From Dalton to Chadwick

Around 400 BCE the Greek thinkers Leucippus and Democritus argued that if you cut a piece of matter in half, and then in half again, you can't keep going forever. Sooner or later you reach a piece that can't be cut. They called it atomos, Greek for "uncuttable".

It was an argument, not an experiment. Nobody could see an atom, weigh one, or test the idea.

Aristotle rejected atoms. He taught that everything is a mixture of four elements, earth, water, air and fire, each a pair of qualities: fire is hot and dry, air is hot and wet, water is cold and wet, and earth is cold and dry. His view dominated for about 2,000 years, and the idea that tiny particles make up everything was pushed aside.

It came back once chemists started weighing what went into a reaction and what came out. In the early 1800s John Dalton turned the atom into a scientific theory.

What was missing Evidence. Dalton's atoms came with masses that experiments could check.

By 1800 chemists were weighing carefully, and three patterns kept turning up.

Conservation of mass (Lavoisier, 1789)

In a chemical reaction, the total mass of the products equals the total mass of the reactants. Lavoisier showed it by weighing reactions in sealed vessels, so no gas could get in or out.

Definite proportions (Proust, 1799)

Also called the law of constant composition: a pure compound always contains the same elements in the same proportions by mass, wherever it comes from. Two samples of water:

Different amounts, same ratio: 7.92 g of oxygen for every gram of hydrogen.

Multiple proportions (Dalton, 1803)

When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers. Copper and bromine form two different substances:

2.52 g ÷ 1.26 g = 2, a whole-number ratio of 2 : 1. There are many more examples on the law of multiple proportions page.

The puzzle Why would nature care about whole numbers? Dalton's answer was atoms.

John Dalton was a Quaker schoolteacher in Manchester who kept daily weather records for decades and studied his own color blindness. He began working out the relative weights of atoms in 1803. His theory first reached print in 1807, in Thomas Thomson's chemistry textbook, and Dalton published it in full in A New System of Chemical Philosophy in 1808. It made five claims:

1. Matter is made of extremely small, indestructible particles called atoms.
2. All atoms of a given element are identical, in mass and in physical and chemical properties.
3. Atoms of different elements have different masses and different physical and chemical properties.
4. Atoms of different elements combine in simple whole-number ratios to form compounds, and a given compound always has the same ratio.
5. Atoms of one element can't be changed into atoms of another. A chemical reaction only rearranges atoms.

How atoms explain the three laws

• Conservation of mass. A reaction only rearranges atoms, and every atom keeps its mass, so the total can't change.
• Definite proportions. A compound always has the same ratio of atoms, and each kind of atom has a fixed mass, so the ratio by mass is fixed too.
• Multiple proportions. Two compounds of the same elements hold different whole-number ratios of atoms, like CO and CO₂, so for the same mass of carbon the oxygen masses come out 1 : 2.

What Dalton got wrong

With no way to find formulas, Dalton assumed the simplest one. He wrote water as HO, one hydrogen atom to one oxygen atom, and from about 7 g of oxygen per gram of hydrogen he gave oxygen an atomic weight of 7. Water is really H₂O. With two hydrogen atoms for every oxygen atom, an oxygen atom is about 2 × 7.9 ≈ 16 times as heavy as a hydrogen atom.

Still standing Statements 3 and 4, and statement 5 for chemical reactions. The others were revised as the experiments below came in; see Dalton's theory, checked.

A cathode ray tube is a glass tube with most of the air pumped out and a metal electrode sealed into each end. With a high voltage across it, a ray streams from the negative electrode, the cathode, toward the positive one, the anode. You can't see the ray itself, but it makes a coated screen glow where it hits.

Thomson showed three things:

• The ray bends toward a positively charged plate, so it is a stream of negatively charged particles.
• A magnet bends it too. By balancing the electric push against the magnetic one, he measured the particles' charge-to-mass ratio: −1.76 × 10⁸ C/g, in coulombs per gram.
• The particles were the same whatever metal the cathode was made of and whatever gas was left in the tube, so they must be part of every atom.

Their charge-to-mass ratio is about 1,800 times that of a hydrogen ion, so they are far lighter than any atom. Thomson called them corpuscles. We call them electrons.

What it changed Atoms have smaller parts, so Dalton's indestructible, uncuttable atom was wrong.

Thomson had measured charge divided by mass, but neither one on its own. Millikan set out to measure the charge.

