Analysis of Situations Encountered During the Storage and Processing of Food

of food. The major cellular pools of carbohydrates, lipids, proteins, and their intermediary metabolites are shown on the lefthand side of the diagram. The exact nature of these pools is dependent on the physiological state of the tissue at the time of processing or storage, and the constituents present in or added to nontissue foods. Each class of compound can undergo its own characteristic type of deterioration. Noteworthy is the role that carbonyl compounds play in many deterioration processes. They arise mainly from lipid oxidation and carbohydrate degradation, and can lead to the destruction of nutritional value, to off-colors, and to off-flavors. Of course these same reactions lead to desirable flavors and colors during the cooking of many foods. 1.3.4 Analysis of Situations Encountered During the Storage and Processing of Food Having before us a description of the attributes of high-quality, safe foods, the significant chemical reactions involved in the deterioration of food, and the relationship between the two, we can now begin to consider how to apply this information to situations encountered during the storage and processing of food. The variables that are important during the storage and processing of food are listed in Table 4. Temperature is perhaps the most important of these variables because of its broad influence on all types of chemical reactions. The effect of temperature on an individual reaction can be estimated from the Arrhenius equation, k = Ae -DE/RT . Data conforming to the Arrhenius equation yield a straight line when logk is plotted versus 1/T. Arrhenius plots in Figure 2 represent reactions important in food deterioration. It is evident that food reactions generally conform to the Arrhenius relationship over a limited intermediate temperature range but that deviations from this relationship can occur at high or low temperatures [21]. Thus, it is important to remember that the Arrhenius relationship for food systems is valid only over a range of temperature that has been experimentally verified. Deviations from the Arrhenius relationship can occur because of the following events, most of which are induced by either high or low temperatures: (a) enzyme activity may be lost, (b) the reaction pathway may change or may be influenced by a competing reaction(s), (c) the physical state of the system may change (e.g., by freezing), or (d) one or more of the reactants may become depleted. Another important factor in Table 4 is time. During storage of a food product, one frequently wants to know how long the food can be expected to retain a specified level of quality. Therefore, one is interested in time with respect to the integral of chemical and/or microbiological changes that occur during a specified storage period, and in the way these changes combine to determine a specified storage life for the product. During processing, one is often Pag e 11 FIGURE 2 Conformity of important deteriorative reactions in food to the Arrhenius relationship. (a) Above a certain value of T there may be deviations from linearity due to a chang e in the path of the reaction. (b) As the temperature is lowered below the freezing point of the system, the ice phase (essentially pure) enlarg es and the fluid phase, which contains all the solutes, diminishes. This concentration of solutes in the unfrozen phase can decrease reaction rates (supplement the effect of decreasing temperature) or increase reaction rates (oppose the effect of declining temperature), depending on the nature of the system (see Chap. 2). (c) For an enzymic reaction there is a temperature in the vicinity of the freezing point of water where subtle chang es, such as the dissociation of an enzyme complex, can lead to a sharp decline in reaction rate. interested in the time it takes to inactivate a particular population of microorganisms or in how long it takes for a reaction to proceed to a specified extent. For example, it may be of interest to know how long it takes to produce a desired brown color in potato chips during frying. To accomplish this, attention must be given to temperature change with time, that is, the rate of temperature change (dT/dt). This relationship is important because it determines the rate at which microorganisms are destroyed and the relative rates of competing chemical reactions. The latter is of interest in foods that deteriorate by more than one means, such as lipid oxidation and nonenzymic browning. If the products of the browning reaction are antioxidants, it is important to know whether the relative rates of these reactions are such that a significant interaction will occur between them. Pag e 12 Another variable, pH, influences the rates of many chemical and enzymic reactions. Extreme pH values are usually required for severe inhibition of microbial growth or enzymic processes, and these conditions can result in acceleration of acid- or basecatalyzed reactions. In contrast, even a relatively small pH change can cause profound changes in the quality of some foods, for example, muscle. The composition of the product is important since this determines the reactants available for chemical transformation. Particularly important from a quality standpoint is the relationship that exists between composition of the raw material and composition of the finished product. For example, (a) the manner in which fruits and vegetables are handled postharvest can influence sugar content, and this, in turn, influences the degree of browning obtained during dehydration or deep-fat frying. (b) The manner in which animal tissues are handled postmortem influences the extents and rates of glycolysis and ATP degradation, and these in turn can influence storage life, water-holding capacity, toughness, flavor, and color. (c) The blending of raw materials may cause unexpected interactions; for example, the rate of oxidation can be accelerated or inhibited depending on the amount of salt present. Another important compositional determinant of reaction rates in foods is water activity (aw). Numerous investigators have shown aw to strongly influence the rate of enzyme-catalyzed reactions [2], lipid oxidation [16,22], nonenzymic browning [10,16], sucrose hydrolysis [23], chlorophyll degradation [17], anthocyanin degradation [11], and others. As is discussed in Chapter 2, most reactions tend to decrease in rate below an aw corresponding to the range of intermediate moisture foods (0.75–0.85). Oxidation of lipids and associated secondary effects, such as carotenoid decoloration, are exceptions to this rule; that is, these reactions accelerate at the lower end of the aw scale. More recently, it has become apparent that the glass transition temperature (Tg) of food and the corresponding water content (Wg) of the food at Tg are causatively related to rates of diffusion-limited events in food. Thus, Tg and Wg have relevance to the physical properties of frozen and dried foods, to conditions appropriate for freeze drying, to physical changes involving crystallization, recrystallization, gelatinization, and starch retrogradation, and to those chemical reactions that are diffusion-limited (see Chap. 2). In fabricated foods, the composition can be controlled by adding approved chemicals, such as acidulants, chelating agents, flavors, or antioxidants, or by removing undesirable reactants, for example, removing glucose from dehydrated egg albumen. Composition of the atmosphere is important mainly with respect to relative humidity and oxygen content, although ethylene and CO2 are also important during storage of living plant foods. Unfortunately, in situations where exclusion of oxygen is desirable, this is almost impossible to achieve completely. The detrimental consequences of a small amount of residual oxygen sometimes become apparent during product storage. For example, early formation of a small amount of dehydroascorbic acid (from oxidation of ascorbic acid) can lead to Maillard browning during storage. For some products, exposure to light can be detrimental, and it is then appropriate to package the products in light-impervious material or to control the intensity and wavelengths of light, if possible. Food chemists must be able to integrate information about quality attributes of foods, deteriorative reactions to which foods are susceptible, and the factors governing kinds and rates of these deteriorative reactions, in order to solve problems related to food formulation, processing, and storage stability. Pag e 13 1.4 Societal Role of Food Chemists 1.4.1 Why Should Food Chemists Become Involved in Societal Issues? Food chemists, for the following reasons, should feel obligated to become involved in societal issues that encompass pertinent technological aspects (technosocietal issues). • Food chemists have had the privilege of receiving a high level of education and of acquiring special scientific skills, and these privileges and skills carry with them a corresponding high level of responsibility. • Activities of food chemists influence adequacy of the food supply, healthfulness of the population, cost of foods, waste creation and disposal, water and energy use, and the nature of food regulations. Because these matters impinge on the general welfare of the public, it is reasonable that food chemists should feel a responsibility to have their activities directed to the benefit of society. • If food chemists do not become involved in technosocietal issues, the opinions of others—scientists from other professions, professional lobbyists, persons in the news media, consumer activists, charlatans, antitechnology zealots—will prevail. Many of these individuals are less qualified than food chemists to speak on food-related issues, and some are obviously unqualified. 1.4.2 Types of Involvement The societal obligations of food chemists include good job performance, good citizenship, and guarding the ethics of the scientific community, but fulfillment of these very necessary roles is not enough. An additional role of great importance, and one that often goes unfulfilled by food chemists, is that of helping determine how scientific knowledge is interpreted and used by society. Although food chemists and other food scientists should not have the only input to these decisions, they must, in the interest of wise decision making, have their views heard and considered. Acceptance of this position, which is surely indisputable, leads to the obvious question, “What exactly should food chemists do to properly discharge their responsibilities in this regard?” Several activities are appropriate. 1. Participate in pertinent professional societies. 2. Serve on governmental advisory committees, when invited. 3. Undertake personal initiatives of a public service nature. The latter can involve letters to newspapers, journals, legislators, government regulators, company executives, university administrators, and others, and speeches before civic groups. The major objectives of these efforts are to educate and enlighten the public with respect to food and dietary practices. This involves improving the public’s ability to intelligently evaluate information on these topics. Accomplishing this will not be easy because a significant portion of the populace has ingrained false notions about food and proper dietary practices, and because food has, for many individuals, connotations that extend far beyond the chemist’s narrow view. For these individuals, food may be an integral part of religious practice, cultural heritage, ritual, social symbolism, or a route to physiological wellbeing—attitudes that are, for the most part, not conducive to acquiring an ability to appraise foods and dietary practices in a sound, scientific manner. One of the most contentious food issues, and one that has eluded appraisal by the Pag e 14 public in a sound, scientific manner, is the use of chemicals to modify foods. “Chemophobia,” the fear of chemicals, has afflicted a significant portion of the populace, causing food additives, in the minds of many, to represent hazards inconsistent with fact. One can find, with disturbing ease, articles in the popular literature whose authors claim the American food supply is sufficiently laden with poisons to render it unwholesome at best, and life-threatening at worst. Truly shocking, they say, is the manner in which greedy industrialists poison our foods for profit while an ineffectual Food and Drug Administration watches with placid unconcern. Should authors holding this viewpoint be believed? It is advisable to apply the following criteria when evaluating the validity of any journalistic account dealing with issues of this kind. • Credibility of the author. Is the author, by virtue of formal education, experience, and acceptance by reputable scientists, qualified to write on the subject? The writer should, to be considered authoritative, have been a frequent publisher of articles in respected scientific journals, especially those requiring peer review. If the writer has contributed only to the popular literature, particularly in the form of articles with sensational titles and “catch” phrases, this is cause to exercise special care in assessing whether the presentation is scholarly and reliable. • Appropriateness of literature citations. A lack of literature citations does not constitute proof of irresponsible or unreliable writing, but it should provoke a feeling of moderate skepticism in the reader. In trustworthy publications, literature citations will almost invariably be present and will direct the reader to highly regarded scientific