Principles of Biochemistry Volume 2 - A. Lehninger 1985

Bioenergetics and Metabolism
The ATP Cycle and Cellular Bioenergetics
First and Second Laws of Thermodynamics

Today, people feel more keenly than ever how essential energy (i.e., the capacity to do work) is to sustaining our entire modern civilization. We need energy to manufacture various goods, transport people and Materials, heat homes and workplaces, and countless other vital tasks. Energy is just as indispensable to the microcosm of the living Cell. Living Cells continuously synthesize new substances, perform mechanical work associated with movement, transport molecules, and generate heat. Over billions of years of evolution, cells have learned to utilize energy much more economically and efficiently than most human-made machines. In fact, we now view living cells as models for designing new, more advanced energy-conversion devices—foremost among them systems for capturing solar energy.

The branch of biochemistry concerned with energy transformation and utilization in living cells is called bioenergetics. We will begin this chapter by examining several fundamental Principles of Thermodynamics, the branch of physics that deals with energy transformations. Following this, we will turn to The ATP system to see how it mediates The transfer of energy in cells from energy-yielding catabolic reactions to those cellular processes that require it.

Energy exists in various forms: we are familiar with electrical, mechanical, chemical, thermal, and radiant energy. We also know that energy can be converted from one form to another. For instance, an electric motor converts electrical energy into mechanical energy, a storage battery transforms chemical energy into electrical energy, and a steam turbine converts heat into mechanical energy. Different forms of energy are related by fixed quantitative ratios: for example, 1 cal of thermal energy is theoretically equivalent to 4.185∙10⁷ of mechanical energy.

It is well known, however, that any energy conversion is accompanied by certain losses. An electric motor converting electrical energy into mechanical work always yields less useful energy than it consumes because friction converts a portion of the energy into heat, which dissipates into the surroundings and can no longer be used. Practically every time energy is used to perform work or is transformed from one form to another, a fraction of the useful energy is lost. In many machines, less than 25% of the consumed energy goes into performing useful work. Numerous quantitative studies on the interconversion of various energy forms conducted by physicists and chemists have led to the formulation of two fundamental Laws of Thermodynamics. We will attempt to present their core essence here in the simplest and most accessible form.

a. The First Law

In any physical or chemical change, the total amount of energy in the universe remains constant.

The first law is the law of conservation of energy; it can also be stated as: energy is neither created nor destroyed. Whenever energy is used to perform work or is transformed from one form to another, the total amount of energy remains unchanged.

b. The Second Law

All Physical and Chemical processes tend to proceed in a direction that corresponds to the irreversible conversion of useful energy into a chaotic, disordered form. A measure of this conversion is a quantity known as Entropy. The process stops when an equilibrium state is reached, at which entropy attains the maximum possible value under the given conditions.

This simplified and somewhat Abstract formulation requires some explanation. First of all, it is necessary to define The concepts of "useful energy" and "entropy" more precisely. There are two kinds of useful energy: (1) Free energy, which can do work at constant Temperature and pressure, and (2) thermal energy, which can do work only when there is A change in temperature and pressure. Entropy is a quantitative characteristic or measure of the disordered (in a sense, useless) energy within a given system. A rigorous definition of entropy requires a mathematical Treatment of the concept of "disorder." Since we cannot do that here, let us use a few simple Examples to characterize METABOLISM/2.html">THE CONCEPT OF entropy qualitatively (Box 14.1).

Box 14-1. The Concept of Entropy

The term "entropy," which literally means "internal transformation" or "turning inward," was first introduced in 1851 by the German physicist Rudolf Clausius, who formulated one of the earliest statements of the second law of thermodynamics. A rigorous quantitative interpretation of entropy can be given based on statistical and probabilistic concepts. The qualitative meaning of this concept can be illustrated by three examples, each characterizing a specific aspect of entropy. The central idea always associated with entropy is system disorder, which may manifest in various ways depending on the case.

