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— CH. 1 · INTRODUCTION —

Atom

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  • An atom is the basic particle of a chemical element, and the swarm of electrons inside it is too small to measure with any available technique. A single human hair is about a million carbon atoms wide. A drop of water holds roughly two sextillion oxygen atoms, and twice that many hydrogen. Nothing in everyday experience prepares you for objects this small. They sit beneath the shortest wavelength of visible light, so no conventional microscope will ever show one. They are so small that classical physics simply cannot predict how they behave. The word itself came from an ancient Greek term, atomos, meaning uncuttable. That name turned out to be wrong. The thing we still call uncuttable is built from protons, neutrons, and electrons, and even those split further into smaller pieces. So how did anyone work out the contents of something they could not see? And why does the count of one particle, the proton, decide whether you are holding sodium or copper? The answers run from a Manchester foil experiment to the cores of stars.

  • Eleven protons make sodium. Twenty-nine make copper. The number of protons inside an atom is what distinguishes one chemical element from another, and nothing else does the job. Atoms with the same proton count but a different number of neutrons are isotopes of the same element. Hydrogen shows this cleanly. Every hydrogen atom carries exactly one proton, yet some have no neutrons, called protium, some have one, called deuterium, and some have two, called tritium. The known elements run from single-proton hydrogen up to oganesson, with 118 protons. Once an atom climbs past 82 protons, every isotope it has is radioactive, though element 83, bismuth, decays so slightly that the effect is practically negligible. Henry Moseley pinned this idea down in 1913. He found that the frequencies of X-ray emissions from an excited atom were a mathematical function of its atomic number. That number, the proton count, matched each element's ordinal place on the periodic table. Ernest Rutherford had already named that count the atomic number, and it gave chemists a clean way to tell the elements apart.

  • More than 99.94% of an atom's mass sits in its nucleus, packed into a tiny volume at the center. Protons carry positive charge and neutrons carry none, so the nucleus is positively charged overall. Electrons carry negative charge, and that opposing pull is exactly what binds them to the nucleus. When the number of protons equals the number of electrons, the atom is electrically neutral. Break that balance and you get an ion. An atom with more electrons than protons carries an overall negative charge and is called an anion. An atom with more protons than electrons carries a positive charge and is called a cation. The electromagnetic force does the binding, pulling electrons toward the protons. Inside the nucleus a different force takes over, the nuclear force, which draws protons and neutrons together. Usually that force overpowers the electromagnetic repulsion that pushes the positive protons apart. Under certain conditions the repulsion wins instead. Then the nucleus splits and leaves behind different elements, a form of nuclear decay. This tug between attraction and repulsion runs straight into the question of which nuclei survive at all.

  • In 1897, J. J. Thomson watched cathode rays bend in electric and magnetic fields. That bending meant the rays were not light but electrically charged particles, and the direction of the deflection showed the charge was negative. He measured these particles at 1,700 times lighter than hydrogen, the lightest atom. He called them corpuscles, and they were later renamed electrons. The discovery meant the atom was not indivisible after all. Thomson pictured the electrons embedded in a sphere of positive charge spread evenly throughout, a layout later nicknamed the plum pudding model, though Thomson and his colleagues never used the phrase. Between 1908 and 1913, Ernest Rutherford, working with Hans Geiger and Ernest Marsden, fired a beam of alpha particles at thin metal foils to study how they scattered. A small number bounced back by angles greater than 90 degrees. Thomson's diffuse charge could never produce a field strong enough for that. Rutherford concluded the positive charge had to be crammed into a tiny central nucleus carrying almost all the atom's mass, with electrons spread around it in a cloud. Years later, a scanning tunneling microscope would finally let researchers visualize atoms at the surfaces of solids, using quantum tunneling rather than light.

  • Back in 1815, William Prout noticed that the atomic weights of many elements were multiples of hydrogen's weight, which holds for all of them once isotopes are counted. In 1919, Rutherford bombarded nitrogen gas with alpha particles and detected hydrogen ions flying out. He decided the alpha particles were splitting the nitrogen nuclei. By 1920 he had accepted the hydrogen nucleus as a single positive particle and named it the proton. One problem nagged at him. The atomic weight of each element ran higher than its proton count, so he proposed an unknown neutral particle with a mass like the proton's to carry the surplus. Walter Bothe found a clue in 1928, when beryllium hit by alpha particles gave off a penetrating, electrically neutral radiation. That radiation could knock hydrogen atoms out of paraffin wax. Many thought it was gamma radiation, but James Chadwick judged the ionization too strong for that if energy and momentum were to be conserved. In 1932, Chadwick exposed elements like hydrogen and nitrogen to the beryllium radiation and measured the recoiling particles. He deduced the radiation was made of neutral particles with a mass close to the proton's. These were Rutherford's neutrons, found at last.

