2.1 Subatomic Particles and Early Atomic Models

Section 1.2 described Dalton's atom: a tiny, hard, indivisible ball, different for each element. That picture explained the laws of chemical combination beautifully — and it was wrong in one important way. By the end of the nineteenth century, experiments on electricity passing through gases, on electrolysis and on radioactivity had shown that atoms can lose or gain tiny charged pieces. Atoms have an inside.

This section tells the story of how that inside was mapped: the electron in 1897, the nucleus in 1911, the neutron in 1932. The experiments are worth learning properly, not only because Board papers ask about them, but because the reasoning — measure a deflection, deduce a charge-to-mass ratio; fire particles at a target, deduce what is inside — is exactly how physicists and chemists still probe matter, from hospital mass spectrometers to the Large Hadron Collider. We finish with the vocabulary every later chapter uses: atomic number, mass number, isotopes and isobars.

Cathode Rays and the Discovery of the Electron

Gases are normally very poor conductors of electricity. In the 1850s–1870s, Michael Faraday, William Crookes and others sealed two metal electrodes into a glass tube, pumped out most of the gas and applied a very high voltage (several thousand volts). At very low pressure the gas stops glowing and a beam travels from the negative electrode (the cathode) towards the positive electrode (the anode). If the anode has a hole in it, the beam passes through and makes the glass — or a coating of zinc sulphide placed behind the anode — glow where it strikes. The old television sets and computer monitors that used picture tubes were exactly such tubes, with the beam steered across a fluorescent screen.

Careful experiments established the following properties of cathode rays:

• they start at the cathode and travel in straight lines (an object in their path casts a sharp shadow);
• they are deflected towards a positively charged plate, so they carry negative charge; a magnet also deflects them, as it deflects any moving charge;
• they are invisible themselves, and are detected by the glow they produce on fluorescent materials;
• their properties do not depend on the metal of the cathode or on the gas in the tube.

The last point is the crucial one. Whatever the electrodes and whatever the gas, the same particles come out — so they must be part of every atom.

Thomson's charge-to-mass ratio

In 1897 J. J. Thomson measured how much a beam of cathode rays is bent by electric and magnetic fields applied at right angles to the beam and to each other. The deflection of a charged particle from its path is larger when

• its charge is larger (the force on it is bigger);
• its mass is smaller (the same force produces a bigger acceleration);
• the field is stronger.

Thomson's method had two stages. First he applied both fields and adjusted them until the beam went straight on: the electric force $eE$ and the magnetic force $evB$ then cancel, which gives the speed of the electrons, $v=E/B$. Then he switched the magnetic field off and measured how far the electric field alone moved the spot. Because the deflection depends only on the ratio of charge to mass, the experiment gives $e/m_{e}$ — not $e$ and $m_{e}$ separately. The modern value is

This ratio is about $1800$ times larger than that of the lightest ion known at the time (the hydrogen ion in electrolysis). Either the electron carries a huge charge, or it is very light. Millikan's experiment settled which.

Millikan's oil-drop experiment: the charge of the electron

Between 1906 and 1914 R. A. Millikan sprayed a fine mist of oil droplets into a chamber between two horizontal metal plates. X-rays passing through the chamber knocked electrons out of air molecules, and some of these electrons stuck to the droplets. Watching one droplet through a microscope, Millikan first timed its steady fall under gravity (which gives its size and mass), then switched on an electric field across the plates and measured how much the field slowed, stopped or reversed the fall. From this he could calculate the charge on the droplet. Repeating the measurement on many drops, he found a striking pattern.

Protons, Neutrons and the Particles of the Atom

Atoms are electrically neutral, so if they contain negative electrons they must also contain an equal amount of positive charge. In 1886 E. Goldstein used a discharge tube with a perforated cathode and saw a glow behind the cathode: rays moving the opposite way, towards and through the cathode. These canal rays (or anode rays) are positively charged ions formed when cathode rays knock electrons out of gas atoms. Unlike cathode rays,

• their charge-to-mass ratio depends on the gas in the tube (they are the positive ions of that gas);
• the charge on them is always a whole-number multiple of $e$;
• they are deflected the opposite way to cathode rays by electric and magnetic fields.

The lightest positive ion, with the largest $e/m$, came from hydrogen: a hydrogen atom that has lost its one electron. This particle, with charge $+e$ and mass about $1836$ times the electron's, was named the proton (Rutherford, around 1919). Atomic masses, however, were too large to be accounted for by protons alone — helium has two protons but a mass of about four units. In 1932 James Chadwick bombarded a thin sheet of beryllium with α-particles and detected electrically neutral particles with a mass slightly larger than the proton's — the neutron:

The Thomson and Rutherford Models of the Atom

Knowing the particles is not the same as knowing how they are arranged. In 1898 Thomson proposed that an atom is a sphere of radius about $10^{-10}\ \mathrm{m}$ in which the positive charge is spread out uniformly, with the electrons embedded in it so as to give the most stable arrangement — like seeds in a watermelon (or plums in a pudding, hence the name "plum pudding model"). In this model the mass of the atom is also spread uniformly through the sphere. It explained why atoms are neutral, but nothing more, and it did not survive its first real test.

Rutherford's α-particle scattering experiment

In 1909–1911 Ernest Rutherford, with Hans Geiger and Ernest Marsden, fired a narrow beam of fast α-particles (helium nuclei, $\mathrm{He^{2+}}$, emitted by a radioactive source) at a very thin gold foil, about $100\ \mathrm{nm}$ thick. A circular screen coated with zinc sulphide surrounded the foil, and every α-particle that struck it produced a tiny flash of light that could be counted through a microscope. If Thomson were right, the positive charge and mass of each gold atom would be too thinly spread to turn a heavy, fast α-particle by more than a fraction of a degree. The results were:

1. most α-particles passed straight through the foil undeflected;
2. a small fraction were deflected through small angles;
3. a very few (about $1$ in $20\,000$) were deflected through large angles, some bouncing almost straight back.

Rutherford said it was "as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you". Something very small, very heavy and strongly positive must be inside each atom.

The three observations follow at once. An α-particle passes through mostly empty space, so it is usually not deflected; one that passes fairly close to a nucleus is pushed sideways by its positive charge and turns slightly; only an α-particle heading almost straight at a nucleus — a rare event, because the nucleus is a tiny target — is turned back. Interactive 2 lets you fire α-particles at a nucleus and compare the two models.

What the Rutherford model could not explain

Atomic Number, Mass Number, Isotopes and Isobars

Rutherford's experiments, followed by Moseley's X-ray measurements in 1913, showed that the positive charge on the nucleus — the number of protons — is what makes an element what it is. Every carbon atom has six protons; every atom with six protons is carbon.

Chemical properties are decided by the number of electrons, which equals the number of protons; neutrons have almost no effect. So isotopes of an element have the same chemical properties (only the rates of their reactions differ slightly, an effect that is noticeable for the light hydrogen isotopes), while isobars, being different elements, behave quite differently. Isotopes are invaluable: carbon-14 dates archaeological remains, cobalt-60 is used in cancer radiotherapy, iodine-131 treats thyroid disease, and the ratio of $\mathrm{{}^{18}O}$ to $\mathrm{{}^{16}O}$ in polar ice records past climates. Section 1.2 showed how the abundances of isotopes give an element's average atomic mass. (A third, less common, term: atoms with the same number of neutrons, such as $\mathrm{{}^{14}_{6}C}$, $\mathrm{{}^{15}_{7}N}$ and $\mathrm{{}^{16}_{8}O}$, are called isotones.)

Exercises