Writing a book requires the help and patience of many colleagues. Thanks to our colleagues at Rose-Hulman and at Indiana University School of Medicine for their ideas and intellectual contributions, most especially, the authors of each chapter. Special thanks to Ellen Hughes, who made valuable comments on the manuscript.

Thanks to Lee's mother, Charlotte Waite, and Gabi's father, Werner Hess, who played an important role in making us who we are. If it is possible for either of us to write a chapter or to edit a book, that ability began at the knees of our parents when they taught us that reading and education are important. You deserve more credit for what is written in this book than you would admit or even realize.

Finally this book is dedicated to the memory of Margarete Hess, Gabi's mother, who died of pancreatic cancer at far too young an age. We miss you very much.

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Chapter 1


Walter X. Balcavage, Ph.D.


■ To understand the role of physical forces for chemical reactions

■ To introduce the specific biological role of water

■ To present the various forms of chemical bonds

■ To introduce the major categories of biomolecules


1.1 Energetics in Biology 2

1.2 Water 9

1.3 Amino Acids, Peptides, and Proteins 14

1.4 Carbohydrates and Their Polymers 20

1.5 Nucleic Acids, Nucleosides, and Nucleotides 24

1.6 Fats and Phospholipids 28

Biomolecules are the fundamental building blocks of all biological matter and are necessary for the existence of all known forms of life. The knowledge of the chemical structures of biological molecules will help in understanding their biological function and the energy flow in cells. Understanding biomolecules is important for the progress of molecular biotechnology, which aims to design new drugs or autonomous nanomachines that heal wounds and perform surgery.

This book is aimed at engineering students and professionals who have only a limited background in the biological sciences but who, through need or personal desire, want a broad exposure to the principles that form the basis of the biological sciences, including medicine. With this brief disclaimer, it should be clear to the reader that this book is not intended as a comprehensive treatise on any of these disciplines but rather as a venue by which the professional nonbiologist can obtain a working knowledge of the life process.

For professionals, the game of bridging the knowledge gap between scientific disciplines is a bit like tourists trying to bridge the gap between cultures. When the tourists are successful they find that they've accomplished one of the most rewarding tasks they've ever encountered. Similarly, when technical experts bridge knowledge gaps such as those between engineering disciplines and biology, the results can lead to exceedingly rewarding personal and professional results. One of the knowledge gaps alluded to is that of the language and syntax gap that is ubiquitous between scientific disciplines. In this regard, it is fortunate that learning the language of biology is no more difficult for the engineer, or scientist from another discipline, than that encountered by a tourist making their way in a foreign country. The difference, of course, is that in the biological sciences the building blocks of our knowledge comprise a well-defined set of atoms, molecules, and chemical reactions rather than letters, words, and sentences. Additionally, the words biological scientists use often have very arcane meanings compared to their conventional usage in everyday language. For example, the term free energy describes a kind of energy that is anything but free.

With this introduction, it is appropriate that the first chapter of this book should focus on the very fundamentals of the language of the biological sciences. As outlined in the following, we will begin by introducing the fundamental thermodynamic principles that help us understand the way in which the flow of energy enables the life process, and then we will go on to define and illustrate the basic molecular building blocks from which all biological structures are built.

1.1 Energetics in Biology 1.1.1 Thermodynamic principles

All chemical, physical, and biological processes are ultimately enabled and regulated by the laws of thermodynamics. Thus, to understand the life processes of cells and higher life forms, we need to develop a working knowledge of thermodynamics and then use this knowledge to understand how biological processes are enabled and regulated according to classical thermodynamic principles.

Classical thermodynamics involves a consideration of the energy content of different states of systems where each system is composed of a number of kinds of molecules or other objects and energy flows between components of the system and between the system and its environment with time. There are two basic kinds of thermodynamic systems: open and closed (Fig. 1.1). Open systems are characterized by a flow of matter (food and excreta in animals) and energy between the system (the body) and the environment. Examples of open systems include individual living cells and the human body, which is an aggregate of cells, and can be considered as an open thermodynamic system. In contrast, in a closed system, such as a bomb calorimeter, only energy is exchanged between the system and its environment. In this discussion of thermodynamic principles, we will review the first and second laws of thermodynamics focusing on their relationship to energy flow in living organisms.

