K Lerner Lee. Scientific Thought: In Context. Editor: K Lee Lerner & Brenda Wilmoth Lerner. Volume 2. Detroit: Gale, 2009.
One early theory of the structure of the atom was the Bohr model, developed in 1913 by Danish physicist Niels Bohr (1885-1962). By using quantum theory, Bohr’s model improved on the earlier atomic models of British physicists J.J. Thomson (1856-1940) and Ernest Rutherford (1871-1937), which were based on classical (Newtonian) physics.
While working on his doctoral dissertation at Copenhagen University, Bohr studied the theory of radiation being developed by German physicist Max Planck (1858-1947). After graduation, Bohr worked in England with Thomson and subsequently with Rutherford. It was while working in England that Bohr developed his model of atomic structure.
Historical Background and Scientific Foundations
Before Bohr, the foremost classical model of the atom was the Saturnian or planetary model. This theory was suggested by Japanese physicist Nagaoka Hantaro (1865-1950) in 1904 and further developed by Rutherford. It proposed that the atom is structured like the Saturnian ring system or the solar system. In those larger systems, small objects orbit a central, massive object: The planetary model suggested that in the atom, small, electrically negative electrons orbit a relatively massive, positively charged nucleus. However, this view was controversial, and in 1907, J.J. Thomson proposed what was called the “plum pudding” model. In this alternative theory, negative electrons are embedded in a uniform globe of positive charge, much like plums in a pudding.
The plum-pudding model was disproved by 1911, when Rutherford showed that alpha particles fired at atoms sometimes bounce right back the way they came, as if they had struck a massive obstacle in the atom—a nucleus. Some form of planetary model was necessary to explain the behavior of atoms.
The classical planetary model of the atom allowed electrons to orbit at any distance from the nucleus. This implied that when an atom—for example, a hydrogen atom—is heated, its electrons should move away from the nucleus, then move back toward it again, giving up energy in the form of electromagnetic radiation (light). This offered a mechanism for the emission of electromagnetic radiation by atoms. However, the model predicted that an atom should produce a continuous spectrum of colors (light at all different rates of vibration) as it cooled because an electron would gradually give up its energy as it spiraled back closer to the nucleus. Yet spectroscopic observations showed that hydrogen atoms produce only certain colors when they are heated, which causes their electrons to shift repeatedly and rapidly between lower- and higher-energy states. Also, the description of electromagnetic radiation by Scottish physicist James Clark Maxwell (1831-1879) predicted that an electron orbiting a nucleus according to Newton’s laws, like a miniature planet, would continuously lose energy and eventually, unless resupplied with energy from some outside external source, fall into the nucleus. In other words, a classical planetary atom could not be stable, though in the real world, most atoms are perfectly stable.
To account for the observed properties of hydrogen, Bohr proposed a radical new form of the planetary model. In Bohr’s model, electrons can exist only in certain orbits—at certain distances from the nucleus—and that, instead of traveling between orbits like tiny spacecraft traversing the solar system, electrons make instantaneous (quantum) leaps or jumps between allowed orbits.
In the Bohr model, the lowest, most-stable energy level is the innermost orbit (which is called an “orbital” or “shell” to distinguish it from the orbits followed by large objects like planets and satellites). The orbitals in any atom are numbered from the inmost outward. The first orbital of the hydrogen atom is spherical and is assigned a principal quantum number (n) of n = 1. Additional orbital shells are assigned values n = 2, n = 3, n = 4, etc. The orbital shells are not spaced at equal distances from the nucleus; rather, the radius of each shell increases as the square of n. Increasing numbers of electrons can fit into these orbital shells according to the formula 2 n 2. Accordingly, the first shell can hold up to 2 electrons, the second shell (n = 2) up to 8 electrons, the third shell (n = 3) up to 18 electrons. Subshells or suborbitals (designated s, p, d, and f), having differing shapes and orientations, allow each element a unique electron configuration.
As electrons move away from the nucleus, they gain potential energy and become less stable, tending to fall back into lower-energy orbitals closer to the nucleus. An atom with all its electrons in their lowest allowable energy orbits is said to be in a ground state, while an atom with one or more electrons raised to higher-energy orbits is said to be in an excited state. Electrons in atoms may acquire energy from thermal collisions (running into other atoms), collisions with subatomic particles, or absorption of a photon (a quantum packet of light energy). Of all the photons that an atom encounters, only those carrying an amount of energy equal to the energy difference between two allowed electron orbits will be absorbed. Atoms release energy by giving off photons as electrons return to lower-energy orbits.
