Atoms and elements

Through most of the 1800s the atom was pictured as a solid ball with nothing inside it. Then, in 1897, the English physicist J. J. Thomson found a particle far lighter than the lightest atom: the electron. The uncuttable atom had parts.

Where the rest of it sat was settled in Ernest Rutherford’s laboratory in Manchester. In 1909 his team fired a stream of fast, heavy particles at gold foil far thinner than paper. Nearly all of them sailed straight through, as if nothing were there. But a very few, about one in twenty thousand, bounced back. Rutherford published his conclusion in 1911. Atoms are mostly empty space, with almost all their mass packed into a tiny, dense centre, now called the nucleus. A direct hit on it was rare, which is why so few particles came back.

The nucleus is itself made of smaller particles, packed tightly together, and there are two kinds. A proton carries a positive electric charge. A neutron is almost exactly as heavy but carries no charge at all, which is where its name comes from: it is neutral. The electron that Thomson found is charged too, negatively, but it lives outside the nucleus.

Like charges push each other apart, so a crowd of positive protons squeezed into so small a space ought to fly apart. It holds because, at very short range, a far stronger pull called the strong force binds protons and neutrons to one another. Neutrons add to that pull without adding any push. Hydrogen, whose nucleus is one lone proton, is the only atom that can do without them. Every heavier nucleus needs neutrons to hold together.

Around the nucleus, and far from it, move the atom’s electrons. Each carries a negative charge exactly as large as a proton’s positive one, and the attraction between the two is what holds an atom together.

Electrons are astonishingly light. A proton has the mass of about two thousand of them, 1,836 to be exact, so the nucleus holds more than 99.9 per cent of an atom’s mass while taking up almost none of its room. If an atom were as wide as a football stadium, its nucleus would be a small pea on the centre spot.

An atom has as many electrons as protons, so the charges cancel and its overall charge is zero. Knock an electron off, or add a spare, and the balance is gone: the charged atom that results is called an ion. Electrons are also what atoms share or hand over when they join.

What makes one atom gold and another oxygen? Only the number of protons in its nucleus. Every hydrogen atom has 1, every carbon atom 6, every oxygen atom 8, every gold atom 79. That count is the atomic number, and a substance built from atoms that all share one atomic number is an element.

Change the count and you change the element. Take one proton out of a gold nucleus and what is left is platinum. This is why the alchemists, for all their skill, never turned lead into gold. Lead has 82 protons. Everything they could do to it, heating, dissolving, mixing, was a chemical reaction, and a chemical reaction rearranges the electrons on the outside of atoms. It leaves every nucleus exactly as it was.

How many elements are there? About ninety occur naturally on Earth, from hydrogen up to uranium, and physicists have made the rest in laboratories, bringing the list to 118. Everything you have ever touched, wood, water, skin, steel, is some combination of those ninety or so kinds of atom, much as every book in English is a combination of twenty-six letters.

Each element has a symbol of one or two letters, written the same way by chemists in every country. Many are obvious: H for hydrogen, O for oxygen, C for carbon. Others look odd until you know that they come from the element’s Latin name. Iron is Fe, from ferrum. Gold is Au, from aurum. Sodium is Na, from natrium. The Swedish chemist Jöns Jacob Berzelius proposed the system in 1813.

Atoms of one element always have the same number of protons, but they need not have the same number of neutrons. Versions of an element that differ only in their neutrons are called isotopes, and each is named by adding its protons and neutrons together. Every carbon atom has 6 protons. Most have 6 neutrons as well, and that isotope is carbon-12. A few have 7, making carbon-13, and a very few have 8, making carbon-14.

About 99 per cent of the carbon in nature is carbon-12. About 1 per cent is carbon-13, and carbon-14 is the merest trace, roughly one atom in a trillion. Isotopes have the same protons and so the same electrons, which means they behave alike in chemical reactions. What differs is their weight and, sometimes, whether the nucleus is stable.

Where did the ninety elements come from? The two simplest, hydrogen and helium, date from the first minutes after the Big Bang. Nearly everything heavier was made later, inside a star.

