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Biology 110 - Basic Concepts and Biodiversity


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Energy I - Thermodynamics


You should have a working knowledge of the following terms:

  • adenosine triphosphate (ATP)
  • chemical work
  • endergonic reaction
  • endothermic reaction
  • energy
  • energy coupling
  • enthalpy (H)
  • entropy (S)
  • exergonic reaction
  • exothermic reaction
  • free energy (G)
  • kinetic energy
  • mechanical work
  • metabolism
  • non-spontaneous reaction
  • potential energy
  • spontaneous reaction
  • thermodynamics
  • transport work

Introduction and Goals

One of the major characteristics of a living organism is that it can obtain and process energy. As you will learn, the basic principles of energy transduction are widely applicable and apply not just to life, but also to other transformations (e.g., those performed by these windmills). The basic molecules that constitute living organisms were discussed in Tutorial 16 (Carbon and Life). This tutorial will concentrate on how some biomolecules are used for energetic purposes, and will provide an overview of the basic rules that govern energy utilization.

By the end of this tutorial you should have a basic working understanding of:

  • What constitutes energy
  • The relationship between chemical bonds and energy
  • The first and second laws of thermodynamics
  • The free energy equation
  • The nature of ATP, and how it can be used to perform work

Metabolism and Energy

Figure 1. Metabolic Pathways in a Cell. (Click to enlarge)  This diagram depicts only a fraction of the metabolic pathways in a cell. Dots represent molecules, and lines represent chemical reactions.

This diagram depicts only a fraction of the metabolic pathways in a cell. Dots represent molecules, and lines represent chemical reactions. The straight line of large dots down the center is the catabolic breakdown of carbohydrates (glycolysis), and the large ring towards the bottom is the citric acid (or Krebs) cycle.

Metabolism refers to all of the chemical reactions that occur in an organism. This figure (left) maps some of the metabolic reactions involved in sugar metabolism. Metabolic reactions can be subdivided into those that result in the formation of molecules (anabolic) versus those that result in the breakdown of molecules (catabolic). Anabolism and catabolism were discussed in Tutorial 16. During photosynthesis, photoautotrophs (e.g., land plants) convert the energy found in sunlight into chemical energy via a series of anabolic reactions that result in starch being formed and stored within the plant. Chemoheterotrophs (e.g., humans) eat starch and break it down via a series of catabolic reactions to obtain the stored chemical energy. Before we discuss how these chemical reactions are used to move energy around the biosphere (and why), we need to define energy.

Energy refers to that which can or does move matter (i.e., the capacity for doing work). Let's consider two forms of energy. Kinetic energy refers to energy that is associated with moving matter. Potential energy refers to energy that is stored. Work is the act of moving matter, whether that matter is as small as an electron or as large as a whale.

Examples of potential and kinetic energy are all around us. For example, the Hoover Dam on the Colorado River(shown here) blocks the normal flow of water and stores potential energy behind its walls. When released, the kinetic energy of the swiftly flowing water is harnessed to provide electricity for 1.3 million people a year.

Figure 2.  The Hoover Dam.  (Click to enlarge)  

There is also a connection between kinetic energy and temperature. The temperature of an object is a reflection of the kinetic energy of its atoms or molecules. Fast molecules = high kinetic energy = high temperature. Boiling water has a much higher temperature, and subsequently a higher kinetic energy than ice.

Energy and Chemical Bonds

To understand how metabolic reactions are used to transduce energy for life's processes, one needs to appreciate some basics about the physical forces that allow two atoms to interact. Chemical bonds hold atoms together, and these bonds form as a result of electron behavior (either directly or indirectly). Electrons have mass, hence their movement requires energy. In addition, electrons contain varying amounts of potential energy. In the vast majority of metabolic reactions, it is the electrons that are the vehicle for moving energy. How does this happen?

Let's start with a common chemical reaction.

