Niels Bohr (1885-1962) was born in Copenhagen, Denmark, to Christian Bohr, a devout Lutheran and professor of physiology at the University of Copenhagen, and Ellen Adler Bohr, from a wealthy Jewish family engaged in Danish banking and politics. Young Niels and his brother, Harald, played soccer with a passion, but he had an even greater interest in science and studied under J. J. Thomson at Trinity College, Cambridge. He received his doctorate from the University of Copenhagen in 1911 with the thesis, “Studies in the Electron Theory of Metals.” Bohr then moved to Manchester, England, for postdoctoral studies under Ernest Rutherford, applying himself to some problems with the new model of the atom.

Rutherford’s proposed idea for the atomic structure had electrons liter­ally orbiting a heavy nucleus, like planets orbiting a sun. It was an appealing model in that it matched well-known configurations, but the astronomical analogy was deeply flawed on the atomic level. Inter-electron interferences, in which like-charged particles get in each other’s way in orbit, could be ignored by considering the simplest case: the hydrogen atom, with only one electron orbiting a simple nucleus consisting of a single positively charged proton. Even this case would not work. Technically, the electron in orbit was always accelerating, as its traveling direction had to constantly change as it maintained a circular path. According to Maxwell’s equations, an acceler­ating electron would emit light. Emitting light would drain the electron of energy, and it would spiral down out of orbit and eventually crash into the nucleus, as well as making hydrogen glow continuously or until its atoms self-destructed. Hydrogen atoms had no tendency to self-destruct. They seemed to last forever, and there was no continuous glow from emitted light.

However, hydrogen could be made to glow under duress. In the great rush to explore vacuum tubes and high-voltage effects in the late 19th century, physicists had filled evacuated tubes with individual gases, such as argon, neon, xenon, and even nitrogen and hydrogen. They found that under the excitation of thousands of volts, gases would glow, with char­acteristic colors. An organized way to classify gases by excitation color was to direct the light through a slit, making a narrow beam, and then through a glass prism, separating it into a spectrum, using a device that had been around since 1859. The results were exciting and useful. A cer­tain gas would throw narrow lines of consistently distinct colors on a scale representing the color spectrum of light.

There was no theory as to why gases such as hydrogen emitted light only in certain colors, but a Swiss mathematician, Johann Balmer (1825­98) came up with a formula that would predict, with amazing accuracy,

the positions of the spectral lines from electrically excited hydrogen on a scale of light wavelength. The lines of hydrogen light were named the Balmer Series in his honor, and in 1885 there was much excitement over this finding but still no ideas as to why this formula worked.

In 1913, Bohr, working in Manchester, stared at Balmer s formula and realized something: To derive the correct light-wavelengths for hydrogen, numbers are plugged into the equation. Not just any numbers are inserted, but integers, such as 3,4,5, and so on. Immediately, the orbital structure of the hydrogen atom became clear to Bohr. The electron bound to the posi­tively charged nucleus was normally in a basic orbit called a ground state. The ground state was the minimum energy a hydrogen electron could have and still be orbiting. Add energy to the atom, by establishing a high voltage across it in a tube, and the electron responds by jumping to a higher orbit. Add more energy, and the electron jumps to an even higher orbit. Remove the excitation, and the electron jumps back down to ground state. The term orbit had been misused. It was not literally a satellite-style orbit, but an energy state. It was when an elec­tron jumps from one energy state to another that it actually experiences an acceleration and therefore radi­ates a Maxwellian particle of light into space. An electron can jump from its highest orbit to ground, from its highest orbit to a semi-high — est orbit, or from a semi-highest orbit to ground, and each abrupt transi­tion results in a different wavelength, or color, of light, as predicted by inte­gers inserted into Balmer s formula.

Each energy transition represents a quantum leap, instead of a continu­ous decay of an orbit, and this new model neatly explains Planck’s find­ing that there is an indivisible quan­tum of light. One electron in one hydrogen atom, making one orbital transition produces one photon of a predictable color, and there is no way to divide that process down to make a smaller quantity of light.

Not only had Bohr explained why Balmer s equation seemed to work and affirmed Plancks findings, he had validated Einsteins energy-packet theory of light, and further experimental confirmation impressed physi­cists worldwide. He had invented quantum mechanics.