The first law of thermodynamics states that the total energy of a system plus its environment remains constant. While not addressing the various forms in which energy can exist, this law declares that energy is neither created nor destroyed and it allows energy to be exchanged between a system and its surroundings. In closed system, like a bomb calorimeter, the only form of energy flow between the system and its environment is heat. Conversely in an open system, like an animal cell, or the human body, energy is most obviously exchanged into and out of the system in the form of heat and energy-rich, reduced carbon-containing molecules (e.g., sugars) and other matter (e.g., the respiratory molecules oxygen

Figure 1.1 Open and closed systems. In open systems mass and energy readily flow in and out of the system as illustrated by mass arrows and energy arrows penetrating the boundary of the open system. In closed systems energy (heat) moves in and out of the system but mass can neither move into the system nor out of the system.

Figure 1.1 Open and closed systems. In open systems mass and energy readily flow in and out of the system as illustrated by mass arrows and energy arrows penetrating the boundary of the open system. In closed systems energy (heat) moves in and out of the system but mass can neither move into the system nor out of the system.

and carbon dioxide). In animals, it is generally the case that matter flowing into the living system contains a high energy potential and matter flowing out of the system is at a lower energy potential. The energy changes that occur between these two mass flow events are used to perform chemical and physical work processes. Some of the work processes, such as pumping molecules from compartments of low concentration to compartments of high concentration and performing biosyntheses, result in some of the energy remaining stored in the body while the remainder is used to perform mechanical work or appears in the environment as a form of heat. In summary, the ingestion of food and excretion of metabolic products represent exchanges of mass with our environment and is a hallmark of an open thermodynamic system.

The process of consuming complex substances from our environment and excreting simpler breakdown products is also a reflection of the second law of thermodynamics. The second law of thermodynamics states that a system and its surroundings always proceed to a state of maximum disorder or maximum entropy, a state in which all available energy has been expended and no work can be performed. Entropy (S) and disorder are synonymous in thermodynamics. In the absence of the transfer of mass (food) from our surroundings into the human body, we soon starve, die, and disintegrate. In the case of plants, the photon energy from the sun powers photosynthesis, providing plants (and, as a consequence, humans) with high energy-potential, reduced-carbon compounds like sugars. In these examples, the plant and the animal systems remain viable as long as a usable form of energy input is available. The systems continue to expend the available potential energy until they proceed to a state of maximum entropy with death being one waypoint on the path to maximum system entropy.

In the conversion of complex foods such as glucose [C6(H2O)6] to simpler products such as CO2 and H2O, energy conversions, allowed by the first law of thermodynamics, take place.

It is these energy changes that are available to perform the chemical and physical work that keep us alive. This energy is known as Gibbs free energy (G) although it might have been more profitably termed usable energy since it certainly is not free but rather is available at the cost of an aging sun. The entropy change associated with glucose oxidation, or any similar reaction, is qualitatively reflected by a change in the ordered spatial relationship of atoms as biochemical reactants are converted to products. In our example, it should be clear that the atoms of glucose [C6(H2O)6] are much more highly structured than the product atoms in CO2 and H2O shown in Eq. (1.1). For any given state of a system the collective organization of the components of the system is related to its entropy. Simultaneously, as a consequence of the same chemical process, the heat content of the molecules, which is the sum of the heat associated with molecular collisions, the motion of bonding, and other electrons in the constituent atoms, also changes. This energy is known as enthalpy (H) and collectively, for a system in a given state, this heat energy is known as the enthalpy of the system.

