BIOENERGETICS AND ENZYMES
Material is presented on energy flow and the laws of thermodynamics. The production of
biologically useful energy is presented from an electrochemical perspective. Examples provided
in lecture are not in the textbook, although additional information about calculating free energy
can be found in the appendices of the textbook. Methane is used as an example since it is the
simplest organic compound.
After attending lecture and reading this chapter you should be able to:
1) Know the meaning of the first and second laws of thermodynamics as they relate to
2) Understand the concept of entropy and free energy and how free energy is calculated.
3) Understand the concept of oxidation/reduction coupled reactions.
4) Discuss reduction potential and its relationship to exergonic and endergonic processes.
5) Describe how enzymes function, their components and how they work.
I. The Laws of Thermodynamics
First law: The total energy in a system remains constant. Energy can be neither created nor
destroyed, but it may be transferred or redistributed from one state to another: e.g. a reaction.
Energy that is consumed or released from the conversion of substrate A to product B can be
kCal: 1 kilocalorie is the amount of heat energy needed to raise the temperature 1 C in 1
kilogram of water. This unit is still used, particularly in nutrition. It relates energy of
compounds to the heat liberated from combustion.
kJ: 1kilojoule is a unit that is normally used in physics. One kJ equals a columb-volt, which is a
very handy unit in converting energy to electrical units that are relevant to biology. This is now
the accepted unit in biology for expressing the "cash" equivalents of energy.
4.18 kJ = 1 kCal
Second law: Physical and chemical processes proceed in such a way that the disorder, i.e.
entropy, of the universe increases to the maximum possible.
Thus, to create order in an otherwise disorderly universe, endergonic (energy requiring)
reactions are coupled to exergonic (energy releasing) reactions.
1 Gibbs Free Energy: Methane heated in the presence of2O to completely combust it t2 CO will
release a finite and constant amount of free energy available for doing work:
CH 4 + 2O 2 -> CO 2 + 2H 2
880 kJ 0 kJ 0 kJ 0 kJ
We subtract the left hand part of the equation from the right hand part, which gives us a negative
value of - 880 kJ. This means that the reaction is exergonic and will proceed from left to right if
a catalyst is present.
Free energy change is referred to as ΔG . If it is negative it is exergonic. In the example above
ΔG = -880 kJ.
In biology, methane is oxidized in stepwise fashion to capture energy as an electrochemical
proton gradient (discussed below). Each of the steps below represents reactions brought about
by methanotrophic bacteria (those that consume methane).
CH 4 + ½ O 2 CH 3H
880 KJ 727 KJ -153
CH 3H + ½ O 2 CH 2 + H2O
727 KJ 568 KJ -159
CH 2 + ½ O 2 HCOOH
568 KJ 255 KJ -313
HCOOH+ ½ O 2 CO 2 + H2O
255 KJ -255
Note that all the reactions above are exergonic and will provide energy to the organism that
carries them out.
II. Oxidation-Reduction Reactions
Oxidation-reduction (redox) reactions involve the transfer of electrons from a donor (reducing
agent or reductant) to an acceptor (oxidizing agent or oxidant). The reactions show the net result
of the two half-reactions. Let us write the last reaction from the previous page in a redox form.
Formic acid, for example, is the substrate that is oxidized to carbon dioxide. It looses electrons:
HCOOH CO 2 + 2H + 2e -
Molecular oxygen is the substrate that is reduced. It gains electrons:
½ O + 2H + 2e - H O
2 When we add the two half reactions, the equation must be balance by: 1) ensuring that all atoms
are represented on both sides, 2) keeping the ionic (e.g. charge) balance equivalent on each side,
and 3) ensuring that all of the electrons and protons are accounted for in the redox reaction. The
final equation is:
HCOOH + ½ O 2 CO 2 + H O2
Why bother dividing the reaction into two half reactions?
The electrons and protons that are released from enzymatic oxidation of formic acid and
virtually all organic compounds do not go directly to molecular oxygen in biological
systems. The heat released from a single step would literally fry the cell. All biological
reactions involving protons and electrons pass through a series of electron carriers, each of which
has a progressively more positive oxidation-reduction (redox) potential.
The standard redox potential (o ) is a measure of the tendency to gain or lose electrons relative
to the hydrogen redox potential (H ). Redox potentials represent equal concentrations of the
oxidized and reduced forms of the molecule: i.e. when they are at equilibrium.
In chemistry, the hydrogen potential is assigned a value of 0 at pH 0. Because most biological
systems operate near neutrality, the E at pH 7.0 is -0.414 volts (or 414 mV).
2H + 2e