Methane (CH4) + 2 oxygen (O2) ----> Carbon dioxide (CO2) + 2 water (H2O)

This reaction takes place every time you turn on a gas stove. The reaction gives off energy in the form of heat and can be used to do work (e.g., boil water). In this reaction the chemical bonds associated with the four hydrogen atoms around the carbon atom must break, then reassociate with the electrons associated with oxygen. Note that it takes energy to break bonds; in this case, 632 kcal of energy are required for every mole of methane that reacts with two moles of oxygen.

If energy is required, how does this reaction evolve heat energy?

The answer is that bond formation produces heat. In this example, 792 kcal of energy evolve as the new bonds form in association with oxygen. The net gain is 792 - 632 = 160 kcal/mole of energy released by the reaction.

Before ending this topic, let's clarify some case-sensitive nomenclature. In everyday usage, energy is written in Calories (Cal, note the uppercase "C"), otherwise known as Nutritional Calories. 1 Calorie = 1,000 calories, or 1 kcal = 1 Calorie.

Laws of Thermodynamics: The First Law

Figure 3.  Heterotrophs need food for energy. (Click to enlarge)

When glucose is broken down, there is an accompanying release of energy. How do organisms use this energy? Or, better yet, why do they even need energy? These not so trivial questions require an understanding of some important general laws that govern energy utilization.

Thermodynamics is the physics of energy transformations that occur in a collection of matter. Formally, any collection of matter under thermodynamic scrutiny is defined as a "system." Bioenergetics is the area of thermodynamics that deals specifically with the energetic reactions that occur in an organism; energetically, an organism is a "system." There are a few laws that apply to energy (both biological and nonbiological). The laws of thermodynamics are actually quite simple, but have some far-reaching implications (not only to life, but to all interactions between energy and matter).

The first law of thermodynamics states that energy is neither created nor destroyed. In other words, the amount of energy in the universe is constant. This concept is straightforward but can be confusing. You might have heard someone say that plants "make energy from the sun." Plants actually convert solar energy (sunlight) into molecular potential energy, in the form of sugar. Remember, plants do not "make" energy. Photosynthesis converts light energy (a form of kinetic energy) into chemical energy. This first law could be considered "bookkeeping." It states that the energy used and released in any reaction must be balanced. Note that the energy in one molecule can be distributed into two or more other molecules.

Laws of Thermodynamics: The Second Law

The second law of thermodynamics deals with the ordering of matter and states that all energy-affected matter in the universe is becoming random. In other words, the total entropy (S) in the universe is increasing. There is a very important relationship between the movement of matter and the ordering of matter. All matter in the universe is tending toward disorder, however, systems can become ordered (become less entropic) with the input of energy.

How can any order be achieved if the tendency for matter in the universe is toward ever-increasing entropy? How can complex (and highly ordered) organisms exist at all? For that matter, how can there be any ordered states in the universe? The answer is that systems can become ordered as long as they are "open" to the universe.

Figure 4.  The Second Law of Thermodynamics - there is a tendency toward disorder. (Click to enlarge)

Think of holding 10 marbles in your hand. Your hand represents a system consisting of 10 marbles. Now, turn your hand over and let gravity (a form of energy from outside the system) act on these marbles; they not only fall out of your hand (converting potential energy into kinetic energy), but they also become relatively disordered. You can reorder these marbles back into your hand, however, it requires the input of energy from outside of the system (i.e., your body as you walk around picking them up). The ordering of the marbles back into your hand is done at the expense of entropy somewhere else in the universe. Keep in mind that life (i.e., each organism) exists as a system that is open to the universe, and ultimately the energy that organisms obtain is used, in a sense, to reverse entropic change. From a thermodynamics standpoint, if you stop eating you will die because there is no input of energy from outside your body (i.e., your system) to reverse the natural tendency of matter to disorder.