publications. When “popular” articles constitute the bulk of the literature citations, the author’s views should be regarded with suspicion. • Credibility of the publisher. Is the publisher of the article, book, or magazine regarded by reputable scientists as a consistent publisher of high-quality scientific materials? If not, an extra measure of caution is appropriate when considering the data. If information in poison-pen types of publications are evaluated by rational individuals on the basis of the preceding criteria, such information will be dismissed as unreliable. However, even if these criteria are followed, disagreement about the safety of foods still occurs. The great majority of knowledgeable individuals support the view that our food supply is acceptably safe and nutritious and that legally sanctioned food additives pose no unwarranted risks [7,13,15,19,24–26]. However, a relatively small group of knowledgeable individuals believes that our food supply is unnecessarily hazardous, particularly with regard to some of the legally sanctioned food additives, and this view is most vigorously represented by Michael Jacobson and his Center for Science in the Public Interest. This serious dichotomy of opinion cannot be resolved here, but information provided in Chapter 13 will help undecided individuals arrive at a soundly based personal perspective on food additives, contaminants in foods, and food safety. In summary, scientists have greater obligations to society than do individuals without formal scientific education. Scientists are expected to generate knowledge in a productive, ethical manner, but this is not enough. They should also accept the responsibility of ensuring that scientific knowledge is used in a manner that will yield the greatest benefit to society. Fulfillment of this obligation requires that scientists not only strive for excellence and conformance to high ethical standards in their day-to-day professional activities, but that they also develop a deep-seated concern for the well-being and scientific enlightenment of the public. Pag e 15 References 1. Accum, F. (1966). A Treatise on Adulteration of Food, and Culinary Poisons, 1920, Facsimile reprint by Mallinckrodt Chemical Works, St. Louis, MO. 2. Acker, L. W. (1969). Water activity and enzyme activity. Food Technol. 23(10):1257–1270. 3. Anonymous (1831). Death in the Pot. Cited by Filby, 1934 (Ref. 12). 4. Beaumont, W. (1833). Experiments and Observations of the Gastric Juice and the Physiology of Digestion, F. P. Allen, Plattsburgh, NY. 5. Browne, C. A. (1944). A Source Book of Agricultural Chemistry. Chronica Botanica Co., Waltham, MA. 6. Chevreul, M. E. (1824). Considérations générales sur l’analyse organique et sur ses applications. Cited by Filby, 1934 (Ref. 12). 7. Clydesdale, F. M., and F. J. Francis (1977). Food, Nutrition and You, Prentice-Hall, Englewood Cliffs, NJ. 8. Davy, H. (1813). Elements of Agricultural Chemistry, in a Course of Lectures for the Board of Agriculture, Longman, Hurst, Rees, Orme and Brown, London, Cited by Browne, 1944 (Ref. 5). 9. Davy, H. (1936). Elements of Agricultural Chemistry, 5th ed., Longman, Rees, Orme, Brown, Green and Longman, London. 10. Eichner, K., and M. Karel (1972). The influence of water content and water activity on the sugar-amino browning reaction in model systems under various conditions. J. Agric. Food Chem. 20(2):218–223. 11. Erlandson, J. A., and R. E. Wrolstad (1972). Degradation of anthocyanins at limited water concentration. J. Food Sci. 37(4):592–595. 12. Filby, F. A. (1934). A History of Food Adulteration and Analysis, George Allen and Unwin, London. 13. Hall, R. L. (1982). Food additives, in Food and People (D. Kirk and I. K. Eliason, eds.), Boyd and Fraser, San Francisco, pp. 148–156. 14. Ihde, A. J. (1964). The Development of Modern Chemistry, Harper and Row, New York. 15. Jukes, T. H. (1978). How safe is our food supply? Arch. Intern. Med. 138:772–774. 16. Labuza, T. P., S. R. Tannenbaum, and M. Karel (1970). Water content and stability of low-moisture and intermediatemoisture foods. Food Techol. 24(5):543–550. 17. LaJollo, F., S. R. Tannenbaum, and T. P. Labuza (1971). Reaction at limited water concentration. 2. Chlorophyll degradation. J. Food Sci. 36(6):850–853. 18. Liebig, J. Von (1847). Researches on the Chemistry of Food, edited from the author’s manuscript by William Gregory; Londson, Taylor and Walton, London. Cited by Browne, 1944 (Ref. 5). 19. Mayer, J. (1975). A Diet for Living, David McKay, Inc., New York. 20. McCollum, E. V. (1959). The history of nutrition. World Rev. Nutr. Diet. 1:1–27. 21. McWeeny, D. J. (1968). Reactions in food systems: Negative temperature coefficients and other abnormal temperature effects. J. Food Technol. 3:15–30. 22. Quast, D. G., and M. Karel (1972). Effects of environmental factors on the oxidation of potato chips. J. Food Sci. 37(4):584–588. 23. Schoebel, T., S. R. Tannenbaum, and T. P. Labuza (1969). Reaction at limited water concentration. 1. Sucrose hydrolysis. J. Food Sci. 34(4):324–329. 24. Stare, F. J., and E. M. Whelan (1978). Eat OK—Feel OK, Christopher Publishing House, North Quincy, MA. 25. Taylor, R. J. (1980). Food Additives, John Wiley & Sons, New York. 26. Whelan, E. M. (1993). Toxic Terror, Prometheus Books, Buffalo, NY Water and Ice OWEN R. FENNEMA University of Wisconsin—Madison, Madison, Wisconsin Prologue: Water—The Deceptive Matter of Life and Death 18 2.1 Introduction 18 2.2 Physical Properties of Water and Ice 19 2.3 The Water Molecule 20 2.4 Association of Water Molecules 22 2.5 Structure of Ice 24 2.5.1 Pure Ice 24 2.5.2 Ice in the Presence of Solutes 29 2.6 Structure of Water 29 2.7 Water-Solute Interactions 30 2.7.1 Macroscopic Level (Water Binding, Hydration, and water Holding Capacity) 30 2.7.2 Molecular Level: General Comments 31 2.7.3 Molecular Level: Bound Water 31 2.7.4 Interaction of Water with Ions and Ionic Groups 32 2.7.5 Interaction of Water with Neutral Groups Possessing HydrogenBonding Capabilities (Hydrophilic Solutes) 33 2.7.6 Interaction of Water with Nonpolar Substances 35 2.7.7 Details of Water Orientation Adjacent to Organic Molecules 37 2.7.8 Hydration Sequence of a Protein 38 2.8 Water Activity and Relative Vapor Pressure 42 2.8.1 Introduction 42 2.8.2 Definition and Measurement 42 2.8.3 Temperature Dependence 44 2.9 Moisture Sorption Isotherms 47 2.9.1 Definition and Zones 47 2.9.2 Temperature Dependence 50 2.9.3 Hysteresis 50 Pag e 18 2.10 Relative Vapor Pressure and Food Stability 52 2.11 Molecular Mobility (Mm) and Food Stability 55 2.11.1 Introduction 55 2.11.2 State Diagrams 57 2.11.3 Nine Key Concepts Underlying the Molecular Mobility Approach to Food Stability 59 2.11.4 Technological Aspects: Freezing 72 2.11.5 Technological Aspects: Air Drying 79 2.11.6 Technological Aspects: Vacuum Freeze-Drying (Lyophilization) 79 2.11.7 Technological Aspects: Other Applications of the Mm Approach (Partial Listing) 80 2.11.8 Technological Aspects: Estimation of Relative Shelf Life 80 2.11.9 Technological Aspects: Relationship of Tg and Mm to Relative Vapor Pressure (p/p0) and Moisture Sorption Isotherms 80 2.11.10 Summary Statements Regarding the Mm Approach to Food Stability 82 2.12 Combined Methods Approach to Food Stability 83 2.13 Concluding Comments About Water 85 Glossary: Molecular Mobility and Food Stability 85 Acknowledgments 87 Abbreviations and Symbols 87 Bibliography 88 References 88 Prologue: Water—The Deceptive Matter of Life and Death Unnoticed in the darkness of a subterranean cavern, a water droplet trickles slowly down a stalactite, following a path left by countless predecessors, imparting, as did they, a small but almost magical touch of mineral beauty. Pausing at the tip, the droplet grows slowly to full size, then plunges quickly to the cavern floor, as if anxious to perform other tasks or to assume different forms. For water, the possibilities are countless. Some droplets assume roles of quiet beauty—on a child’s coat sleeve, where a snowflake of unique design and exquisite perfection lies unnoticed; on a spider’s web, where dew drops burst into sudden brilliance at the first touch of the morning sun; in the countryside, where a summer shower brings refreshment; or in the city, where fog gently permeates the night air, subduing harsh sounds with a glaze of tranquility. Others lend themselves to the noise and vigor of a waterfall, to the overwhelming immensity of a glacier, to the ominous nature of an impending storm, or to the persuasiveness of a tear on a woman’s cheek. For others the role is less obvious but far more critical. There is life—initiated and sustained by water in a myriad of subtle and poorly understood ways—or death inevitable, catalyzed under special circumstances by a few hostile crystals of ice; or decay at the forest’s floor, where water works relentlessly to disassemble the past so life can begin anew. But the form of water most familiar to humans is none of these; rather, it is simple, ordinary, and uninspiring, unworthy of special notice as it flows forth in cool abundance from a household tap. “Humdrum,” galunks a frog in concurrence, or so it seems as he views with stony indifference the watery milieu on which his very life depends. Surely, then, water’s most remarkable feature is deception, for it is in reality a substance of infinite complexity, of great and unassessable importance, and one that is endowed with a strangeness and beauty sufficient to excite and challenge anyone making its acquaintance. 2.1 Introduction On this planet, water is the only substance that occurs abundantly in all three physical states. It is the only common liquid and is the most widely distributed pure solid, being ever present Pag e 19 TABLE 1 W ater Contents of Various Foods Food W ater content (%) Meat Pork, raw, composite of lean cuts 53–60 Beef, raw, retail cuts 50–70 Chicken, all classes, raw meat without skin 74 Fish, muscle proteins 65–81 Fruit Berries, cherries, pears 80–85 Apples, peaches, orang es, g rapefruit 90–90 Rhubarb, strawberries, tomatos 90–95 Veg etables Avocado, bananas, peas (g reen) 74–80 Beets, broccoli, carrots, potatoes 85–90 Asparag us, beans (g reen), cabbag e, cauliflower, lettuce 90–95 somewhere in the atmosphere as suspended ice particles, or on the earth’s surface as various types of snow and ice. It is essential to life: as an important governor of body temperature, as a solvent, as a carrier of nutrients and waste products, as a reactant and reaction medium, as a lubricant and plasticizer, as a stabilizer of biopolymer conformation, as a likely facilitator of the dynamic behavior of macromolecules, including their catalytic (enzymatic) properties, and in other ways yet unknown. It is truly remarkable that organic life should depend so heavily on this small inorganic molecule, and, perhaps even more remarkable, that so few scientists are aware of this fact. Water is the major component of many foods, each having its own characteristic allotment of this component (Table 1). Water in the proper amount, location, and orientation profoundly influences the structure, appearance, and taste of foods and their susceptibility to spoilage. Because most kinds of fresh foods contain large amounts of water, effective forms of preservation are needed if long-term storage is desired. Removal of water, either by conventional dehydration or by separation locally in the form of pure ice crystals (freezing), greatly alters the native properties of foods and biological matter. Furthermore, all attempts (rehydration, thawing) to return water to its original status are never more than partially successful. Ample justification exists, therefore, to study water and ice with considerable care. 2.2 Physical Properties of Water and Ice As a first step in becoming familiar with water, it is appropriate to consider its physical properties as shown in Table 2. By comparing water’s properties with those of molecules of similar molecular weight and atomic composition (CH4, NH3, HF, H2S, H2Se, H2Te) it is possible to determine if water behaves in a normal fashion. Based on this comparison, water is found to melt and boil at unusually high temperatures; to exhibit unusually large values for surface tension, permittivity (dielectric constant), heat capacity, and heats of phase transition (heats of fusion, vaporization, and sublimation); to have a moderately low value for density; to exhibit an unusual attribute of expanding upon solidification; and to possess a viscosity that in light of the foregoing oddities, is surprisingly normal. In addition, the thermal conductivity of water is large compared to those of other liquids, and the thermal conductivity of ice is moderately large compared to those of other nonmetallic Pag e 20 TABLE 2 Physical Properties of W ater and Ice Property Value Molecular weig ht 18.0153 Phase transition properties Melting point at 101.3 k Pa (1 atm) 0.000°C Boiling point at 101.3 k Pa (1 atm) 100.000°C Critical temperature 373.99°C Critical pressure 22.064 MPa (218.6 atm) Triple point 0.01°C and 611.73 Pa (4.589 mm Hg ) Enthalpy of fusion at 0°C 6.012 kJ (1.436 kcal)/mol Enthalpy of vaporization at 100°C 40.657 kJ (9.711 kcal)/mol Enthalpy of sublimination at 0°C 50.91 kJ (12.16 kcal)/mol Temperature Other properties 20°C 0°C 0°C (ice) -20°C (ice) Density (g /cm3 0.99821 0.99984 0.9168 0.9193 Viscosity (pa·sec) 1.002×10-3 1.793×10-3 — — Surface tension ag ainst air (N/m) 72.75×10-3 75.64×10-3 — — Vapor pressure (kPa) 2.3388 0.6113 0.6113 0.103 Heat capacity (J/g ·K) 4.1818 4.2176 2.1009 1.9544 Thermal conductivity (liquid) (W /m·K) 0.5984 0.5610 2.240 2.433 Thermal diffusitity (m2 /s) 1.4×10-7 1.3×10-7 11.7×10-7 11.8 ×10-7 Permittivity (dielectric constant) 80.20 87.90 ~90 ~98 Source: Mainly Ref. 69. solids. Of greater interest is the fact that the thermal conductivity of ice at 0°C is approximately four times that of water at the same temperature, indicating that ice will conduct heat energy at a much greater rate than will immobilized water (e.g., in tissue). The thermal diffusivities of water and ice are of even greater interest since these values indicate the rate at which the solid and liquid forms of HOH will undergo changes in temperature. Ice has a thermal diffusivity approximately nine times greater than that of water, indicating that ice, in a given environment, will undergo a temperature change at a much greater rate than will water. These sizable differences in thermal conductivity and thermal diffusivity values of water and ice provide a sound basis for explaining why tissues freeze more rapidly than they thaw, when equal but reversed temperature differentials are employed. 