He sprayed a fine mist of oil into a chamber above two horizontal metal plates, and a few drops drifted down through a small hole in the top plate. The drops picked up extra electrons, from the spraying and from X-rays that knocked electrons loose from the air. The top plate was positive, so it pulled upward on a negatively charged drop.

Watching a single drop through a microscope, Millikan adjusted the voltage until the upward electric pull balanced the drop's weight and it hung still. From that voltage, and from how fast the drop fell with the voltage off, he worked out the charge on the drop.

Every drop's charge was a whole-number multiple of the same amount, 1.602 × 10⁻¹⁹ C. That is the charge of one electron, −1.602 × 10⁻¹⁹ C. Dividing it by Thomson's charge-to-mass ratio gives the electron's mass:

mass = (−1.602 × 10⁻¹⁹ C) ÷ (−1.76 × 10⁸ C/g) = 9.10 × 10⁻²⁸ g

What it changed Charge comes in whole electrons. An electron has a definite charge and a tiny, definite mass, about 1/1,800 of a hydrogen atom's.

Atoms are electrically neutral, but electrons are negative. So every atom must also hold positive charge to cancel them out, and because electrons are so light, most of an atom's mass must be somewhere else too. The question was where.

In 1903 Hantaro Nagaoka pictured a large, heavy, positive center with electrons circling it in a ring, like Saturn and its rings.

In 1904 Thomson proposed that the positive charge is spread evenly through a sphere the size of the whole atom, with the electrons dotted through it like plums in a pudding, or chocolate chips in a cookie. This plum pudding model became the favorite picture of the atom.

The plum pudding model made a prediction that could be tested. If the positive charge is spread thin, nothing inside an atom is dense enough to stop a fast, heavy particle, so a beam of them should pass through atoms almost undisturbed.

What came next Rutherford had the fast, heavy particles to test it with: alpha particles from radioactive elements.

In 1896 Henri Becquerel found that uranium salts darkened a photographic plate wrapped in black paper, in the dark. The uranium was giving off radiation all by itself. Marie and Pierre Curie named this radioactivity, and in 1898 they discovered two new radioactive elements, polonium and radium.

The radiation turned out to be three different kinds, told apart by how they bend between charged plates:

• Alpha (α) particles are positive (2+) and heavy, about 7,300 times the mass of an electron. Each one is a helium nucleus.
• Beta (β) particles are fast electrons: negative and very light, so they bend the most.
• Gamma (γ) rays are high-energy light with no charge, so they don't bend at all.

In 1902 Ernest Rutherford and Frederick Soddy showed that as a radioactive element gives off radiation, its atoms turn into atoms of a different element.

What it changed Atoms of one element can become another through radioactive decay, a nuclear change. Chemical reactions still can't do it. And alpha particles gave Rutherford the probe he needed.

In Rutherford's lab, Hans Geiger and Ernest Marsden aimed a narrow beam of alpha particles from radium at a sheet of gold foil, which can be hammered thinner than any other metal. Around it was a screen coated with zinc sulfide that gave a tiny flash wherever an alpha particle landed. Sitting in the dark, they counted the flashes through a microscope.

If the plum pudding model were right, every alpha particle should go through with only the slightest deflection. Most did pass straight through. But a few were knocked off at large angles, and a very few bounced almost straight back. Rutherford said it was nearly as incredible as firing an artillery shell at tissue paper and having it come back at you.

In 1911 Rutherford explained it. Nearly all of an atom's mass and all of its positive charge are packed into a tiny, dense center, the nucleus. The nucleus is about 10⁻¹⁴ m across, while the atom is about 10⁻¹⁰ m, 10,000 times wider. The rest of the atom, where the electrons are, is mostly empty space. An alpha particle bounces back only when it heads almost straight at a nucleus, which is rare.

What it changed The plum pudding model was out. An atom is a tiny positive nucleus with electrons in the mostly empty space around it.

Positive particles had turned up before. In 1886 Eugen Goldstein saw rays in a discharge tube traveling the opposite way to cathode rays, and in 1898 Wilhelm Wien showed they were positively charged and far heavier than electrons.

A mass spectrometer sorts positive ions by mass. The ions fly through a magnetic field, which bends light ions more than heavy ones, so ions of different masses land in different places. The lightest positive particle is a hydrogen atom with its electron removed. It carries the same size of charge as an electron, but positive, and it is about 1,800 times as heavy (1,836 times, measured today). The masses of other atoms came out close to whole-number multiples of hydrogen's.

