Travel Through Time

EARLY HISTORY
In the 18th and 19th centuries, Newtonian, or classical, mechanics appeared to provide a wholly accurate description of the motions of bodies—for example, planetary motion. In the late 19th and early 20th centuries, however, experimental findings raised doubts about the completeness of Newtonian theory. Among the newer observations were the lines that appear in the spectra of light emitted by heated gases, or gases in which electric discharges take place. From the model of the atom developed in the early 20th century by the English physicist Ernest Rutherford, in which negatively charged electrons circle a positive nucleus in orbits prescribed by Newton’s laws of motion, scientists had also expected that the electrons would emit light over a broad frequency range, rather than in the narrow frequency ranges that form the lines in a spectrum.
Another puzzle for physicists was the coexistence of two theories of light: the corpuscular theory, which explains light as a stream of particles, and the wave theory, which views light as electromagnetic waves. A third problem was the absence of a molecular basis for thermodynamics. In his book Elementary Principles in Statistical Mechanics (1902), the American mathematical physicist J. Willard Gibbs conceded the impossibility of framing a theory of molecular action that reconciled thermodynamics, radiation, and electrical phenomena as they were then understood.

EINSTEIN’S CONTRIBUTION
The next important developments in quantum mechanics were the work of German-born American physicist and Nobel laureate Albert Einstein. He used Planck’s concept of the quantum to explain certain properties of the photoelectric effect—an experimentally observed phenomenon in which electrons are emitted from metal surfaces when radiation falls on these surfaces.

According to classical theory, the energy, as measured by the voltage of the emitted electrons, should be proportional to the intensity of the radiation. The energy of the electrons, however, was found to be independent of the intensity of radiation—which determined only the number of electrons emitted—and to depend solely on the frequency of the radiation. The higher the frequency of the incident radiation, the greater is the electron energy; below a certain critical frequency no electrons are emitted. These facts were explained by Einstein by assuming that a single quantum of radiant energy ejects a single electron from the metal. The energy of the quantum is proportional to the frequency, and so the energy of the electron depends on the frequency.