of current research, Readings and Resources points toward primary sources rather than a fixed list, since a fixed list of "the latest" chemistry would be out of date the moment it was written.