THE BIG PICTURE:  AN OVERVIEW

The nature of energy

To discuss energy in a biological context, it seems appropriate to first address the concept of energy in general.  Most of us have a general intuitive idea about what energy is:  it’s what you need to make things happen, whether you’re talking about inside a cell or a more macroscopic context.  Energy is the ability to do work, or in other words to move matter around.  Energy, however, is less a “thing” in and of itself than it is a measurable, quantifiable quality or characteristic of something else.   And, there are certain rules that seem to regulate how things lose and acquire this characteristic.

First of all, although the amount of energy possessed by any given object at any given time can change, both theories and experimental evidence indicate that the total amount of energy in the universe does not ever change (This rule is called the Law of Conservation of Energy).  So, it appears that the quantity of energy in our universe is fixed, and we can’t go around creating new energy from nothing.  Likewise, energy can’t ever just disappear.  If an object gains energy it must have “come from” someplace else, and if it loses energy, that energy must have “gone to” someplace else. 

It is also clear that energy seems to come in different forms.  One form of energy, for example, relates to the relative location of the object that possesses it.  This type of energy is generally called potential energy.  Potential energies derive from forces, which are fundamental interactions between objects.   For example, objects with mass are attracted to other objects with mass, a phenomenon that we describe as the gravitational force.   In a similar vein, objects with a property we call electrical charge (which comes in two forms, positive and negative) are repelled by particles of like charge and attracted to particles of opposite charge. This interaction is termed the electromagnetic force.  This force helps hold atoms together (by providing attraction between positively charged protons and negatively charged electrons) and also is the basis for the electrical and magnetic phenomena that we see around us daily.  Two other fundamental forces also exist.  One, the strong force, acts over only very short distances and describes the interaction between quarks (based on a property of quarks that has been whimsically termed “color”). The other force, the weak force, describes the way particles possessing a particular “spin” interact.  Current theories have developed that portray the electromagnetic and weak forces as aspects of the same underlying “electroweak” force, and work continues to describe all four forces in terms of the same “unified field theory”. 

The strong and weak forces, acting as they do over such miniscule distances, are not directly observable by us.  The gravitational and electromagnetic forces, however, are.  On the cellular level, though, we really don’t have to deal much with gravity.   This is because the strength of the gravitational force is proportional to the mass of the objects involved, and the masses of most cellular constituents are very small indeed.  In fact, you need an object the size of the Earth for the gravitational force to be large enough to be significant.  The force that is most obvious at the cellular level, then, is the electromagnetic force, the effect that dominates on that scale and distance. 

When matter is actually set in motion by one of these attractive or repulsive forces it acquires another form of energy:  kinetic energy, or energy of motion.  Objects that are in motion possess energy in proportion to their mass and the square of the velocity at which they move.  Kinetic energy and potential energy are interchangeable:  for example, if we drop an object that is being held above the surface of the earth, its potential energy will diminish as its kinetic energy increases. 

There are two important rules that regulate the flow of energy through the universe.  The first rule (which we've already mentioned) states that although potential energy can be transformed into kinetic energy and vice versa, and both types can be transferred from one object to another, the total amount of energy in the universe must remain the same:  in other words, energy cannot be created or destroyed.  The second rule involving energy (or second law of thermodynamics) describes what happens when energy is transferred from one form to another or from one object to another.  When this occurs, even though no energy is destroyed, the energy available to “do something” will always diminish.  This is because some of the organized, useful energy (usable to do work) will always be converted into disorganized, unuseful energy (which is not able to be harnessed to do work).  So, the second law says that the total disorder in the universe is always increasing.  It is possible, however, to produce a local decrease in disorder—but the catch is that if disorder decreases in one area, there must be a compensatory increase in disorder someplace else. So, with every transfer of energy, there are some changes in the nature of the energy involved. 

Energy that is in a useful form and that can be used to do work is called free energy.    The amount of disordered energy, on the other hand, is measured by a quantity called entropy.  So, another way to look at all this is to say that the total amount of free energy in the universe is always decreasing and the entropy of the universe is always increasing. 

Where does this disorganized energy go?  In most cases, disorganized energy takes the form of thermal energy, or heat.  Heat is the measurable reflection of random, disordered molecular motion.  It is easy to convert useful energy into thermal energy, but difficult to reverse the process.

 

Energy in chemical systems

So what we’re really interested in, in terms of cellular function, is not really the total energy, but rather the free energy that is available to do the work in the cell.   To understand how free energy is managed by cells, we first have to look at the role of energy in chemical reactions.  Atoms and molecules (reactants) react together, break old bonds and create new bonds to form new molecules (products).  Breaking a molecular bond requires energy to pull against the forces holding the molecule together.  When a molecular bond is formed, on the other hand, energy is released as the potential energy between the reacting atoms decreases with the decreasing distance between them.  The amount of energy required for breaking a particular bond is the same as the amount released on formation of that bond and is called the bond energy.

Molecules will react together spontaneously if the reaction leads to a more stable arrangement for the atoms that make them up.  Systems (such as the chemical system of reactants and products) tend to move from a state of higher towards a state of lower total energy.  This means that the total energy of the resulting molecules is usually less than that of the molecules you began with.

One way to think of this situation is that the energy required to break the old bonds is less than the energy released when the new bonds are formed.  If this is the case, the reaction will give off excess energy (since energy is never destroyed).  On the other hand, if the energy required to break the old bonds is more than the energy released when the new bonds are formed, the reaction would require the input of energy before it can proceed and is much less likely to occur spontaneously. 

            Also, entropy plays a role in determining how easily chemical reactions take place.  In general, reactions in which the local levels of entropy increase are favored over reactions in which the local entropy decreases.  (Remember, though, although levels of entropy in isolated areas may go up or down, the total levels of entropy in the universe are always increasing).

Now here is the somewhat complicated part.  These two conditions (the total energy decreasing and the entropy increasing) do not have to BOTH be favorable in order for a reaction to occur spontaneously.   A large enough decrease in total energy (a favorable condition) may be big enough to outweigh a small decrease in entropy (an unfavorable condition) and a large increase in entropy may be big enough to outweigh a small increase in total energy.

How can we tell when these combinations are right?  By calculating the free energy.   Free energy is found by subtracting the disordered, less useful energy from the total energy of the molecules involved.  The bottom line for a reaction is that the free energy of the products must be less than the free energy of the original reactants.  Assuming that this condition is favorable, then, the reaction will tend to proceed. 

But although the energy criteria that we just talked about tend to push a reaction in one direction or the other, it is important to remember that most reactions can actually go in both directions:  reactants to products, and products back to reactants.  So, as a reaction occurs, what is really happening is that reactants are constantly forming products, and products are constantly forming reactants until a state called equilibrium is reached.  At equilibrium the rate of formation of the products from reactants is exactly equal to the rate of formation of reactants from products.  The energy criteria tell you what you will have at equilibrium:  more product or more reactant.

The relative amounts of reactants and products at equilibrium will be constant for a particular reaction.   For some reactions, the amount of products will be much higher than the amount of reactants at equilibrium; for other reactions the reverse may be true.  Basically, equilibrium is reached at the point where the total free energy for all molecules is at a minimum. In other words, the free energy for that combination of reactants and products is less than for either reactants alone or products alone.  Any further change would lead to an increase in free energy and thus would be unfavorable.  

 

Further reading on this topic

Brown, Guy.  1999.  The Energy of Life.  The Free Press, NY.

 Feynman, Richard Phillips, Leighton, Robert B., and Sands, Matthew.   1964.  The Feynman Lectures on Physics.   Addison-Wesley Pub Co., Reading, Mass.