How do we measure disorder? The means of doing this came originally
from physics, not chemistry. During the middle of the last century,
physicists were interested in the nature of heat and its manipulation,
an understandable bias in an era of steam power. James Joule, Julius
Mayer, and others concluded after careful experimental measurements
that heat, work, and energy all were merely different aspects of
the same thing.
William Thomson (later Lord Kelvin, of the Kelvin or absolute temperature
scale) and Rudolf Clausius were struck by the fact that the interconversion
of heat and work is a one-way street. It is easy to convert the
energy of work completely into heat, but the reverse transformation
is never complete. Thomson's version of the second law of thermodynamics
states that it is impossible by any cyclic, repeatable process to
take heat and convert it entirely into work without losing some
of this heat to a reservoir at a lower temperature. There can be
no steam engines without condensing cylinders, and part of the available
heat always is lost to the condenser instead of being converted
to useful work. The second law in any of its forms makes heat look
like the lowest or most degraded form of energy: easy to obtain
but hard to reconvert.
We know what Thomson and Clausius a century ago did not. On a molecular
level, kinetic energy is the coordinated motion of all of the molecules
in a solid in the same direction (right). Heat in a solid is the
disunited motion of individual moleculed about their equilibrium
positions. Kinetic energy is organised, coherent motion and heat
is random incoherent motion.
It is easy to change coordinated motion into random motion
but impossible to turn uncoordinated motion completely back
into uniform motion. When we heat a can of soup, all the molecules
begin moving faster but in a random manner. What is the probablility
that, purely by chance, all the molecules in the soup will
begin to move faster in the same direction, taking
the pan and the kitchen wall with them.