The first and second laws are intimately related, and later you will learn how this relationship can be quantitatively stated. In a general sense, life channels energy thermodynamically (via a complex series of energy transformations) for the purpose of decreasing entropy (creating order). For the majority of life forms on Earth, this energy ultimately comes from the sun. Indeed, all the order you see around you is a direct result of the small fraction of total solar radiation that is absorbed by the our planet.

There is a special relationship between temperature and entropy. The entropic state of a given system is proportional to temperature. At absolute zero (in degrees Kelvin), the entropic state of any system is zero. Although absolute zero has never been achieved (-273 deg Centigrade), the relationship between temperature and entropy is important.

Free Energy Embodies the First and Second Laws

Free energy (G) is the energy available (or required) to do work in a given system. If a given system releases free energy, then it can do work. Conversely, if it absorbs free energy, then work can be done on it.

Let's return to the example of marbles being held in a hand. We will define the system as the person holding the marbles and the marbles themselves. When the marbles are held, they are relatively ordered (they have low entropy) but unstable (simply opening the hand will cause a spontaneous change in the system). The potential energy of the marbles is also relatively high. When the hand is opened, however, several things happen. First, potential energy is converted to kinetic energy. Second, friction is produced and released in the form of heat as the marbles fall through the air. Third, the system becomes disordered as the marbles bounce around (entropy increases). If the person walks around and picks up the marbles, then the thermodynamic state changes again. First, the kinetic energy exerted by the person picking up the marbles is converted to gravitational potential energy as the marbles become elevated above the ground, and second, the marbles become more ordered (less entropic). Importantly, the ordered state can be restored, but energy (in the form of the person picking up the marbles) is required to order matter.

The change in free energy (delta G) is endergonic if energy enters the system, and exergonic if it leaves the system. Moreover, an exergonic reaction is unstable, has a negative delta G, and is therefore a spontaneous reaction.

Lastly, in this example one can see why energy flow is not 100% efficient; as the marbles fall through the air there is a production of frictional heat (which, in this example, does no useful work and represents waste). All energy transfers have some inefficiencies, which is why reactions do not transduce 100% of the available energy.

Organisms can only live at the expense of free energy (G). The free energy changes (delta G) associated with life's metabolic energy involve the movement of matter. This free energy comes from a series of metabolic reactions that result in work being done at the molecular level (the movement of electrons, atoms, or molecules). Recall the relationship above, between free energy and stability; a given reaction (a system) that has the potential to do a lot of work (release a lot of free energy) is inherently unstable; it typically has a low relative entropy and tends to change spontaneously to a more stable, disordered state. In fact, the concept of spontaneity actually defines whether free energy is made available to do work (or if work is required).

Free energy is more than a change in entropic state because each given system has a certain amount of total energy. However, not all of this total energy is available to do work. Free energy is a function of the total energy change of a system and the entropic change.

The Free Energy Equation

A reaction actually describes a change (delta) from one state to another. The change in free energy is denoted as delta G. If this value is negative, then free energy will leave the system, work can be done, and the reaction will occur spontaneously.

Enthalpy (H) is the total energy in a system. If this value is negative, then some energy (typically heat) will leave the system. If this value is positive, then energy will enter the system (typically heat will be absorbed from outside). An exothermic reaction has a negative delta H and will release heat, whereas an endothermic reaction has a positive delta H and will absorb heat.

The entropic state of the system is denoted S. If the reaction results in an increase in entropy, then this value is positive. If the reaction decreases in entropy, then this value is negative.

Free energy is also dependent on temperature. (Recall the relationship between temperature and entropy) The value "T" is given in degrees Kelvin. (To convert Centigrade to Kelvin, add 273.)

