Atomic structure and spectra

Atomic structure and spectra The idea that matter is subdivided into discrete and further indivisible building blocks called atoms dates back to the Greek philosopher Democritus, whose teachings of the 5th century B.C. are commonly ac­cepted as the earliest authenticated ones concerning what has come to be called atom­ism by students of Greek philosophy. The weaving of the philosophical thread of atom­ism into the analytical fabric of physics began in the late 18th and the 19th centuries. Robert Boyle is generally credited with introducing the concept of chemical elements, the irreducible units of which are now recognized as individual atoms of a given el­ement. In the early 19th century John Dalton developed his atomic theory, which postulated that matter consists of indivisible atoms as the irreducible units of Boyle’s elements, that each atom of a given element has identical attributes, that differences among elements are due to fundamental differences among their constituent atoms, that chemical reactions proceed by simple rearrangement of indestructible atoms, and that chemical compounds consist of molecules which are reasonably stable aggregates of such indestructible atoms. See CHEMISTRY.

The work of J. J. Thomson in 1897 clearly demonstrated that atoms are electromag­netically constituted and that from them can be extracted fundamental material units bearing electric charge that are now called electrons. The electrons of an atom account for a negligible fraction of its mass. By virtue of overall electrical neutrality of every atom, the mass must therefore reside in a compensating, positively charged atomic component of equal charge magnitude but vastly greater mass. See ELECTRON.

Thomson’s work was followed by the demonstration by Ernest Rutherford in 1911 that nearly all the mass and all of the positive electric charge of an atom are con­centrated in a small nuclear core approximately 10,000 times smaller in extent than an atomic diameter. Niels Bohr in 1913 and others carried out some remarkably suc­cessful attempts to build solar system models of atoms containing planetary pointlike electrons orbiting around a positive core through mutual electrical attraction (though only certain “quantized” orbits were “permitted”). These models were ultimately super­seded by nonparticulate-matter wave quantum theories of both electrons and atomic nuclei.

The modern picture of condensed matter (such as solid crystals) consists of an aggre­gate of atoms or molecules which respond to each other’s proximity through attractive electrical interactions at separation distances of the order of 1 atomic diameter (ap­proximately 10-¹⁰ m) and repulsive electrical interactions at much smaller distances. These interactions are mediated by the electrons, which are in some sense shared and exchanged by all atoms of a particular sample, and serve as a kind of interatomic glue which binds the mutually repulsive, heavy, positively charged atomic cores together.

The hydrogen atom is the simplest atom, and its spectrum ( or pattern of light frequen­cies emitted) is also the simplest. The regularity of its spectrum had defied explanation until Bohr solved it with three postulates, these representing a model which is useful, but quite insufficient, for understanding the atom.

Postulate 1: The force that holds the electron to the nucleus is the Coulomb force between electrically charged bodies.

Postulate 2: Only certain stable, nonradiating orbits for the electron’s motion are possible, those for which the angular momentum is an integral multiple of h/2rr: (Bohr’s quantum condition on the orbital angular momentum). Each stable orbit represents a discrete energy state.

Postulate 3: Emission or absorption of light occurs when the electron makes a tran­sition from one stable orbit to another, and the frequency v of the light is such that the difference in the orbital energies equals hv (A. Einstein’s frequency condition for the photon, the quantum of light).