2.3 The Water Molecule Water’s unusual properties suggest the existence of strong attractive forces among water molecules, and uncommon structures for water and ice. These features are best explained by considering the nature of first a single water molecule and then small groups of molecules. To form a molecule of water, two hydrogen atoms approach the two sp3 bonding orbitals of oxygen and form two covalent sigma (s) bonds (40% partial ionic character), each of which has a dissociation energy of 4.6×102 kJ/mol (110 kcal/mol). The localized molecular orbitals remain symmetrically oriented about the original orbital axes, thus retaining an approximate Pag e 21 tetrahedral structure. A schematic orbital model of a water molecule is shown in Figure 1A and the appropriate van der Waals radii are shown in Figure 1B. The bond angle of the isolated water molecule (vapor state) is 104.5° and this value is near the perfect tetrahedral angle of 109°28′. The O-H internuclear distance is 0.96 Å and the van der Waals radii for oxygen and hydrogen are, respectively, 1.40 and 1.2 Å. At this point, it is important to emphasize that the picture so far presented is oversimplified. Pure water contains not only ordinary HOH molecules but also many other constituents in trace amounts. In addition to the common isotopes 16O and 1H, also present are 17O, 18O, 2H FIGURE 1 Schematic model of a sing le HOH molecule: (a) sp 3 config uration, and (b) van der W aals radii for a HOH molecule in the vapor state. Pag e 22 (deuterium) and 3H (tritium), giving rise to 18 isotopic variants of molecular HOH. Water also contains ionic particles such as hydrogen ions (existing as H3O+ ), hydroxyl ions, and their isotopic variants. Water therefore consists of more than 33 chemical variants of HOH, but the variants occur in only minute amounts. 2.4 Association of Water Molecules The V-like form of an HOH molecule and the polarized nature of the O-H bond result in an unsymmetrical charge distribution and a vapor-state dipole moment of 1.84D for pure water. Polarity of this magnitude produces intermolecular attractive forces, and water molecules therefore associate with considerable tenacity. Water’s unusually large intermolecular attractive force cannot, however, be fully accounted for on the basis of its large dipole moment. This is not surprising, since dipole moments give no indication of the degree to which charges are exposed or of the geometry of the molecule, and these aspects, of course, have an important bearing on the intensity of molecular association. Water’s large intermolecular attractive forces can be explained quite adequately in terms of its ability to engage in multiple hydrogen bonding on a three-dimensional basis. Compared to covalent bonds (average bond energy of about 335 kJ/mol), hydrogen bonds are weak (typically 2–40 kJ/mol) and have greater and more variable lengths. The hydrogen bond between oxygen and hydrogen has a dissociation energy of about 13–25 kJ/mol. Since electrostatic forces make a major contribution to the energy of the hydrogen bond (perhaps the largest contribution), and since an electrostatic model of water is simple and leads to an essentially correct geometric picture of HOH molecules as they are known to exist in ice, further discussion of the geometrical patterns formed by association of water molecules will emphasize electrostatic effects. This simplified approach, while entirely satisfactory for present purposes, must be modified if other behavioral characteristics of water are to be explained satisfactorily. The highly electronegative oxygen of the water molecule can be visualized as partially drawing away the single electrons from the two covalently bonded hydrogen atoms, thereby leaving each hydrogen atom with a partial positive charge and a minimal electron shield; that is, each hydrogen atom assumes some characteristics of a bare proton. Since the hydrogen—oxygen bonding orbitals are located on two of the axes of an imaginary tetrahedron (Fig. 1a), these two axes can be thought of as representing lines of positive force (hydrogen-bond donor sites). Oxygen’s two lone-pair orbitals can be pictured as residing along the remaining two axes of the imaginary tetrahedron, and these then represent lines of negative force (hydrogen-bond acceptor sites). By virtue of these four lines of force, each water molecule is able to hydrogen-bond with a maximum of four others. The resulting tetrahedral arrangement is depicted in Figure 2. Because each water molecule has an equal number of hydrogen-bond donor and receptor sites, arranged to permit threedimensional hydrogen bonding, it is found that the attractive forces among water molecules are unusually large, even when compared to those existing among other small molecules that also engage in hydrogen bonding (e.g., NH3, HF). Ammonia, with its tetrahedral arrangement of three hydrogens and one receptor site, and hydrogen fluoride, with its tetrahedral arrangement of one hydrogen and three receptor sites, do not have equal numbers of donor and receptor sites and therefore can form only twodimensional hydrogen-bonded networks involving less hydrogen bonds per molecule than water. Conceptualizing the association of a few water molecules becomes considerably more complicated when one considers isotopic variants and hydronium and hydroxyl ions. The Pag e 23 hydronium ion, because of its positive charge, would be expected to exhibit a greater hydrogen-bond donating potential than nonionized water (dashed lines are hydrogen bonds). STRUCTURE 1 Structure and hydrog en-bond possibilities for a hydronium ion. Dashed lines are hydrog en bonds. The hydroxyl ion, because of its negative charge, would be expected to exhibit a greater hydrogen-bond acceptor potential than un-ionized water (XH represents a solute or