The relationship between these terms is expressed by the free energy equation:

  • delta G = delta H - T(delta S)

This equation reveals that not all of the energy stored in a system is available for work; free energy is less than the total energy of a system. The free energy concept can be used to determine whether a specific process or reaction will occur spontaneously. Reactions that occur spontaneously tend to give off heat (have a negative delta H, or change in heat) and result in the disordering of matter (have a positive delta S). However, reactions that absorb heat (have a positive delta H) can occur if there is a large increase in entropy. The free energy equation takes both of these factors into consideration. Keep in mind, the more negative the value of delta G, the more free energy released by the reaction and the more work that can be done. Conversely, as delta G becomes progressively more positive, the energy required for the reaction to proceed also increases.

Free Energy and Metabolic Reactions

When the energy in a system at the start of a reaction is greater than the energy at the end of the reaction, the delta G is negative and the reaction is an exergonic reaction. This means that energy is released and can be used to do work. For example, the reactions involved in breaking down glucose to retrieve energy during cellular respiration are exergonic. Cellular respiration (discussed in Tutorial 22 and Tutorial 23) releases energy that the cell can then use to do work. We can express this reaction as:

  • C6H12O6 + 6 O2 ----> 6 CO2 + 6 H2O

As you determined, this reaction has a ?G of -686 kcal/mol. For every 180 gms of glucose (1 mole), 686 kcals become available to do work.

On the other hand, endergonic reactions require an input of energy. Photosynthesis, the process by which plants convert CO2 and H2O into sugars, requires an input of energy from the sun, so it is an endergonic process overall. In other words, photosynthesis can be described as:

  • 6 CO2 + 6 H2O ----> C6H12O6

This reaction has a delta G of +686 kcal/mole and requires that work be done. This work is done via energy provided by the sun. Photosynthesis will be discussed in Tutorial 28 and Tutorial 29.

ATP and Cellular Work

 Figure 5.  The structure of ATP.  (Click to enlarge)

How does the cell use energy? The short answer is to do work. A cell does three types of work: mechanical (e.g., contracting muscle cells), transport (e.g., moving substances across the cell membrane), and chemical (e.g., non-spontaneous reactions between molecules; discussed in the next section). A major source of chemical energy for this work is adenosine triphosphate (ATP), which is illustrated in the figure above.

ATP is a 5-carbon sugar (ribose) attached to a nitrogenous base (adenine; recall our discussion of the nucleotides DNA and RNA) and a group of three phosphates. The three phosphates are the triphosphate component of adenosine triphosphate, and they are very unstable. This instability is due to the three negative charges that induce an intramolecular strain in one area of the molecule. Most reactions that involve ATP depend on the hydrolysis of the third phosphate to liberate the potential energy that can be used to do work. In addition to free energy that can do work, ADP (adenosine diphosphate) and inorganic phosphate (PO4; abbreviated Pi or a circled P) are released during hydrolysis. The overall reaction is summarized in the figure above. The amount of energy that is made available to do work during this process is variable (depending on a number of factors, including temperature).

Figure 6.  The hydrolysis of ATP releases energy that cells can use to do cell work.  (Click to enlarge)

How Does ATP Perform Work?

Figure 7. Energy coupling drives the conversion of glutamic acid to glutamine.  (Click to enlarge)

Energy is made available from the hydrolysis of ATP to ADP + Pi. This energy provides energy for endergonic reactions. The many mechanisms that different reactions use to obtain energy from ATP are remarkable. Generally they all involve tapping into high-energy electrons found associated with the terminal phosphate group. The use of an exergonic (energy-releasing) process to drive an endergonic (energy-requiring) process is called energy coupling. There are many ways that this can be done. Examine how energy coupling can be used to allow an energetically unfavorable (non-spontaneous) reaction to take place.

Glutamine (an amino acid) synthesis is an unfavorable reaction that can be described as:

  • Glu + NH3 ----> Glu-NH2

The delta G for this reaction is + 3.4 kcal/mol, therefore, it is a non-spontaneous reaction. However, this reaction does occur with the help of ATP. This is accomplished by coupling the removal of the terminal phosphate of ATP and the addition of ammonia to glutamic acid. In the first step, glutamic acid becomes an unstable phosphorylated intermediate when ATP is broken down. In the second step, ammonia displaces the phosphate group and glutamine is formed. This figure summarizes how ATP can be coupled to a reaction, such that an unfavorable reaction becomes exergonic.