At the centre of a star, heat and pressure are so great that light nuclei are forced together into heavier ones. This is fusion, and it releases the energy that makes the star shine. The Sun fuses hydrogen into helium. Bigger stars go further, building carbon, oxygen and heavier elements up to iron. When such a giant runs out of fuel it collapses and blows apart in an explosion called a supernova, which scatters its elements across space, and from that debris new stars and planets form. Elements heavier than iron are built in other ways, some slowly inside ageing stars and some, such as gold, in the collision of two dead stars.

Add up all the ordinary matter in the universe and the result is lopsided. About 74 per cent of it, by mass, is hydrogen, the simplest atom there is: one proton and one electron. About 24 per cent is helium, the next simplest. Every other element put together, all the carbon, oxygen, iron and gold, comes to about 2 per cent. Stars have been fusing hydrogen for more than thirteen billion years and have barely dented the supply.

Helium is so light that Earth’s gravity cannot hold on to it, so it is scarce here, and it was found first in the Sun. In 1868 astronomers studying sunlight saw a yellow line that matched no element then known. The new element was named after helios, the Greek word for the Sun. Nobody isolated helium on Earth until 1895.

Closer to home the recipe changes completely. Earth’s crust, the rocky outer layer we live on, is nearly half oxygen by mass, about 46 per cent, and more than a quarter silicon, about 28 per cent. The two are locked together: sand and most rock are silicon bonded to oxygen. Almost none of that oxygen is the gas we breathe.

Your body follows a different recipe again. About 65 per cent of your mass is oxygen, because you are mostly water and oxygen is the heavy part of every water molecule. Carbon comes next at about 18 per cent, then hydrogen at 10 and nitrogen at 3. Just these four elements make up 96 per cent of you. Hydrogen, the lightest element of all, supplies the most atoms but little of the weight.

Not every nucleus lasts. Some combinations of protons and neutrons are unstable, and sooner or later such a nucleus throws out a particle or a burst of energy and settles into a steadier form. This spontaneous breakdown is called radioactivity. Henri Becquerel came upon it in Paris in 1896, when uranium salts fogged a wrapped photographic plate, and Marie Curie gave it its name.

Nothing triggers a decay and nothing can hurry it. Nobody can say when one particular nucleus will go. Yet a large crowd of them keeps a strict rule: in a fixed length of time, half of them decay. That time is the isotope’s half-life. After one half-life, 50 per cent of the atoms are left. After two, 25 per cent. After three, 12.5 per cent. Half-lives run from a fraction of a second to billions of years.

What comes out of a decaying nucleus? Early experimenters found three kinds of radiation and, not yet knowing what they were, named them after the first three letters of the Greek alphabet.

Alpha radiation is a heavy particle, two protons and two neutrons, the same as a helium nucleus. It hits hard but is stopped by a sheet of paper or the outer layer of your skin. Beta radiation is a fast electron thrown out of the nucleus. It gets through paper, and it takes a few millimetres of aluminium to stop it. Gamma radiation is not a particle at all but a wave, from the same family as light and X-rays, with far more energy than light. It passes through the body, and only thick lead or concrete blocks it. The lighter the radiation and the less charge it carries, the further it travels.

Carbon-14 is radioactive, with a half-life of 5,730 years, and more of it is made all the time, high in the air, where rays from space turn nitrogen atoms into carbon-14. Plants take it in along with ordinary carbon, and animals eat the plants. While a plant or animal lives it keeps absorbing fresh carbon, so the small share of carbon-14 in its body stays level even though those atoms are steadily decaying.

Death stops the supply. From that moment the carbon-14 only decays, halving every 5,730 years. Measure how much is left in a piece of bone, wood or cloth and you can work out how long ago it stopped living. The method is called carbon dating, and the American chemist Willard Libby worked it out in the 1940s. It reaches back about fifty thousand years. Beyond that, too little carbon-14 remains to measure.

Radiation damages living cells, and the same reach that makes it dangerous makes it useful. In a hospital, a patient can be given a tiny dose of a radioactive substance called a tracer, which travels to a chosen organ. The tracers used give off gamma radiation, because gamma rays escape the body and reach a camera outside, which builds a picture of blood flow or of a hidden tumour. Alpha radiation would be useless here: it would be stopped inside the body and never reach the camera. Stronger beams, carefully aimed, are used to kill cancer cells.

People who work with radioactive sources protect themselves in three ways. They put shielding between themselves and the source, they keep their time near it short, and they keep their distance, because radiation spreads out as it travels and weakens quickly.