a water molecule). STRUCTURE 2 Structure and hydrog en-bond possibilities for a hydroxyl ion. Dashed lines are hydrog en bonds and XH represents a solute or a water molecule. Water’s ability to engage in three-dimensional hydrogen bonding provides a logical explanation for many of its unusual properties; its large values for heat capacity, melting point, boiling point, surface tension, and enthalpies of various phase transitions all are related to the extra energy needed to break intermolecular hydrogen bonds. The permittivity (dielectric constant) of water is also influenced by hydrogen bonding. Although water is a dipole, this fact alone does not account for the magnitude of its permittivity. Hydrogen-bonded clusters of molecules apparently give rise to multimolecular di- Pag e 24 FIGURE 2 Hydrog en bonding of water molecules in a tetrahedral config uration. Open circles are oxyg en atoms and closed circles are hydrog en atoms. Hydrog en bonds are represented by dashed lines. poles, which effectively increase the permittivity of water. Water’s viscosity is discussed in a later section. 2.5 Structure of Ice The structure of ice will be considered before the structure of water because the former is far better understood than the latter, and because ice’s structure represents a logical extension of the information presented in the previous section. 2.5.1 Pure Ice Water, with its tetrahedrally directed forces, crystallizes in an open (low density) structure that has been accurately elucidated. The O-O internuclear nearest-neighbor distance in ice is 2.76 Å and the O-O-O bond angle is about 109°, or very close to the perfect tetrahedral angle of 109°28′ (Fig. 3). The manner in which each HOH molecule can associate with four others (coordination number of four) is easily visualized in the unit cell of Figure 3 by considering molecule W and its four nearest neighbors 1, 2, 3, and W’. When several unit cells are combined and viewed from the top (down the c axis) the hexagonal symmetry of ice becomes apparent. This is shown in Figure 4a. The tetrahedral FIGURE 3 Unit cell of ordinary ice at 0°C. Circles represent oxyg en atoms of water molecules. Nearest-neig hbor internuclear O-O distance is 2.76 Å; q is 109°. substructure is evident from molecule W and its four nearest neighbors, with 1, 2, and 3 being visible, and the fourth lying below the plane of the paper directly under molecule W. When Figure 4a is viewed in three dimensions, as in Figure 4b, it is evident that two planes of molecules are involved (open and filled circles). These two planes are parallel, very close together, and they move as a unit during the “slip” or flow of ice under pressure, as in a glacier. Pairs of planes of this type comprise the “basal planes” of ice. By stacking several basal planes, an extended structure of ice is obtained. Three basal planes have been combined to form the structure shown in Figure 5. Viewed down the c axis, the appearance is exactly the same as that shown in Figure 4a, indicating that the basal planes are perfectly aligned. Ice is monorefringent in this direction, whereas it is birefringent in all other directions. The c axis is therefore the optical axis of ice. With regard to the location of hydrogen atoms in ice, it is generally agreed that: 1. Each line connecting two nearest neighbor oxygen atoms is occupied by one hydrogen atom located 1 ± 0.01 Å from the oxygen to which it is covalently bonded, and 1.76 ± 0.01 A° from the oxygen to which it is hydrogen bonded. This configuration is shown in Figure 6A. 2. If the locations of hydrogen atoms are viewed over a period of time, a somewhat different picture is obtained. A hydrogen atom on a line connecting two nearest neighbor oxygen atoms, X and Y, can situate itself in one of two possible positions—either 1 Å from X or 1 A° from Y. The two positions have an equal probability of being occupied. Expressed in another way, each position will, on the average, be occupied half of the time. This is possible because HOH molecules, except at extremely low temperatures, can cooperatively rotate, and hydrogen atoms can “jump” be- Pag e 26 FIGURE 4 The “basal plane” of ice (combination of two layers of slig htly different elevation). Each circle represents the oxyg en atom of a water molecule. Open and shaded circles, respectively, represent oxyg en atoms in the upper and lower layers of the basal planes. (a) Hexag onal structure viewed down the c axis. Numbered molecules relate to the unit cell in Fig ure 3. (b) Three-dimensional view of the basal plane. The front edg e of view b corresponds to the bottom edg e of view a. The crystallog raphic axes have been positioned in accordance with external (point) symmetry. tween adjacent oxygen atoms. The resulting mean structure, known also as the half-hydrogen, Pauling, or statistical structure, is shown in Figure 6B. With respect to crystal symmetry, ordinary ice belongs to the dihexagonal bipyramidal class of the hexagonal system. In addition, ice can exist in nine other crystalline polymorphic structures, and also in an amorphous or vitreous state of rather uncertain but largely noncrystal- Pag e 27 FIGURE 5 The extended structure of ordinary ice. Only oxyg en atoms are shown. Open and shaded circles, respectively, represent oxyg en atoms in upper and lower layers of a basal plane. line structure. Of the eleven total structures, only ordinary hexagonal ice is stable under normal pressure at 0°C. The structure of ice is not as simple as has been indicated. First of all, pure ice contains not only ordinary HOH molecules but also ionic and isotopic variants of HOH. Fortunately, the isotopic variants occur in such small amounts that they can, in most instances, be ignored, leaving for major consideration only HOH, H+ (H3O+ ), and OH- . Second, ice crystals are never perfect, and the defects encountered are usually of the orientational (caused by proton dislocation accompanied by neutralizing orientations) or ionic types (caused by proton dislocation with formation of H3O+ and OH- ) (see Fig. 7). The presence of these defects provides a means for explaining the mobility of protons in ice and the small decrease in dc electrical conductivity that occurs when water is frozen. In addition to the atomic mobilities involved in crystal defects, there are other types of activity in ice. Each HOH molecule in ice is believed to vibrate with a root mean amplitude (assuming each molecule vibrates as a unit) of about 0.4 Å at -10°C. Furthermore, HOH molecules that presumably exist in some of the interstitial spaces in ice can apparently diffuse slowly through the lattice. Ice therefore is far from static or homogeneous, and its characteristics are dependent on temperature. Although the HOH molecules in ice are four-coordinated at all temperatures, it is necessary to lower the temperature to about -180°C or lower to “fix” the hydrogen atoms in one of the many possible configurations. Therefore, only at temperatures near -180°C or lower will Pag e 28 FIGURE 6 Location of hydrog en atoms ( ) in the structure of ice. (A) Instantaneous structure. (B) Mean structure [known also as the half-hydrog en ( ), Pauling , or statistical structure]. Open circle is oxyg en. FIGURE 7 Schematic representation of proton defects in ice. (A) Formation of orientational defects. (B) Formation of ionic defects. Open and shaded circles represent oxyg en and hydrog en atoms, respectively. Solid and dashed lines represent chemical bonds and hydrog en bonds, respectively. Pag e 29 all hydrogen bonds be intact, and as the temperature is raised, the mean number of intact (fixed) hydrogen bonds will decrease gradually. 