Entropy is defined as a “thermodynamic property that
is the measure of a system’s thermal energy per unit temperature that is
unavailable for doing useful work”.
A common example of increasing entropy is ice melting in
a warm room, described in 1862 by Rudolf Clausius as an increase in the
“disgregation” (the magnitude of the degree in which the molecules of a body
are separated from each other) of the water molecules in ice: order leading to
disorder (more random).
In a system isolated from the environment, the entropy of
that system will tend to not decrease. In addition, it is impossible for any
device operating on a cycle to produce net work from a single temperature
reservoir. The production of net work requires flow of heat from a hotter to
colder reservoir. As a result, there is no perpetual motion system. A reduction
in the increase of entropy in a specific process, such as a chemical reaction,
means that it is energetically more efficient.
The entropy of a system that is not isolated may
decrease. For example, an air conditioner may cool the air in a room, reducing
the entropy of the air in the room. The heat expelled from the room, which the
air conditioner transports and discharges to the outside air, will always make
a bigger contribution to the entropy of the environment than will decrease the
entropy of the air in the room. Therefore, the total entropy of the room plus
the entropy of the environment increases. This example proves the second law of
thermodynamics.
References
Second law of thermodynamics. (n.d.). Retrieved
from http://hyperphysics.phy-astr.gsu.edu/hbase/thermo/seclaw.html
Benson, T. (2008, July 11). Second law of
thermodynamics. Retrieved from http://www.grc.nasa.gov/WWW/k-12/airplane/thermo2.html
Laws of thermodynamics. (2010, May 18). Retrieved
from http://www.emc.maricopa.edu/faculty/farabee/biobk/biobookener1.html
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