The leaping of electrons between orbits in the Bohr model accounted for Planck’s observation that atoms emit and absorb electromagnetic radiation only in certain fixed energy units called quanta, rather than in a smooth range of in-between values. Bohr’s model also explained quantum aspects of the photoelectric effect described by Albert Einstein in 1905.
According to the Bohr model, when an electron is excited by energy it jumps from its ground state to an excited state (i.e., a higher energy orbital). The excited atom can then emit energy only in certain fixed quantities as electrons jump back to lower-energy orbits located closer to the nucleus. This energy is emitted as quanta of electromagnetic radiation (photons, particles of light) that have the same energy as the difference in energy between the orbits jumped by the electron. Electron movements between different orbitals explain certain features of electromagnetic spectra. For example, in hydrogen, when an electron returns to the second orbital (n = 2) from a higher orbital, it emits a photon with energy that corresponds to a particular color or spectral line found in the Balmer series of bright lines in the visible portion of the electromagnetic spectrum. The particular color of the photon emitted depends on which higher orbital the electron jumps from. When the electron returns all the way to the innermost orbital (n = 1), the photon emitted has more energy and helps form a line in the Lyman series of bright lines, which is found in the higher-energy, ultraviolet portion of the spectrum.
When the electron returns to the third quantum shell (n = 3), the photon emitted has less energy and helps form a line in the Paschen series, which is found in the lower-energy, infrared portion of the spectrum.
Modern Cultural Connections
Later, more mathematically complex models based on the work of French physicist Louis Victor de Broglie (1892-1987) and Austrian physicist Erwin Schrödinger (1887-1961) that took into account the particle-wave duality of electrons proved more useful to describe atoms with more than one electron. According to these later models, an electron does not move in a circular path around the nucleus like a planet or moon: Rather, it is smeared out in space so that there is some probability of finding it at any points in its orbital. An electron is therefore most accurately imagined as occupying a fuzzy shell centered on the nucleus. Such shells have a variety of shapes, from simple spheres, to two-lobed hourglass shapes, to complex, many-lobed shapes. Today’s standard model of the atom, which incorporates the particles known as quarks (whose existence was first proposed in 1961), further refines the Bohr model so that its properties match all those that have been observed so far in experiments. Despite the need for such improvements, Bohr’s original, simpler model remains fundamental to the study of chemistry. For instance, the valence shell concept used to predict an element’s reactive properties is derived from the Bohr model of the atom.
The Bohr model remains a landmark in scientific thought that poses profound questions for scientists and philosophers. Along with related claims about the physical world made by quantum theory, it has called common-sense concepts of space, time, and language into question. For example, the idea that electrons make quantum leaps from one orbit to another, as opposed to simply moving between orbits, seems counter-intuitive—that is, outside the range of human experience or imagination. How can an object move from A to B without occupying a continuous series of locations between A and B? Yet electrons and other subatomic particles do just this. Bohr once said, “Anyone who is not shocked by quantum theory has not understood it.”
Because quantum physics describes an atomic world that does not obey the rules of everyday physical common sense, the details of atomic structure cannot be portrayed in drawings without leaving out some aspect of their nature. This is why almost all pictures or drawings showing the structure of the atom show a simplified version of the Bohr model, with electrons orbiting the nucleus like tiny, round planets. These pictures can be an aid to elementary understanding of the atom, but do not show what atoms look like. In fact, in quantum physics, atoms do not look like anything. They do not have any appearance at all because they cannot interact with light as do larger objects. This fact arises from their very nature, not from a mere lack of appropriate lighting, so atoms cannot be thought of as invisible or hard to see: they simply lack any visual appearance whatever, seen or unseen.
The phrase “quantum leap” (or “quantum jump”) has entered everyday English as referring to a very large, sudden change. In the physics of the Bohr model of the atom, however, this phrase refers to the shifting of an electron from one orbital to another without traveling through any intermediate point. Such instantaneous shifts can be made only by very small particles moving over extremely small distances—almost the opposite of the popular phrase’s meaning. This is an example of how ideas from physics often enter popular culture in distorted forms. Another example is the saying, “Einstein showed that everything is relative.” He did not: He showed that some things are relative and others are not relative.