Case 1. The Teakettle and Heat Dissipation. It is well known that steam generated by boiling Water can perform useful work. Suppose, however, that as soon as the water temperature in the kettle (i.e., the "system") reaches 100 ˚C, we turn off the heat and simply let it cool in the kitchen (i.e., the "surrounding environment"). No work will be performed in the process. Instead, heat will flow from the kettle into the surroundings, gradually raising the temperature of the environment (i.e., the kitchen) until complete thermal equilibrium is finally reached. At that point, all parts of our kettle and the kitchen will be at practically the same temperature. The free

energy that was concentrated in the kettle when it was filled with water heated to 100 ˚C—and which had the potential to do work—has vanished. An equivalent amount of thermal energy remained in the "kettle + kitchen" system (i.e., the "universe") after the kettle cooled, but it became redistributed randomly—or, in other words, uniformly—among the various PARTS OF THE system. This energy is no longer available to do work because there are no longer any temperature gradients within the kitchen. Moreover, the increase in entropy in the kitchen (the "environment") caused by the cooling of the kettle is irreversible. Indeed, we know well from everyday experience that heat will never spontaneously flow backward—that is, from the kitchen to the cooled kettle—to reheat the water inside it to 100 ˚C.

Case 2. Glucose Oxidation. Entropy characterizes the state of matter as well as energy. Aerobic organisms extract free energy from glucose obtained from their environment. To harness this energy, they oxidize glucose with molecular oxygen, which is also taken up from the environment. The End products of glucose oxidative metabolism, CO₂ and H₂O, are returned to the environment. During this process, the entropy of the surroundings increases, while the Organism itself remains in a steady state, and the degree of its internal order does not change. The increase in entropy in this case is partly related to heat dissipation, but another kind of disorder also arises, as illustrated by the overall equation for glucose oxidation in living organisms: C₆H₁₂O₆ + 6O₂ → 6СО₂ + 6Н₂О. This process can be represented schematically as follows:

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The atoms that previously made up a single glucose molecule and six oxygen molecules—totaling seven molecules—became more evenly distributed As a result of the reaction, since twelve molecules (6СО₂ + 6Н₂О) are now formed from the original seven.

Whenever a chemical reaction increases the number of molecules, or whenever a solid substance (such as glucose) is converted into liquid or gaseous products whose molecules possess a greater number of degrees of freedom and can move more freely in space compared to the solid, the degree of molecular disorder increases and, consequently, entropy rises.

Case 3. Information and Entropy. In Shakespeare's *Julius Caesar* (Act IV, Scene 3), Brutus, upon learning that Mark Antony is advancing on him with his army, speaks the following lines:

There is a tide in the affairs of men,

Which taken at the flood, leads on to fortune;

Omitted, all the voyage of their life

Is bound in shallows and in miseries¹.

Here we have an information-rich message encoded using letters of the English alphabet; there are 125 of them in total. Beyond its literal meaning, these words carry a secondary, hidden significance. They reflect not only the complex sequence of events in the play, but also the author's thoughts on clashing interests, ambition, and the lust for power. One Senses Shakespeare's profound insight into human nature. Thus, the Amount of Information contained within them is very large.

Let us now imagine that the 125 letters making up this quotation are scattered in complete disorder, as shown here in the figure

The entire meaning was lost. In this form, these 125 letters carry virtually no information, yet their entropy is quite high. From this, we can conclude that information is a form of energy; it is sometimes referred to as “negative entropy.” Indeed, information theory—the branch of mathematics that underpins computer software logic—is closely tied to thermodynamic theory. Living organisms are highly ordered structures containing a colossal amount of information and, consequently, possessing very low entropy.

¹ There is a tide in the affairs of men,

Which, taken at the flood, leads on to fortune;

Omitted, all the voyage of their life

Is bound in shallows and in miseries.

(W. Shakespeare. The Complete Works.)

There is another aspect of the second law that must be considered to understand how this law operates, particularly in biological systems. First, let us introduce the concept of a reaction system, which refers to the set of substances involved in a given chemical or physical process.

Such a system might be, for example, an animal organism, a single cell, or two reacting compounds. Next, we must introduce the Concept of the surroundings, with which the reaction system can exchange energy. The combination of the reaction System and Its surroundings constitutes what we call the “Universe.”