1.1.2 Relationship between entropy (S), enthalpy (H), and free energy (E)

The quantitative relationship between the different forms of energy in a system, or in a reaction, going from one state to another is given by Eq. (1.2), the Gibbs equation, where A represents the quantitative difference in the energy forms G (Gibbs free energy), H (enthalpy), or S (entropy) between any two states of a system:

The Gibbs equation applies to all reactions and processes. A general example is the equilibrium equation (1.3):

In the Gibbs expression [Eq. (1.2)], where T is the Kelvin temperature of the system, it is clear that the magnitude of the entropic contribution to the free or usable energy is dependent on temperature (TAS). The enthalpy, or heat content of the system, is in principal also dependent on temperature, but in our biological world, and especially in the human body where reactions take place at constant temperature, molecular motions and collisions remain the same from one state of the system to the next. As a consequence, these kinds of contributions to enthalpy change are generally considered to be negligible. At constant temperature, the remaining and principal enthalpic source of energy is that associated with the chemical bonding between atoms in systems. This energy is known as internal energy (E) and in organic molecules it can be recognized as the covalent bonds, or forces, that stabilize atoms in the molecule. At constant temperature, enthalpy can be taken to be equal to internal energy, and thus the relationship between internal energy and enthalpy changes between two states can be expressed as shown in Eq. (1.4).

Equation (1.5) states that as a consequence of a reaction or process going from one constant temperature state to another, the available useful energy (AG) equals the difference between the changes in internal or bonding energy (AE) and the changes in organization (AS) of the atoms involved in the reaction.

The sign, + or —, and the magnitude of each term in the Gibbs equation is important in determining if a specified reaction or process described according to the Gibbs equation will proceed spontaneously. If TAS is positive and greater than AE, then AG will be negative and reactions such as the reaction in Eq. (1.3) will proceed spontaneously to the right as written. Reactions or processes having a negative AG are called exergonic. Reactions having a positive AG are called endergonic and these reactions require the input of energy in some form for the reaction to proceed in the direction written. For example, Eq. (1.1), the oxidation of glucose, can be written in the reverse direction, shown in Eq. (1.6), as a synthesis reaction in which CO2 and H2O are combined to form glucose.


However, notice that in this case we invoke the energy of sunlight (and implicitly all the photosynthetic machinery of a green plant) to reverse the entropic and enthalpic changes that result from oxidizing glucose. In biochemical systems, the free energy decrease of exergonic reactions is usually associated with a corresponding increase in entropy, although internal energy changes can also be important. Table 1.1 summarizes the preceding relationships.

TABLE 1.1 Relationship Between Thermodynamic Constants Keq, AG0, and the Terms Exergonic, Endergonic, and Spontaneous.

Exergonic reactions

Endergonic reactions

AG° is negative

Equilibrium constant (Keq) is greater than 1 Spontaneous as written

AG° is positive

Equilibrium constant (Keq) is less than 1

Spontaneous in the reverse of the direction written

The equilibrium constant (Keq) is described in Sec. 1.1.3, Eq (1.10). The terms exergonic, endergonic, and spontaneous are often used to describe the thermodynamic character of reactions. For example, if the reaction A ^ B has a negative AG° or an equilibrium constant greater than 1, it is called exergonic and will proceed to equilibrium spontaneously, as written, provided a reaction pathway is available.

To illustrate these thermodynamic relationships, we can consider more carefully the oxidation of 1 mol of glucose [Eq. (1.6)] where the initial state of the system/reaction is at the so-called standard concentration state and proceeds to the equilibrium concentration state under standard conditions of temperature and pressure according to Eq. (1.7).

