The discovery of the structure of the atom sharpened the understanding
of chemical energy.
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.