Atoms, Molecules, and Ions

This is Unit Two of Chemistry, Part One: the atom and its parts. The unit opens with three laws established before anyone had seen a subatomic particle, from mass measurements alone. The third of them, multiple proportions, is the one worth slowing down for — it compares masses across two different compounds, not atoms within one, and confusing those two misreads the law entirely. Then three experiments. Thomson found a particle far lighter than any atom. Millikan found the smallest charge, without ever isolating a single electron — he inferred it as the common factor across many droplets. Rutherford expected no deflection at all and got a rare, sharp bounce, which forced the nuclear model. Then the three particles themselves, and a comparison entry the book flags as load-bearing. The intuition to correct is that the heavy particle is the charged one outside. It is the reverse: the light particle is outside, and both heavy ones sit together in the core. The part closes on the two ways an atom varies. Change the neutrons and you have an isotope. Change the electrons and you have an ion. The protons never change in either case. By the end you will be able to do five things. One: say what multiple proportions compares. Two: say what each experiment established. Three: place, charge and weigh each particle. Four: get neutron count from two numbers. Five: tell an isotope from an ion.

Four entries covering what careful weighing established on its own, decades before any particle was found. Dalton's theory came first, and two mass laws supported it. The procedure at the end is where the third law is actually applied, and it has a step people skip.

1807.

Chem 0181 Dalton's Atomic Theory

Dalton's atomic theory of 1807 holds that matter is made of indivisible atoms, each element's atoms identical in mass and unchanged by reactions — copper keeps its identity through every chemical change. Dalton called atoms indivisible, and later discoveries added subatomic particles without discarding his other four postulates.

Always the same ratio.

Chem 0184 The Law of Definite Proportions

The law of definite proportions states that every sample of a pure compound contains the same elements in the same fixed mass ratio — isooctane shows a carbon-to-hydrogen mass ratio of five point three three to one wherever it comes from. Sharing a mass ratio does not prove two samples are the same compound; other compounds can match it.

Two compounds, compared.

Chem 0187 The Law of Multiple Proportions

When two elements form more than one compound, the second element's combining masses per fixed mass of the first form a small whole-number ratio — chlorine combines with copper in one ratio in one compound and twice that in another. The law compares masses across two compounds, not atoms within one.

Applying it.

Chem 0189 Applying the Law of Multiple Proportions

The procedure fixes one element's mass across two compounds, then reduces the other element's two masses to a small whole-number ratio. Dividing oxygen by carbon in each of two compounds gives about one point three three and about two point six seven; dividing the larger by the smaller gives almost exactly two. That confirms the law and shows the second compound carries twice the oxygen per unit of carbon — carbon monoxide and carbon dioxide fit the pattern. Stopping at the two separate ratios looks like progress and proves nothing; the law is confirmed only once those ratios are compared to each other.

Four entries. Three experiments across sixteen years and the model that came out of the last one. Each experiment's warning is about how narrow its result actually was — worth holding onto, because the chain reads as one smooth discovery otherwise.

1897.

Chem 0193 J. J. Thomson and the Discovery of the Electron (1897)

In 1897 J. J. Thomson used a cathode ray tube to show cathode rays were negatively charged particles lighter than any atom, calculating a charge-to-mass ratio that came out identical whichever metal formed the electrodes. The particle was not a smaller atom — being far lighter than any atom is what marked it a new subatomic constituent.

1913.

Chem 0196 Millikan's Oil-Drop Experiment (1913)

In his 1913 published results, Robert Millikan showed each oil droplet's charge was a whole-number multiple of one fixed smallest charge, balancing gravity against an adjustable electric field to read each drop. He never isolated a single electron — he inferred the fundamental charge as the smallest factor shared across many droplets.

1911.

Chem 0199 Rutherford's Gold Foil Experiment (1911)

Rutherford, with Geiger and Marsden, fired alpha particles at gold foil: most passed straight through and a few bounced back, meaning they had struck something small, dense and positively charged hidden in mostly empty space. He had expected no deflection at all under the earlier model — the rare bounce is what forced a new one.

The model.

Chem 0202 The Nuclear Model of the Atom

Rutherford's nuclear model places a small, dense, positively charged nucleus at an atom's centre, surrounded by mostly empty space holding the electrons — scale the nucleus to a blueberry and the atom is about a football stadium. Mostly empty does not mean unstable; the nucleus's concentrated mass and charge govern how the atom behaves chemically.

Three entries, and the last is a comparison the book calls out as what most other atomic reasoning depends on. Three properties, three particles — location, charge and relative mass — and one intuition that runs backwards.

Positive, inside.

Chem 0205 Proton

A proton is a positively charged subatomic particle in the nucleus, with a mass far greater than an electron's — every hydrogen nucleus is a single proton. Its charge is fixed at one positive unit; only how many protons are present varies, never the charge each one carries.

Neutral, inside.