In Eq. (1.7), the term AG0 has an added prime mark. However, recall that in classical physical chemistry AG° is the symbol for the standard free energy change of a system that proceeds from the standard chemical state (as defined in physical chemistry) to the equilibrium state. However, bio-scientists have defined a somewhat different set of standard state conditions that reflect the aqueous pH-neutral conditions under which the life process takes place. Thus, standard biological conditions are defined as 760 mm Hg (1 atm), a hydrogen ion concentration of 10"7 molar (M) (i.e., pH 7.0), 298 Kelvin (K), 55.5 M water, and 1 M concentration of all other reactants and products. The symbols for energy changes that take place going from these standard biological conditions to the equilibrium state are given a prime mark as indicated in Eq. (1.7) (i.e., AGO') to signify that they refer to reactions taking place at 55.5 M water, pH 7.0, and 298 K. Thus, the biochemical standard free energy (AG°') is that available as the reaction proceeds from the biological standard state (1 M glucose, 55.5 M water, 1 atm oxygen, 1 atm CO2, pH 7.0, 298 K) to the chemical equilibrium state under otherwise standard biological conditions.

The units of free energies are either calories per mole (cal/mol) or joules per mole (J/mol). Since calories and joules are both currently in common use, it is important to recall that 1 cal is equal to 4.184 J. Thus, the change in free energy, or AG°', for Eq. (1.7) is approximately -686,000 cal/mol, or -2,870,000 J/mol. The value of -686 kcal/mol (—2870 kJ/mol) for glucose oxidation is a large negative standard free energy change, which indicates that if a reaction mechanism or pathway is available, the reaction will proceed vigorously in the forward direction (as written). This is of course also the direction in which it proceeds in living organisms that possess abundant mechanistic pathways that enable organisms to utilize glucose as a source of energy. Standard free energy changes have been tabulated for most of the known biochemical reactions and can be found in many reference texts as well as on the Internet.

1.1.3 Entropy as driving force in chemical reactions

As a consequence of the preceding considerations, it follows that energy available from a reaction having an initial concentration state different than the standard concentration state will not be equal to that specified as AG°'. To evaluate the actual energy available from a reaction that is at other than the standard state, the free energy needs to be evaluated taking into account the prevailing reaction conditions. Thus, the free energy of a reaction when the initial reactant concentrations are other than 1 M (AG') is given by Eq. (1.8).


In this expression, the brackets signify that we mean concentration, in molarity, and R is the gas constant, 1.987 cal/molK, or 8.134 J/molK. An important observation related to Eq. (1.8) is that when the reaction is in the standard state, the ratio of reactants to products is 1, the log of 1 is 0 and thus AG' = AG°'!

Equation (1.8) reflects, in mathematical terms, the fact that the actual energy available from a reaction depends on its standard free energy plus (or minus) an energy contribution determined by the prevailing concentration of reactants and products. Clearly, in the human body, the concentrations of reactants and products for metabolic reactions are almost never that of the standard state; therefore, in vivo AG°' is almost never representative of the actual energy available from a metabolic reaction.

Reactions that are at equilibrium under biological standard state conditions of temperature, pressure, and pH cannot proceed spontaneously to any new biologically relevant state, and thus they cannot provide any useful biological work. Consequently, it can be said that the thermodynamic cause of death is that the reactions responsible for maintaining life have come to biological equilibrium. A more precise way of expressing these ideas is to note that at equilibrium, AG' = 0. As a consequence of this relationship, Eq. (1.8) can be algebraically modified to yield Eq. (1.9).


Since, at equilibrium, the ratio of [product]/[reactant] is equal to the equilibrium constant (Kgq), Eq. (1.9) is often written as shown in Eq. (1.10).

From Eq. (1.10), it is apparent that at 298 K, AG°' is equal to K^ multiplied by a collection of constants. In Eq. (1.11),we explicitly identify all the constants that impact Keq. and then simplify Eq. (1.11) as shown in Eq. (1.12).

joules degrees X mole

AG°' = -2.303 X 8.134 J _ „ , X 298 degrees X log Keq (1.11)

joules mole

Thus, at room temperature, or 25°C, the standard free energy of a biological reaction is simply —5.58 kJ/mol multiplied by the log of the equilibrium constant. The equivalent value for body temperature, 37°C or 310 K, is 5.80 kJ/mol.

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