Energy and the
The concept of energy is fundamental
to all life processes. We usually express
energy as the capacity to do work, that
is, to bring about change. Yet energy is
a somewhat abstract quantity that is
difficult to define and elusive to measure.
Energy cannot be seen; it can
only be defined and described by how
it affects matter.
|Figure 4-1 Solar energy sustains
virtually all life on earth.
each energy transfer, however,
of the energy is lost
Energy can exist in either of two
states: kinetic or potential. Kinetic
is the energy of motion. Potential energy
is stored energy,
energy that is not doing work but has
the capacity to do so. Energy can be
transformed from one state to another.
Especially important for living organisms
is chemical energy, a form of
potential energy that is stored in chemical
bonds of molecules. Chemical
energy can be tapped when bonds are
rearranged to release kinetic energy.
Much of the work done by living
organisms involves the conversion of
potential energy to kinetic energy.
The conversion of one form of
energy to another is governed by the
two laws of thermodynamics. The first
law of thermodynamics
energy cannot be created or destroyed.
It can change from one form to
another, but the total amount of
energy in a system remains the same.
In short, energy is conserved. If we
burn gasoline in an engine, we do not
create new energy but merely convert
the chemical energy in gasoline to
another form, in this example,
mechanical energy and heat. The second
law of thermodynamics
in the prologue to this section,
concerns the transformation of energy.
This fundamental law states that a
closed system moves toward increasing
disorder, or entropy, as energy is
dissipated from the system (Figure 4-
2). Living systems, however, are open
systems that not only maintain their
organization but also increase it, as
during the development of an animal
from egg to adult.
To describe the energy changes that
take place in chemical reactions, biochemists
use the concept of free energy.
Free energy is simply the energy
in a system available for doing work. In
a molecule, free energy equals the
energy present in chemical bonds
minus the energy that cannot be used.
The majority of reactions in cells release
free energy and are said to be exergonic
, out, + ergon
Such reactions are spontaneous and
always proceed “downhill” since free
energy is lost from the system. Thus:
However, many important reactions
in cells require the addition of free
energy and are said to be endergonic
, within, + ergon
Such reactions have to be “pushed
uphill” because they end up with more
energy than they started with:
As we will see in a later section, ATP is
the ubiquitous, energy-rich intermediate
used by organisms to power important
uphill reactions such as those required
for active transport of molecules across
membranes and cellular synthesis.
|Figure 4-2 Diffusion of a solute through a solution, an example of entropy.
When the solute (sugar molecules) is first introduced into a
solution, the system is ordered and unstable (B). Without
energy to maintain this order, the solute particles become
distributed into solution, reaching a state of disorder
(equilibrium) (D). Entropy has increased from left diagram to