An atom is the smallest unit of stable matter that we speak of in normal conversation. This does not mean that it is the smallest type of matter: atoms themselves are made of smaller components, namely protons, neutrons, and electrons. However, the atom is the smallest unit out of which larger structures of matter are directly built. An atom, in other words, cannot be broken down through chemical reactions.
– Composition –
Atoms actually consist of several subatomic structures of matter: protons, electrons, and neutrons. Within the nucleus, or core, of the atom are the two heaviest, the neutrons and protons. Protons are positively charged subatomic particles; the number of protons contained within an atom determines what element it will be. An atom with just one proton will be hydrogen; an atom with eight protons will be oxygen.
Neutrons have no charge, and therefore do not determine what element an atom will form; however, they do affect the overall weight of the atom, and therefore determine what isotope it will be. Certain isotopes of atoms – that is, certain combinations of protons and neutrons – are less stable than others. Unstable atomic cores tend to fly apart periodically, blasting fragments in all directions. We know this process as radiation. Some of the heaviest metals, like plutonium, are always radioactive. In most cases, however, only some isotopes of a given element will be radioactive. For example, carbon-14 is somewhat radioactive, but carbon-12 (with less neutrons) is entirely stable.
Electrons have virtually no mass, but they do carry a negative charge, which roughly counters the charge of a single proton. Chemical bonding – the reactions in which elements of different atoms combine together to form compounds – occurs when electrons are exchanged between the outer shells of different atoms. Compounds formed through this process and encountered frequently in everyday life include table salt (sodium chloride, or one sodium atom and one chlorine atom) and water (H2O, or two hydrogen atoms and one oxygen atom).
– Origins –
Most of the universe consists of free-floating clouds of hydrogen, which is an unusual element in that it consists only of a single proton, sometimes with one or two neutrons attached, and one orbiting electron. The lightest and smallest atoms, like hydrogen and helium, are believed to have been formed almost immediately after the Big Bang. Since then, subsequent generations of stars have gradually pushed together larger and larger atoms, until we are left with our current periodic table of elements. At this time, naturally occurring elements range from one proton (hydrogen) all the way up to 92 protons (uranium), beyond which a variety of elements can be created artificially in laboratories and reactors, but are too radioactive to survive long in nature.
– The Discovery of Atoms –
Ancient Greek and Indian philosophers were aware that, at least in theory, if you took a sample of physical matter and kept cutting it down into smaller and smaller chunks, eventually you would reach a point at which it was impossible to cut any further. These “atoms” or “uncuttables,” as they were so named by the Greek Leucippus and Democritus, therefore existed in theory long before they existed in any confirmable form in the notebooks or laboratories of early chemists.
During the 1600s and 1700s, however, the first generations of chemists began to piece together a real theory of the atom based on their observations of basic chemical reactions. To Robert Boyle, John Dalton, and Anthony Lavoisier, for example, atoms were a useful counterweight to the pre-existing and now debunked theory of the four elements (earth, air, fire and water). Instead, as Lavoisier pointed out, chemistry could now point to dozens of elements, each of which were later identified as distinct types of atoms. The eventual publication of the first periodic table in the 1860s, by Siberian scientist Dmitri Mendeleev, would not have been possible without at least a basic concept of the atom.
This basic concept, however, said very little about what atoms actually were: it simply pointed out that, whatever they were made up of, they formed discrete and distinguishable elements that behaved very differently under observable conditions, like oxygen, iron, and helium. By the 1890s, new speculation had developed over what atoms actually were. New theories in the early twentieth century, by Ernest Rutherford and Niels Bohr, completed the basic picture of the atom which is still taught to high school science students today: a small nucleus full of neutrons and positively charged protons, surrounded by shells of nearly weightless and negatively charged electrons.
Subsequent developments in quantum physics have led to new understandings of the subatomic world, but these would be too complex to go into without separate articles. Essentially, the work done by Rutherford and Bohr between 1911 and 1915 paved the way for more than a century of popular scientific understanding to come.