2.5.2 Ice in the Presence of Solutes The amount and kind of solutes present can influence the quantity, size, structure, location, and orientation of ice crystals. Consideration here will be given only to the effects of solutes on ice structure. Luyet and co-workers [75,77] studied the nature of ice crystals formed in the presence of various solutes including sucrose, glycerol, gelatin, albumin, and myosin. They devised a classification system based on morphology, elements of symmetry, and the cooling velocity required for development of various types of ice structure. Their four major classes are hexagonal forms, irregular dendrites; coarse spherulites, and evanescent spherulites. The hexogonal form, which is most highly ordered, is found exclusively in foods, provided extremely rapid freezing is avoided and the solute is of a type and concentration that does not interfere unduly with the mobility of water molecules. Gelatin at high concentrations will, for example, result in more disordered forms of ice crystals. 2.6 Structure of Water To some, it may seem strange to speak of structure in a liquid when fluidity is the essence of the liquid state. Yet it is an old and well-accepted idea [96] that liquid water has structure, obviously not sufficiently established to produce long-range rigidity, but certainly far more organized than that of molecules in the vapor state, and ample to cause the orientation and mobility of a given water molecule to be influenced by its neighbors. Evidence for this view is compelling. For example, water is an “open” liquid, being only 60% as dense as would be expected on the basis of close packing that can prevail in nonstructured liquids. Partial retention of the open, hydrogen-bonded, tetrahedral arrangement of ice easily accounts for water’s low density. Furthermore, the heat of fusion of ice, while unusually high, is sufficient to break only about 15% of the hydrogen bonds believed to exist in ice. Although this does not necessarily require that 85% of the hydrogen bonds existing in ice be retained in water (for example, more could be broken, but the change in energy could be masked by a simultaneous increase in van der Waals interactions), results of many studies support the notion that many water-water hydrogen bonds do exist. Elucidation of the structure of pure water is an extremely complex problem. Many theories have been set forth, but all are incomplete, overly simple, and subject to weaknesses that are quickly cited by supporters of rival theories. That is, of course, a healthy situation, which will eventually result in an accurate structural picture (or pictures) of water. In the meantime, few statements can be made with any assurance that they will stand essentially unmodified in years to come. Thus, this subject will be dealt with only briefly. Three general types of models have been proposed: mixture, interstitial, and continuum (also referred to as homogeneous or uniformist models) [5]. Mixture models embody the concept of intermolecular hydrogen bonds being momentarily concentrated in bulky clusters of water molecules that exist in dynamic equilibrium with other more dense species—momentarily meaning ~10- 11 sec [73]. Continuum models involve the idea that intermolecular hydrogen bonds are distributed uniformly throughout the sample, and that many of the bonds existing in ice simply become distorted rather than broken when ice is melted. It has been suggested that this permits a continuous network of water molecules to exist that is, of course, dynamic in nature [107,120]. The interstitial model involves the concept of water retaining an ice-like or clathrate-type Pag e 30 structure with individual water molecules filling the interstitial spaces of the clathrates. In all three models, the dominant structural feature is the hydrogen-bonded association of liquid water in ephemeral, distorted tetrahedra. All models also permit individual water molecules to frequently alter their bonding arrangements by rapidly terminating one hydrogen bond in exchange for a new one, while maintaining, at constant temperature, a constant degree of hydrogen bonding and structure for the entire system. The degree of intermolecular hydrogen bonding among water molecules is, of course, temperature dependent. Ice at 0°C has a coordination number (number of nearest neighbors) of 4.0, with nearest neighbors at a distance of 2.76 Å. With input of the latent heat of fusion, melting occurs; that is, some hydrogen bonds are broken (distance between nearest neighbors increases) and others are strained as water molecules assume a fluid state with associations that are, on average, more compact. As the temperature is raised, the coordination number increases from 4.0 in ice at 0°C, to 4.4 in water at 1.50°C, then to 4.9 at 83°C. Simultaneously, the distance between nearest neighbors increases from 2.76 Å in ice at 0°C, to 2.9 Å in water at 1.5°C, then to 3.05 Å at 83°C [7,80]. It is evident, therefore, that the ice-to-water transformation is accompanied by an increase in the distance between nearest neighbors (decreased density) and by an increase in the average number of nearest neighbors (increased density), with the latter factor predominating to yield the familiar net increase in density. Further warming above the melting point causes the density to pass through a maximum at 3.98°C, then gradually decline. It is apparent, then, that the effect of an increase in coordination number predominates at temperatures between 0 and 3.98°C, and that the effect of increasing distance between nearest neighbors (thermal expansion) predominates above 3.98°C. The low viscosity of water is readily reconcilable with the type of structures that have been described, since the hydrogenbonded arrangements of water molecules are highly dynamic, allowing individual molecules, within the time frame of nano- to picoseconds, to alter their hydrogen-bonding relationships with neighboring molecules, thereby facilitating mobility and fluidity. 