Attempts at understanding matter date back to ancient Greek philosophers, who viewed matter as being composed of elements, or simpler substances. Two models were developed that considered matter to be (1) continuous, or infinitely divisible, or (2) discontinuous, made up of particles called atoms.

In the early 1800s, Dalton published an atomic theory, reasoning that matter was composed of hard, indivisible atoms that were joined together or dissociated during chemical change.

When a good air pump to provide a vacuum was invented in 1885, cathode rays were observed to move from the negative terminal in an evacuated glass tube. The nature of cathode rays was a mystery. The mystery was solved in 1887 when Thomson discovered they were negatively charged particles now known as electrons. Thomson had discovered the first elementary particle of which atoms are made and measured their charge-to-mass ratio.

Rutherford developed a solar system model based on experiments with alpha particles scattered from a thin sheet of metal. This model had a small, massive, and positively charged nucleus surrounded by moving electrons. These electrons were calculated to be at a distance from the nucleus of 100,000 times the radius of the nucleus, so the volume of an atom is mostly empty space. Later, Rutherford proposed that the nucleus contained two elementary particles: protons with a positive charge and neutrons with no charge. The atomic number is the number of protons in an atom. Atoms of elements with different numbers of neutrons are called isotopes. The mass of each isotope is compared to the mass of carbon-12, which is assigned a mass of exactly 12.00 atomicmass units. The mass contribution of the isotopes of an element according to their abundance is called the atomic weight of an element. Isotopes are identified by their mass number, which is the sum of the number of protons and neutrons in the nucleus. Isotopes are identified by their chemical symbol, with the atomic number as a subscript and the mass number as a superscript.

Bohr developed a model of the hydrogen atom to explain the characteristic line spectra emitted by hydrogen. His model specified that (1) electrons can move only in allowed orbits, (2) electrons do not emit radiant energy when they remain in an orbit, and (3) electrons move from one allowed orbit to another when they gain or lose energy. When an electron jumps from a higher orbit to a lower one, it gives up energy in the form of a single photon. The energy of the photon corresponds to the difference in energy between the two levels. The Bohr model worked well for hydrogen but not for other atoms.

Schrödinger and others used the wave nature of the electron to develop a new model of the atom called wave mechanics, or quantum mechanics. This model was found to confirm exactly all the experimental data as well as predict new data. The quantum mechanical model describes the energy state of the electron in terms of quantum numbers based on the wave nature of the electron. The quantum numbers defined the probability of the location of an electron in terms of fuzzy regions of space called orbitals.

The periodic table has horizontal rows of elements called periods and vertical columns of elements called families. Families have the same outer orbital electron configurations, and it is the electron configuration that is mostly responsible for the chemical properties of an element.

Summary of Equations