From 1917 to 1919 Rutherford fired alpha particles into nitrogen gas and detected hydrogen nuclei flying out: he had knocked them out of nitrogen nuclei. He concluded that the hydrogen nucleus is a building block of every nucleus, and in 1920 he named it the proton, from the Greek for "first".

What it changed A nucleus is built from protons, each with a charge of 1+. But the masses didn't add up: most nuclei weighed about twice what their protons could explain.

Studying the elements formed by radioactive decay, Frederick Soddy found substances that were chemically identical, so no chemical method could separate them, yet had different masses. In 1913 he named them isotopes, from the Greek for "same place", because they share one place in the periodic table. The name was suggested to him by Margaret Todd, a Scottish doctor and writer.

The same year Thomson found two kinds of neon atoms in his positive-ray tube, with masses of about 20 and 22. From 1919 Francis Aston's mass spectrograph turned up isotopes of element after element. Once the neutron was found in 1932 the reason was clear: isotopes of an element have the same number of protons and electrons, but different numbers of neutrons.

A mass spectrometer also measures how common each isotope is. The atomic mass on the periodic table is the weighted average of the isotope masses. For neon:

(0.9048)(19.992436) + (0.0027)(20.993843) + (0.0925)(21.991383) = 20.18 amu

What it changed Atoms of the same element are not all identical in mass, so Dalton's second statement had to be revised.

Mendeleev's periodic table (1869 to 1871) put the elements roughly in order of atomic mass, but a few neighbors had to be swapped to keep the families together. Tellurium (127.60) sits before iodine (126.90), and cobalt (58.93) before nickel (58.69).

Henry Moseley, who worked in Rutherford's group, fired electrons at samples of different elements and measured the X-rays they gave off. The square root of the X-ray frequency rose by the same step from each element to the next, and each step meant one more unit of positive charge in the nucleus. That count is the atomic number: the number of protons.

In order of atomic number no swaps are needed, and the gaps showed exactly which elements were still missing, such as numbers 43, 61, 72 and 75. Moseley was killed at Gallipoli in 1915, at the age of 27.

What it changed An element is defined by its atomic number, not its mass. Every atom of an element has the same number of protons.

Protons alone couldn't account for the mass of a nucleus. A helium nucleus has 2 protons but about 4 times the mass of a hydrogen nucleus, and for most atoms the protons explained only about half the mass. In 1920 Rutherford suggested that the nucleus also holds a neutral particle. But a particle with no charge is hard to catch: electric and magnetic fields don't bend it, and it leaves no trail in a detector.

In 1930 Walther Bothe and Herbert Becker found that beryllium struck by alpha particles from polonium gives off a very penetrating radiation, which they took to be gamma rays. Early in 1932 Irène Joliot-Curie and Frédéric Joliot found that this radiation knocks protons out of paraffin wax at high speed.

Chadwick showed that gamma rays couldn't do that. He measured how hard the radiation struck protons and nitrogen atoms, and from those collisions worked out that it was a stream of neutral particles about as heavy as a proton: neutrons. He won the 1935 Nobel Prize in Physics for the discovery.

What it changed A nucleus holds protons and neutrons. Isotopes of an element have the same number of protons but different numbers of neutrons.

An atom is a tiny, dense nucleus of protons and neutrons, together called nucleons, with electrons around it. The nucleus is about 10⁻¹⁴ m across and holds almost all of the mass; the electrons take up a region about 10⁻¹⁰ m across. A neutral atom has the same number of protons and electrons.

Masses this small are given in atomic mass units. 1 amu is defined as 1/12 of the mass of one carbon-12 atom, which is 1.6605 × 10⁻²⁴ g. The same unit is also called the dalton (Da), after John Dalton, and the unified atomic mass unit (u).

The atomic number (Z) is the number of protons, and it decides the element. The mass number (A) is the number of protons plus neutrons, so neutrons = A − Z. Carbon-12 has 6 protons and 6 neutrons; carbon-14 has 6 protons and 8 neutrons.

The picture of electrons in energy levels and clouds came later, starting with Niels Bohr in 1913 and quantum mechanics in the 1920s.

In one line Protons decide the element, neutrons decide the isotope, and electrons decide the chemistry.