MODERN BOTANY - P. RAVEN - 1990

SECTION II. ENERGY AND LIVING CELLS

CHAPTER 5. THE FLOW OF ENERGY

Life on Earth depends on the flow of energy generated by thermonuclear reactions in the core of the Sun (Fig. 5-1). The amount of energy radiated by the Sun toward Earth is 13 • 10²³ cal/year. This is indeed staggering to imagine. For example, the amount of energy striking Earth each day is roughly equivalent to a million atomic bombs similar to the one dropped on Hiroshima.

About 1/3 of the solar energy is immediately reflected back into space as light (much like moonlight). Of the remaining 2/3, the greater part is absorbed by Earth and converted into heat. Part of this absorbed thermal energy goes toward evaporating Water in the oceans, driving The formation of clouds and, consequently, rain and snow. Solar energy, acting in concert with other factors, is responsible for the movement of air and water masses, thereby shaping the diverse climates across Earth's surface.

Less than 1% of the solar energy reaching Earth is captured by plants and other photosynthetic organisms, sustaining the life processes of all PROKARYOTES AND EUKARYOTES inhabiting our planet. Living systems transform energy from one form to another, converting solar energy into chemical and mechanical energy (Fig. 5-2).

Class="center">Fig. 5-2. An example of biological Energy Flow. The radiant energy of sunlight originates from thermonuclear reactions taking place deep within the Sun. METABOLISM/14.html">Chloroplasts, present in all photosynthetic Eukaryotic Cells, absorb this energy and use it to convert water and carbon dioxide into CARBOHYDRATES such as glucose, starch, and other storage compounds. Oxygen is released into the environment as a byproduct of photosynthetic reactions. Cell/35.html">Mitochondria, Organelles found in all eukaryotic cells, break down these carbohydrates and store the released energy in the form of ATP molecules. This process, known as cellular Respiration, consumes oxygen and produces carbon dioxide and water, thus completing the cycle.

The concepts of the vital interrelationships between plants and animals, and between energy and life, arose relatively recently. They form an integral part of Thermodynamics, the science of energy exchange. In our energy-dependent world, the word "energy" is used constantly, making it surprising that it entered common usage only about 100 years ago. Let us first examine a few key Principles of Thermodynamics, and then look at the mechanism by which Enzymes catalyze many intracellular reactions.

Laws of Thermodynamics

Energy is an elusive concept, typically defined by the work it performs. Nearly 200 years ago, heat (the most thoroughly studied form of energy) was viewed as a separate, weightless substance called "caloric." An object was considered cold or hot depending on how much caloric it contained; when a cold object was placed near a hot one, caloric flowed from the hot body to the cold one. When a piece of metal was struck with a hammer, it grew warm as caloric was driven to the surface. Although the caloric theory was incorrect, the concept proved remarkably useful.

The First Law of Thermodynamics

In the final quarter of the 18th century, the rapid improvement of the steam engine greatly contributed to a shift in scientific understanding of The Nature of energy. Energy became associated with work, while heat and motion were recognized as forms of energy. On this foundation, the fundamental laws of thermodynamics were formulated.

The first law of thermodynamics is quite simple: "Energy can be transformed from one form to another, but can neither be created nor destroyed."

In machines, for instance, chemical energy (from coal or gasoline) is converted into thermal energy, which is then partially transformed into mechanical (kinetic) energy. Part of this energy is reconverted into heat through friction generated during movement, while another portion remains in the machine as waste. Unfortunately, the heat of an engine or steam engine, as well as the heat generated by friction or left over as waste, cannot perform work. It cannot turn gears because it dissipates into the surrounding environment, yet it remains part of the overall equation. Indeed, engineers have calculated that most of the energy produced by machines is randomly dissipated as heat; the efficiency, or coefficient of performance (COP), of most of them does not exceed 25%.

THE CONCEPT OF potential energy emerged during studies of engine efficiency. It was believed that a barrel of oil (142 L) or a ton of coal possessed a specific amount of potential energy, which could be expressed in terms of the heat released during their combustion.

The efficiency of converting potential energy into useful work depends on the type of energy-transforming system.

Although these principles were formulated for heat engines, they apply equally to other systems. For example, a boulder perched at the top of a hill possesses potential energy. With a slight nudge (activation energy), it rolls down the hill, its potential energy being converted into kinetic energy of motion and frictional heat. As mentioned earlier, water also possesses potential energy under certain conditions. When a large mass of water plunges from the top of a waterfall or dam, it can perform work by driving various mechanisms, such as those used for grinding grain. Thus, the potential energy of the water in this system is transformed into the kinetic energy of wheels and gears, as well as the heat generated by the movement of the water itself and the Rotation of the wheels and machinery.

Light and electricity are Other forms of energy. Light can be converted into electrical energy, and electrical energy can be converted into light, for instance, by passing a current through the tungsten filament of an incandescent light bulb.

The first law of thermodynamics states that during energy exchanges and interconversions, the energy of the reaction products plus The energy released in those reactions is always equal to the sum of the energies possessed by the initial reactants.