The hydrolysis of ATP to ADP and Pi releases energy (-7.3 kcal/mol) in excess of the energy required (+3.4 kcal/mol) for the synthesis of glutamine. Therefore, the overall delta G for the two-step process is -3.9 kcal/mol.

This type of energy coupling procedure is extremely important, and it allows many endergonic anabolic reactions to occur in the cell.

A Cautionary Note

The field of thermodynamics has a novel history. It began in the 19th century when engineers were trying to understand how to make more efficient steam-driven engines. What was discovered in this area of science turned out to have an impact on other disciplines (a phenomenon frequently observed in science). The laws of thermodynamics apply to many situations that involve the interaction of matter and energy.

Perhaps in an effort to reach even greater levels of understanding, attempts have been made to apply the laws of thermodynamics to all situations. My kids once tried to talk me out of cleaning their room by citing the second law; it didn't work (but they did). There are even some popular books that have attempted to explain human societies using thermodynamics (e.g., all societies tend toward disorder unless energy is provided). These analogies can be attractive and can even be quite entertaining (e.g., search the internet for "Thermodynamics of Hell"). However, many of these efforts are flawed and unfounded. While reading, be watchful about how these laws can, and cannot, be reasonably applied.


Energy and its relationship with matter were examined in this tutorial. Life exists within the context of the physical rules of the universe. Life does not violate any of these rules; rather, these rules can be used to understand the functioning of life. It is important to understand what energy really is; briefly, it can be defined as that which does, or can, move matter. Matter can be things as small as electrons or as large as whales. Most of this tutorial dealt with movement of matter at the chemical level, but these first principles apply to all levels of organization.

The breakage and formation of chemical bonds at a subatomic level deals with the movement of electrons. Therefore, it is not surprising that these processes involve energetic considerations. Generally, energy is required to break bonds, but energy is released during bond reformation. In a chemical reaction with two reactants, bonds typically must be broken before new bonds form. For example, when methane combines with oxygen, carbon dioxide and water are formed. In this reaction, more energy is released as the bonds reform than was required to break the original bonds. This energy released is available to do work. You will learn that those chemical reactions in the cell that release energy are an important component of energy transduction and are associated with catabolic reactions.

Some reactions do not release energy. Rather, energy is required in order for these reactions to take place. In fact, anabolic reactions typically require energy (this is why food is required for growth). The chemicals in food are used to supply energy, which are used, in part, to build new biomolecules.

It is important to recognize the behavior of energy and matter in the universe. These two quantities are described by the laws of thermodynamics. The first law states that energy cannot be created or destroyed; rather, energy is conserved. The second law deals with matter and states that matter is becoming more disordered.

The two laws of thermodynamics can be related into a quantity known as free energy. The free energy equation is useful because it predicts whether a given reaction will occur spontaneously or non-spontaneously. If the change in free energy is negative, then the reaction will occur spontaneously. If the change in free energy is positive, then the reaction will be non-spontaneous. Virtually all anabolic reactions have an overall positive free energy, but nonetheless occur with the help of energy.

If a reaction requires energy, where does it come from? It depends on the species. Chemoheterotrophic organisms obtain energy from various high-energy carbon compounds. Photoautotrophic organisms obtain their energy from the sun. Ultimately most chemoheterotrophs depend on photoautotrophs for food, so most life on the planet depends on the sun as an energy source.

This concludes the first of three tutorials on energy and metabolism. This tutorial covered the concepts of energy, entropy, and thermodynamics. We also began examining ATP and its energetic relationship within the cell. Next we will begin our discussion of one way that ATP is produced.

Added by DENISE WOODWARD , last edited by STEPHEN WADE SCHAEFFER on Jun 10, 2009 10:59