22 Chapter 1. What the ancients knew[[Student version, December 8, 2002]]
Figure 1.5:(Molecular structure sketches.) (a)The original molecule shown is chiral; it cannot be rotated into
any orientation that makes it identical to its mirror image. The original and its mirror (or “enantiomeric”) form are
chemically different, even though they have the same atoms, the same bonds, and the same bond angles. (b)This
molecule, in contrast, is nonchiral: The original molecule can be rotated until it coincides with its mirror image.
[Copyrighted figure; permission pending.]
1.5.3 Molecules have definite internal energies
Section 1.1.2 briefly alluded to the chemical energy stored in a match. Indeed the atoms making
up a molecule carry a definite amount of stored energy, which is said to reside in “chemical bonds”
between the atoms. The chemical bond energy drives toward lower values just as any other form of
stored energy (for example the potential energy of the weight in Figure 1.3). Indeed the chemical
bond energy is just another contribution to the quantityEappearing in Equation 1.4 on page 7.
Molecules generally prefer to indulge in heat-liberating (exothermic)reactions over heat-accepting
(endothermic)ones, but we can nevertheless get them to adopt higher-energy states by adding
energy from outside. For example, we can split (orhydrolyze)water by passing electrical current
through it. More precisely, Chapter 8 will show thatchemical reactions proceed in the direction
that tends to lower the free energy,just as in the osmotic machine.
Even an unstable molecule may not spontaneously split up until a large “activation energy” is
supplied; this is how explosives store their energy until they are detonated. The activation energy
can be delivered to a molecule mechanically, by collision with a neighbor. But this is not the only
possibility. In one of his five historic papers written in 1905, Albert Einstein showed thatlight, too,
comes in packets of definite energy,calledphotons.Amolecule can absorb such a packet and then
hop over its activation energy barrier, perhaps even ending in a higher-energy state than initially.
The explanations for all of the familiar facts in this subsection and the previous one come from
abranch of physics called “quantum mechanics.” Quantum mechanics also explains the numerical
values of the typical atomic sizes and bond energies in terms of a fundamental physical constant,
the Planck constant.Inthis book we will take all these values as just experimentally determined
facts, sidestepping their quantum origins altogether.
How can there be a “typical” bond energy? Don’t some reactions (say, in a stick of dynamite)