The Second Law of Thermodynamics

From a biological standpoint, the second law of thermodynamics is the most fascinating. It predicts the direction of all events involving energy exchange and is therefore known as the "arrow of time." One way to state the second law is as follows: "In all energy exchanges or transformations, barring any loss of energy or external input into the system, the potential energy of the final state will always be less than the potential energy of the initial state." The second law aligns with everyday experience. A rock will fall down a mountain, but it will never roll back up. Heat will flow from a warmer body to a colder one, but never in the reverse direction. Human cells can enzymatically utilize glucose to produce CO₂ and water, but because we cannot store solar energy the way plants do, we are unable to enzymatically synthesize glucose from carbon dioxide and water.

In any process that results in a decrease in potential energy, energy is released (otherwise, the first law would be violated). Such a process is called an exergonic reaction. According to the second law, only exergonic reactions can occur "spontaneously," meaning on their own. (The word "spontaneously" says nothing about the reaction rate, only whether or not it can occur.) Conversely, endergonic reactions require an input of energy—the greater the energy difference between the final product and the reactants, the larger the input required.

Whether a given reaction is exergonic or not is judged by the value of ΔH, which is The change in the heat content of the system (Δ meaning change, and H representing heat content). In short, the change in heat content is approximately equal to the change in potential energy. For example, the energy change that occurs during The oxidation of glucose can be measured in a calorimeter and expressed as ΔH. The oxidation of 1 mole of glucose yields 673 kcal. (A mole is the amount of a substance in grams equal to its molecular weight. For instance, the atomic mass of carbon is 12 and that of oxygen is 16, so the Molecular Weight of CO₂ is 44, and 1 mole of CO₂ equals 44 g.)

C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O

∆Н = — 673 kcal/mol.

The minus sign indicates that energy is released; in other words, a chemical reaction with a negative ∆H value is an exergonic reaction.

Another factor related to heat gain or loss can determine the direction of a process. This factor is Entropy, which serves as a measure of a system's disorder or randomness. For instance, the transition of water from a solid to a liquid and subsequently to a gaseous state (water vapor) involves endothermic processes that absorb a significant amount of heat from the environment. Under the right conditions, these processes occur spontaneously. The key driver in all such cases is the increase in entropy.

When ice melts, water transitions from a solid to a liquid state, breaking the bonds that hold water molecules within the ice crystals. When water evaporates, the remaining Hydrogen Bonds are broken, and individual molecules enter the vapor phase. In both instances, the disorder of the system increases.

In our daily lives, we observe that the greater the number of objects and the smaller their size, the higher the degree of disorder. The potential for disorder is greater if There are 20 sheets of paper on a desk compared to 2 or 10. If you cut 20 sheets of paper in half, the entropy of the system increases, thereby enhancing the potential for disorder. It is a general rule that a fundamental relationship exists between energy and entropy. If you need to tidy up a room or alphabetize books, you understand that someone must perform work and expend energy. Similarly, organizing papers on a desk requires a definite input of energy.

Fig. 5-3. Illustrations of the second law of thermodynamics. Stored energy—whether in the form of heated copper plates, pressurized gas molecules, or neatly stacked books—inevitably dissipates. In nature, all processes tend toward disorder or randomness. Only by applying energy can this trend be reversed and the original state restored. Ultimately, however, disorder will prevail because the total amount of energy in the universe is finite.

Let us now turn to the energy changes that govern the course of Chemical Reactions. As discussed earlier, both the change in the heat content of a system (∆H) and the change in entropy (∆S) contribute to the overall change in its energy state. This combined energy change is known as the change in Free energy and is denoted as ∆G in honor of the American physicist J. Willard Gibbs (1839–1903), who made profound contributions to The Development of thermodynamics.

Having established the concept of ∆G, let us return to the example of glucose oxidation. In this reaction, ∆H is — 673 kcal/mol, and ∆G is — 686 kcal/mol. Thus, the entropy factor contributes —13 kcal/mol to the free energy value. The combined effects of heat and entropy changes result in a lower energy state for the products of this reaction.

The magnitude of free energy ∆G and its change allow us to predict the nature of a process even when ∆H is zero or positive. We know that heat flows from a warmer body to a cooler one, and that ink molecules disperse evenly throughout a Glass of water. In each of these processes, the final state possesses higher entropy and consequently lower potential energy than the initial state.

The relationship between ∆G, ∆H, and entropy can be expressed

by the following equation:

∆G = ∆Н - Т∆S.

This equation indicates that the change in free energy equals the heat change (a negative value in exothermic reactions where heat is released) minus the entropy change, which is scaled by the absolute Temperature T.

In exergonic reactions, ∆G is always negative, whereas ∆H can be zero or positive. Because T is always positive, the larger the entropy change, the greater the magnitude of ∆G. This leads to an alternative formulation of the second law: "All natural processes are exergonic."

The Biological Law of Thermodynamics

The laws of thermodynamics are just as crucial to biology as they are to physics and chemistry. They provide a framework for understanding the principles underlying a vast array of biological processes and phenomena. Furthermore, as we will see, they enable a unique form of "bookkeeping" in biochemistry.

The most fascinating application of the second law to biology concerns the relationship between entropy on the one hand, and order and Organization on the other. Living systems continuously expend large amounts of energy to maintain order—that is, to preserve a state far from equilibrium. If equilibrium were reached, chemical reactions within The Cell would cease, and no work could be performed. At equilibrium, a cell would quickly perish.