  • Classical mechanics carried a fatal flaw for the atom. An accelerating charged particle radiates electromagnetic energy and loses speed, and circular motion counts as acceleration. So an orbiting electron should spiral straight into the nucleus. In 1913, Niels Bohr proposed that electrons could occupy only a finite set of orbits and jump between them only in discrete energy steps, each step matching the absorption or release of a photon. This explained why orbits stayed stable and why elements emit and absorb radiation in discrete spectra. Bohr's model worked only for hydrogen and failed for atoms with more than one electron. The fix arrived through waves. In 1924, Louis de Broglie suggested all particles behave like waves to some degree. In 1925, Werner Heisenberg published the first consistent mathematical formulation of quantum mechanics, matrix mechanics. In 1926, Erwin Schrodinger used de Broglie's idea to build his equation, describing electrons as three-dimensional waveforms rather than points. A direct consequence followed in 1927, Heisenberg's uncertainty principle, which makes it impossible to know both an electron's position and its momentum precisely at once. The planetary picture gave way to atomic orbital zones, regions where an electron is most likely to be found. Each orbital corresponds to a particular energy level, and the gaps between those levels produce the spectral lines that fingerprint every element.

  • About 339 nuclides occur naturally on Earth, and 251 of them, roughly 74%, have never been observed to decay. Only 90 are theoretically stable, while another 161 have simply never been seen to break down, even though decay is energetically possible for them. The numbers grow stranger with the long-lived. An additional 35 radioactive nuclides have half-lives longer than 100 million years, old enough to have survived since the Solar System's birth. Together these 286 form the primordial nuclides. A further 53 short-lived nuclides appear as daughters of primordial decay, such as radium from uranium, or from natural energy processes like cosmic ray bombardment that makes carbon-14. Stability hinges on the ratio of protons to neutrons and on certain magic numbers that fill quantum shells. Tin shows this vividly. Its filled shell of 50 protons gives it the largest number of stable isotopes of any element, ten. Of the 251 stable nuclides, only four have both an odd proton count and an odd neutron count, hydrogen-2, lithium-6, boron-10, and nitrogen-14. The heaviest stable atom is lead-208, with a mass of 207.9766521. Beyond a certain point no nucleus holds together. Theorists imagine an island of stability for superheavy elements with atomic numbers 110 to 114, with predicted half-lives ranging from a few minutes to millions of years.

  • Electrons are thought to have existed since the early stages of the Big Bang, but neutral atoms came much later. They began to dominate over charged particles 380,000 years after the Big Bang, an epoch called recombination, when the cooling universe finally let electrons attach to nuclei. In its first three minutes, Big Bang nucleosynthesis produced most of the helium, lithium, and deuterium in the universe, and perhaps some beryllium and boron. It produced no carbon and no heavier elements at all. Those came from stars. At the Sun's core, protons need energies of 3 to 10 keV to push past their mutual repulsion, the coulomb barrier, and fuse. Fusion that builds nuclei lighter than iron and nickel, up to a total of about 60 nucleons, releases more energy than it takes, which is what keeps fusion in stars self-sustaining. Past that point the process turns endothermic, so the heaviest nuclei cannot fuel a star. Elements heavier than iron form elsewhere, in supernovae and colliding neutron stars through the r-process, and in AGB stars through the s-process, both involving neutron capture. The atoms of Earth and its inhabitants were mostly present in the nebula that collapsed to form the Solar System. One leftover detail closes the circle. Helium was discovered in the spectrum of the Sun 23 years before anyone found it on Earth.

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Common questions

What is an atom made of?

An atom consists of a nucleus of protons and generally neutrons, surrounded by an electromagnetically bound swarm of electrons. More than 99.94% of an atom's mass sits in the nucleus. In the Standard Model, protons and neutrons are themselves composed of quarks, while electrons are elementary particles with no internal structure.

How does the number of protons define a chemical element?

The number of protons in an atom, called its atomic number, distinguishes each chemical element from every other. Any atom with 11 protons is sodium, and any atom with 29 protons is copper. Atoms with the same proton count but different neutron counts are isotopes of the same element.

Who discovered the electron and the atomic nucleus?

J. J. Thomson discovered the electron in 1897 when he found that cathode rays could be deflected by electric and magnetic fields. Between 1908 and 1913, Ernest Rutherford, with Hans Geiger and Ernest Marsden, discovered the nucleus by firing alpha particles at thin metal foils and observing some deflect by angles greater than 90 degrees.

When were the proton and neutron discovered?

Ernest Rutherford named the proton in 1920 after concluding in 1919 that alpha particles splitting nitrogen nuclei released hydrogen nuclei. James Chadwick discovered the neutron in 1932 by exposing elements to beryllium radiation and measuring the recoiling particles to find they were neutral with a mass close to the proton's.

How small is an atom?

Atoms are typically around 100 picometers across, and a human hair is about a million carbon atoms wide. They are smaller than the shortest wavelength of visible light, so they cannot be seen with conventional optical microscopes, though individual atoms can be observed with a scanning tunneling microscope.

Where do atoms come from?

Big Bang nucleosynthesis produced most of the helium, lithium, and deuterium in about three minutes but no carbon or heavier elements. Carbon up to iron formed in stars through nuclear fusion, while elements heavier than iron formed in supernovae and colliding neutron stars through the r-process and in AGB stars through the s-process.

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