Chem 0208 Neutron

A neutron is an uncharged subatomic particle in the nucleus, close in mass to a proton, confirmed by Chadwick in 1932 — two atoms of one element can carry different numbers of them, changing mass without changing chemical behaviour. Uncharged does not mean massless; neutrons contribute to mass fully as much as protons.

All three, compared.

Chem 0211 Comparing the Three Subatomic Particles

Protons and neutrons both sit in the nucleus with roughly the same mass as each other; protons carry one positive unit of charge and neutrons none. Electrons sit outside, occupying nearly all of the atom's volume while carrying almost none of its mass — a proton or neutron outweighs an electron more than a thousand times over. A neutral atom always holds as many electrons as protons, which is what keeps it balanced. The common error is assuming the heavier particle must be the charged one outside; it is the reverse — the light particle is outside and both heavy ones sit together at the centre.

Three entries on how to write an element and how to count what is in it. Two numbers do the work here: one fixes which element you have, and the other, with a subtraction, gives the neutrons.

Writing it.

Chem 0214 Chemical Symbols

A chemical symbol abbreviates an element's name, first letter capitalised and any remaining letters lowercase — cobalt is C-o, while the different capitalisation C-O names carbon monoxide instead. Many symbols come from Latin rather than English: sodium and lead take their symbols from older names.

Which element.

Chem 0217 Atomic Number

The atomic number is the number of protons in an atom's nucleus, and it alone defines which element the atom is — any atom with six protons is carbon, whatever its neutrons or electrons. Neutron count never sets identity; atoms with the same protons and different neutrons are the same element.

How heavy.

Chem 0219 Mass Number

The mass number is the total of protons plus neutrons in the nucleus, and subtracting the atomic number gives the neutron count — atomic number seventeen and mass number thirty-five means eighteen neutrons. It is a count, not a mass in grams: it ignores electrons and the slight difference between proton and neutron mass.

Three entries on the distinction the rest of the unit rests on. Two of an element's three particle counts can change, and each change has its own name. Protons are the one thing that never varies.

Neutrons differ.

Chem 0221 Isotope

Isotopes are atoms of the same element sharing the same number of protons but differing in neutrons — carbon-twelve and carbon-fourteen both carry six protons, and the second carries two more neutrons. Isotope differences never change chemical identity or ordinary chemical behaviour.

Electrons differ.

Chem 0222 Ion

An ion is an atom or molecule whose electron count no longer equals its proton count, giving it a net charge — a neutral sodium atom with eleven protons and eleven electrons becomes a positive ion after losing one. An ion's proton count never changes, only its electrons; isotopes instead vary in neutrons and never in charge.

Writing an isotope.

Chem 0225 Isotope Notation (Nuclear Symbol)

Isotope notation writes an element's mass number as a left superscript before its symbol, naming that specific isotope — magnesium-twenty-four has twelve protons like every magnesium atom and exactly twelve neutrons. The atomic number, sometimes shown as a left subscript, is often dropped, since the element symbol already fixes it.

Two entries closing the part, and the second carries a memory aid rather than a correction — which is unusual for this book and a signal that these two get swapped more than most.

Lost electrons.

Chem 0228 Cation

A cation is a positively charged ion, formed when an atom or molecule loses one or more electrons — a calcium atom losing two becomes a two-plus cation, with two more protons than electrons. The name refers to the missing electrons, not to extra protons; proton count never changes when an ion forms.

Gained electrons.

Chem 0229 Anion

An anion is a negatively charged ion, formed when an atom or molecule gains one or more electrons — a bromine atom gaining one carries one more electron than it has protons. The two names are easy to swap under time pressure, and the book's aid is that an anion's gained electrons pull toward the negative sign.

Five things to carry out of Part One. The law of multiple proportions compares masses across two compounds, not atoms within one. Fix one element's mass in both, take the other element's two masses, and reduce those two ratios against each other to a small whole number. Stopping at the separate ratios proves nothing — the ratio of ratios is the whole test. Each experiment established one thing. Thomson found a negatively charged particle far lighter than any atom. Millikan found the smallest unit of charge, inferred as the common factor across many droplets rather than isolated. Rutherford's rare, sharp deflections forced the nuclear model on him. Three particles, three properties. Protons and neutrons sit in the nucleus with roughly equal mass; protons carry one positive charge, neutrons none. Electrons sit outside, filling nearly all the volume and carrying almost none of the mass. The light particle is the one outside. Neutron count is mass number minus atomic number. The atomic number counts protons and fixes which element you have. The mass number counts protons plus neutrons and is a count rather than a mass in grams. And an isotope differs from an ion by which count changed. Different neutrons gives an isotope, with the same chemical behaviour and no charge. Different electrons gives an ion, with a net charge — positive when electrons are lost, negative when they are gained. Protons never change in either case. Part Two turns to measuring atomic mass, and to the formulas that describe compounds.