∆С = ∆Н - Т∆S

(Fig. 14-1) and which, generally speaking, includes the Earth and outer space. Certain chemical or physical processes can, of course, occur in closed systems that are incapable of exchanging energy with their surroundings. However, in the real world, and especially in the biological world, the systems in which chemical and physical processes take place do exchange energy with their surroundings. We will soon see how crucial this distinction between the system and its surroundings is when discussing energy exchange.

Fig. 14-1. Schematic representation of a reaction system and its surroundings. In reactions proceeding at constant temperature and constant pressure, energy can be exchanged between the system and its surroundings, but such exchange must comply with the laws of thermodynamics. The first law states that the total amount of energy in the “Universe” (system + surroundings) remains constant. According to the second law, any physical or chemical change within a system increases the entropy of the Universe; simultaneously, the Free energy of the reaction system decreases. Along with these changes, heat may be transferred from the system to the surroundings or vice versa, as expressed by the relation

Changes in free energy, heat, and entropy in Chemical Reactions occurring at constant temperature and constant pressure—conditions characteristic of biological systems—are quantitatively related by the following equation:

∆С = ∆Н — T∆S,  (1)

where ∆С is The change in free energy of the reaction system, ∆Н is the change in its heat content, or enthalpy (from the Greek enthalpein, “to warm”), Т is the absolute temperature at which the process takes place, and ∆S is the change in entropy of the “Universe,” which includes both the system and its surroundings. As a chemical reaction approaches equilibrium, the entropy of the Universe (system + surroundings) increases. Therefore, the value of AS in the real world is always positive. In principle, within some ideal system, a reaction might proceed without an increase in entropy. According to Equation (1), any increase in the entropy of the Universe during a reaction must be matched by a decrease in the free energy of the reaction system. Consequently, the value of AG for the reaction system is always negative. The change in enthalpy ∆Н is defined as The amount of heat that the reaction system releases to or absorbs from the surroundings at constant temperature and constant pressure. If the reaction system loses (i.e., releases) heat, ∆Н is negative; if the system absorbs heat from the surroundings, ∆Н is positive.

Another vital feature of entropy changes applies specifically to biological systems. According to the second law of thermodynamics, chemical reactions or physical processes increase the entropy of the Universe. This law does not imply, however, that the increase in entropy must necessarily occur within the reaction system itself; it can take place anywhere else in the Universe. In living organisms, metabolic processes—the chemical transformations that nutrients undergo—do not lead to an increase in the internal disorder, or entropy, of the organisms themselves. Everyday observation tells us that any organism, whether a fly or an elephant (i.e., our definition of a “system”), maintains its complex and ordered Structure throughout all life processes. As a result of life processes, it is the entropy of the surroundings that increases, not that of the living organisms themselves. Living organisms preserve their internal order by acquiring free energy in the form of nutrients (or sunlight) from their surroundings and returning an equivalent amount of energy in a less useful form, primarily as heat, which is dissipated into the rest of the Universe.

In Conclusion, it should be emphasized that the increase in entropy, or rising degree of disorder, is not entirely futile. Because the increase in the entropy of the Universe during biological processes is irreversible, it provides the driving force and sets the direction for all forms of biological activity. Living organisms continuously increase the entropy of their surroundings, and this is the price the Universe pays to maintain their internal order.

Here, it is probably useful to examine a specific cellular chemical reaction to gain a sense of the magnitude of energy changes involving different forms of energy. In aerobic cells, glucose (С₆Н₁₂О₆) is oxidized to СО₂ and Н₂О at constant temperature and constant pressure:

С₆Н₁₂О₆ + 6О₂→ 6СО₂ + 6Н₂О.

Assuming this reaction takes place under standard conditions—which in thermodynamic calculations means a temperature of 25 ˚С (or 298 K) and a pressure of 1 atm (760 mm Hg)—then for every mole of glucose oxidized,

The increase in molecular disorder, or entropy, that accompanies The oxidation of glucose can be readily visualized using the example given in Box 14.1.