Journey Inside the Atom

Class 8 explained that all matter is built from tiny particles. If a particle belongs to an element, it is an atom, and atoms are so small that a stack of about a million of them, side by side, would only be as thick as a single sheet of paper. No microscope, then or now, can simply look inside one and see what's there directly. Everything scientists know about the inside of an atom had to be worked out indirectly, by firing things at atoms and carefully reasoning backward from what happened.

In 1808, John Dalton proposed the first genuinely scientific theory of the atom: that all matter is made of tiny, indivisible particles called atoms, which cannot be created, destroyed, or split into anything smaller. Atoms of the same element, Dalton said, are identical to each other in mass and properties, while atoms of different elements differ. For its time, this was a huge step forward, since it finally gave scientists a solid, testable idea of what matter was built from. But it left an obvious question unanswered: if atoms are the smallest possible thing, why do different elements behave so differently from each other, and what makes one atom different from another on the inside?

In 1897, J. J. Thomson passed a high voltage through a gas at very low pressure inside a glass tube and studied the rays that streamed from the negative electrode to the positive one. Using electric and magnetic fields, he showed these rays were actually a stream of tiny, negatively charged particles, far lighter than any atom, later named electrons. This was the first evidence that atoms are not truly indivisible: they contain even smaller pieces inside them. But this raised a puzzle: atoms overall have no charge, so if electrons carry negative charge, where is the positive charge that must be balancing them out?

Thomson's own answer was that the atom is a sphere of positive charge with electrons scattered through it like plums in a pudding, or seeds in a watermelon, so the positive and negative charges balance out overall. This model held until 1911, when Rutherford's students Geiger and Marsden fired a beam of tiny, positively charged alpha particles at an extremely thin sheet of gold foil. If Thomson's model were correct, with positive charge spread evenly through the atom, every alpha particle should have passed straight through, or been deflected only slightly. Instead, while most did pass straight through, a small number were deflected sharply, and a few even bounced straight back, something Thomson's spread-out model could not explain at all.

Rutherford concluded that an atom's positive charge, and almost all of its mass, must be concentrated in an extremely small region at the centre, which he called the nucleus, with electrons orbiting it rather like planets around the Sun. Since most alpha particles passed straight through the gold foil undeflected, most of an atom, Rutherford reasoned, must be empty space. The size difference is hard to imagine: an atom is about 10 to the power -10 metres across, while its nucleus is only about 10 to the power -15 metres, a hundred thousand times smaller. If an atom were blown up to the size of a cricket ground, its nucleus would be a single peppercorn sitting at the centre.

Rutherford's model had its own unanswered puzzle: an orbiting electron should, according to the physics of the time, continuously lose energy and spiral into the nucleus within a fraction of a second, meaning atoms should collapse instantly, which plainly does not happen. In 1913, Niels Bohr proposed that electrons do not orbit freely at all, but travel only in certain fixed circular paths, called shells or energy levels, labelled K, L, M, N and so on, moving outward from the nucleus. While an electron stays within one of these allowed shells, Bohr proposed, it does not lose energy at all; it can only jump between shells by absorbing or releasing an exact, fixed amount of energy. This single idea explained why atoms are stable, and why each shell can only hold a fixed maximum number of electrons.

As chemists began identifying more and more elements, Dalton first tried representing each with a small picture, but this quickly became impractical to draw and remember. In 1813, Berzelius suggested using letters from each element's name instead, usually the first letter, or first two letters, capitalised only at the start (hydrogen, H; aluminium, Al, not AL), with some symbols instead coming from Latin, Greek or German names, such as iron's symbol Fe, from the Latin ferrum. Today, an international body called IUPAC approves the official name and symbol of every element, so that scientists anywhere in the world, regardless of language, can read and write the same chemistry.

The number of protons in an atom's nucleus is called its atomic number, given the symbol Z, and it is what identifies an element: every atom with 6 protons is carbon, no matter how many neutrons or electrons it might have, and every atom with 11 protons is sodium. Since a neutral atom has no overall charge, its number of electrons always equals its number of protons. The total count of protons and neutrons together, the particles that actually make up almost all of an atom's mass, is called its mass number, given the symbol A.

Dalton believed all atoms of one element were exactly identical, but scientists later found atoms of the same element that always have the same number of protons, and therefore the same atomic number, yet can have different numbers of neutrons, and so different mass numbers. These variants of the same element are called isotopes. Naturally occurring hydrogen, for example, is almost entirely a form called protium, with no neutrons at all, mixed with a tiny fraction of deuterium, which has one neutron, and a trace of tritium, which has two; all three have exactly one proton and one electron, and so behave identically in chemical reactions, since chemical behaviour depends on electrons, not on neutrons.

Electrons in the outermost shell of an atom, called valence electrons, are the ones involved in bonding. Atoms are most stable when their outermost shell holds 8 electrons (or 2, if that outermost shell is the very first shell), a pattern called the octet rule, achieved by losing, gaining, or sharing electrons with another atom. The number of electrons an atom must lose, gain, or share to complete its octet is called its valency: sodium (electronic configuration 2,8,1) loses its 1 outer electron easily, so its valency is 1; oxygen (2,6) needs 2 more electrons to complete an octet, so its valency is 2; carbon (2,4), unable to easily lose or gain 4 electrons, instead shares 4, giving it a valency of 4.

Knowing what is inside an atom, and how many electrons it needs to feel 'complete', is only useful once you ask what atoms actually do with that knowledge: how they share or transfer those electrons to bond with each other, and how those bonds let us predict and write down the exact formula of a compound before ever making it in a lab. That is exactly where this trail goes next.