Non equilibrium
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Generally the systems treated with the conventional chemical thermodynamics
are either at equilibrium or near equilibeium.
Ilya Prigogine developed the thermodynamic treatment of
open systems that are far from equilibrium. In doing so he has discovered
phenomena and structures of completely new and completely unexpected types. His
generalized, nonlinear and irreversible thermodynamics has found surprising
applications in a wide variety of fields.
The non equilibrium thermodynamics has been applied for explaining how
ordered structures e.g. the biological systems, can develop from disorder. Even
if Onsager's relations are utilized, the classical principles of equilibrium in
thermodynamics still show that linear systems close to equilibrium always
develop into states of disorder which are stable to perturbations and cannot
explain the occurrence of ordered structures.
Prigogine called these systems
dissipative systems, because they are formed and maintained by the
dissipative processes which take place because of the exchange of energy between
the system and its environment and because they disappear if that exchange
ceases. They may be said to live in symbiosis with their environment.
The method which Prigogine used to study the stability of the dissipative
structures to perturbations is of very great general interest. It makes it
possible to study the most varied problems, such as city traffic problems, the
stability of insect communities, the development of ordered biological
structures and the growth of cancer cells to mention but a few examples.
System constraints
In this regard, it is crucial to understand the role of walls and other
constraints, and the distinction between independent processes and
coupling. Contrary to the clear implications of many reference sources, the
previous analysis is not restricted to
homogenous,
isotropic bulk systems which can deliver only PdV work to the
outside world, but applies even to the most structured systems. There are
complex systems with many chemical "reactions" going on at the same time, some
of which are really only parts of the same, overall process. An independent
process is one that could proceed even if all others were unaccountably
stopped in their tracks. Understanding this is perhaps a �thought
experiment� in
chemical kinetics, but actual examples exist.
A gas reaction which results in an increase in the number of molecules will
lead to an increase in volume at constant external pressure. If it occurs inside
a cylinder closed with a piston, the equilibrated reaction can proceed only by
doing work against an external force on the piston. The extent variable for the
reaction can increase only if the piston moves, and conversely, if the piston is
pushed inward, the reaction is driven backwards.
Similarly, a redox
reaction might occur in an
electrochemical cell with the passage of
current in wires
connecting the
electrodes. The half-cell reactions at the
electrodes are constrained if no current is allowed to flow. The current
might be dissipated as
joule
heating, or it might in turn run an electrical device like a
motor doing
mechanical work. An
automobile
lead-acid
battery can be recharged, driving the chemical reaction backwards. In this
case as well, the reaction is not an independent process. Some, perhaps most, of
the Gibbs free energy of reaction may be delivered as external work.
The
hydrolysis of
ATP to
ADP and
phosphate can drive the
force times
distance
work delivered by living
muscles, and
synthesis of ATP is in turn driven by a redox chain in
mitochondria and
chloroplasts, which involves the transport of
ions across the
membranes of these
cellular
organelles. The coupling of processes here, and in the previous examples, is
often not complete. Gas can leak slowly past a piston, just as it can slowly
leak out of a
rubber
balloon. Some reaction may occur in a battery even if no external current is
flowing. There is usually a coupling
coefficient, which may depend on relative rates, which determines what
percentage of the driving free energy is turned into external work, or captured
as "chemical work", a misnomer for the free energy of another chemical process.
Quote
In the preface section to popular book Basic Chemical Thermodynamics
by physical chemist Brian Smith, originally published in 1973, and now in the
5th edition, we find the following overview of the subject as it is perceived in
college:
| The first time I heard
about chemical thermodynamics was when a second-year undergraduate
brought me the news early in my freshman year. He told me a
spine-chilling story of endless lectures with almost three-hundred
numbered equations, all of which, it appeared, had to be committed
to memory and reproduced in exactly the same form in subsequent
examinations. Not only did these equations contain all the normal
algebraic symbols but in addition they were liberally sprinkled with
stars, daggers, and circles so as to stretch even the most powerful
of minds. |
|
