In 1904 the German chemist Richard Abegg (1869-1910) first suggested that atoms were held together through the transfer of electrons from one atom to another.
To see how this worked, one began by noting that electrons in an atom existed in a series of shells. The innermost shell could hold only 2 electrons, the next 8, the next 18 and so on. It turned out that some electron arrangements were more stable than others. If only the innermost shell contained electrons and it were filled with the 2 electrons that were all it could hold, then that was a stable arrangement. If an atom contained electrons in more than one shell and the outermost shell that held electrons held 8, that was a stable arrangement, too.
Thus, the helium atom has 2 electrons only, filling the innermost shell, and that is so stable an arrangement that helium undergoes no chemical reactions at all. The neon atom has 10 electrons—2 in the innermost shell, and 8 in the next—and it does not react. The argon atom has 18 electrons—2, 8, and 8—and it too is very stable.
But what if an atom did not have its electron shell so neatly filled. The sodium atom has 11 electrons—2, 8, and 1—while the fluorine atom has 9 electrons—2 and 7. If the sodium atom passed one of its electrons to a fluorine atom, both would have the stable configuration of neon—2 and 8. This, therefore, ought to have a great tendency to happen.
If it did happen, though, the sodium atom, minus 1 electron, would have a unit positive charge and would be Na⁺, a positively charged ion. Fluorine with 1 electron in excess would become F⁻, a negatively charged ion. The 2 ions, with opposite charges, would cling together, since opposite charges attract, and thus the molecule of sodium fluoride (NaF) would be formed.
In 1916 the American chemist Gilbert Newton Lewis (1875-1946) carried this notion farther. Atoms could cling together not only as a result of the outright transfer of 1 or more electrons, but through sharing pairs of electrons. This sharing could only take place if the atoms remained close neighbors, and it would take energy to pull them apart and break up the shared pool, just as it would take energy to pull 2 ions apart against the attraction of opposite charges.
In this way the vague notions of atoms clinging together in molecules and being forced apart gave way to a much more precise picture of electrons being transferred or shared. The electron shifts could be dealt with mathematically by a system that came to be called “quantum mechanics” and chemistry was thus made a more exact science than it had ever been before.
The Energy of the Sun
The most serious problem raised by the law of conservation of energy involved the sun. Until 1847, scientists did not question sunlight. The sun radiated vast quantities of energy but that apparently was its nature and was no more to be puzzled over than the fact that the earth rotated on its axis.
Once Helmholtz had stated that energy could neither be created nor destroyed, however, he was bound to ask where the sun’s energy came from. It had, to man’s best knowledge, been radiating heat and light, with no perceptible change, throughout the history of civilization and, from what biologists and geologists could deduce, for countless ages earlier. Where, then, did that energy come from?