2.7 Water-Solute Interactions 2.7.1 Macroscopic Level (Water Binding, Hydration, and Water Holding Capacity) Before dealing with water-solute interactions at the molecular level, it is appropriate to discuss water-related phenomena referred to by terms such as water binding, hydration, and water holding capacity. With respect to foods, the terms “water binding” and “hydration” are often used to convey a general tendency for water to associate with hydrophilic substances, including cellular materials. When used in this manner, the terms pertain to the macroscopic level. Although more specialized terms, such as “water binding potential,” are defined in quantitative terms, they still apply only to the macroscopic level. The degree and tenacity of water binding or hydration depends on a number of factors including the nature of the nonaqueous constituent, salt composition, pH, and temperature. “Water holding capacity” is a term that is frequently employed to describe the ability of a matrix of molecules, usually macromolecules present at low concentrations, to physically entrap large amounts of water in a manner that inhibits exudation. Familiar food matrices that entrap water in this way include gels of pectin and starch, and cells of tissues, both plant and animal. Pag e 31 Physically entrapped water does not flow from tissue foods even when they are cut or minced. On the other hand, this water behaves almost like pure water during food processing; that is, it is easily removed during drying, is easily converted to ice during freezing, and is available as a solvent. Thus, its bulk flow is severely restricted, but movement of individual molecules is essentially the same as that of water molecules in a dilute salt solution. Nearly all of the water in tissues and gels can be categorized as physically entrapped, and impairment of the entrapment capability (water holding capacity) of foods has a profound effect on food quality. Examples of quality defects arising from impairment of water holding capacity are syneresis of gels, thaw exudate from previously frozen foods, and inferior performance of animal tissue in sausage resulting from a decline in muscle pH during normal physiological events postmortem. Gel structures and water holding capacity are discussed more fully in other chapters. 2.7.2 Molecular Level: General Comments Mixing of solutes and water results in altered properties of both constituents. Hydrophilic solutes cause changes in the structure and mobility of adjacent water, and water causes changes in the reactivity, and sometimes structure, of hydrophilic solutes. Hydrophobic groups of added solutes interact only weakly with adjacent water, preferring a nonaqueous environment. The bonding forces existing between water and various kinds of solutes are of obvious interest, and these are summarized in Table 3. 2.7.3 Molecular Level: Bound Water Bound water is not a homogeneous, easily identifiable entity, and because of this, descriptive terminology is difficult, numerous definitions have been suggested, and there is no consensus about which one is best. This term is controversial, frequently misused, and in general, poorly understood, causing increasing numbers of scientists to suggest that its use be terminated. TABLE 3 Classifications of Types of W ater-Solute Interactions Type Example Streng th of interaction compared to water-water hydrog en bond a Dipole-ion W ater-free ion W ater-charg ed g roup on org anic molecule Greaterb Dipole-dipole W ater-protein NH W ater-protein CO W ater-sidechain OH Approx. equal Hydrophobic hydration W ater + Rc R(hydrated) Much less (DG>0) Hydrophobic interaction R(hydrated) + R(hydrated) R2 (hydrated) + H2O Not comparable d (>hydrophobic interaction; DG<0) aAbout 12-25 kJ/mol. bBut much weaker than streng th of sing le covalent bond. cR is alkyl g roup. dHydrophobic interactions are entropy driven, whereas dipole-ion and dipole-dipole interactions are enthalpy driven. Pag e 32 Although this may be desirable, the term “bound water” is so common in the literature that it must be discussed. The numerous definitions proposed for “bound water” should indicate why this term has created confusion [3,51]: 1. Bound water is the equilibrium water content of a sample at some appropriate temperature and low relative humidity. 2. Bound water is that which does not contribute significantly to permittivity at high frequencies and therefore has its rotational mobility restricted by the substance with which it is associated. 3. Bound water is that which does not freeze at some arbitrary low temperature (usually-40°C or lower). 4. Bound water is that which is unavailable as a solvent for additional solutes. 5. Bound water is that which produces line broadening in experiments involving proton nuclear magnetic resonance. 6. Bound water is that which moves with a macromolecule in experiments involving sedimentation rates, viscosity, or diffusion. 7. Bound water is that which exists in the vicinity of solutes and other nonaqueous substances and has properties differing significantly from those of “bulk” water in the same system. All of these definitions are valid, but few will produce the same value when a given sample is analyzed. From a conceptual standpoint it is useful to think of bound water as “water that exists in the vicinity of solutes and other nonaqueous constituents, and exhibits properties that are significantly altered from those of ‘bulk water’ in the same system.” Bound water should be thought of as having “hindered mobility” as compared to bulk water, not as being “immobilized.” In a typical food of high water content, this type of water comprises only a minute part of the total water present, approximately the first layer of water molecules adjacent to hydrophilic groups. The subject of bound water (hindered water) will be discussed further in the section dealing with molecular mobility (Mm) in frozen systems. Interactions between water and specific classes of solutes will now be considered. 2.7.4 Interaction of Water with Ions and Ionic Groups Ions and ionic groups of organic molecules hinder mobility of water molecules to a greater degree than do any other types of solutes. The strength of water-ion bonds is greater than that of water-water hydrogen bonds, but is much less than that of covalent bonds. The normal structure of pure water (based on a hydrogen-bonded, tetrahedral arrangement) is disrupted by the addition of dissociable solutes. Water and simple inorganic ions undergo dipole-ion interactions. The example in Figure 8 involves hydration of the NaCl ion pair. Only first-layer water molecules in the plane of the paper are illustrated. In a dilute solution of ions in water, second-layer water is believed to exist in a structurally perturbed state because of conflicting structural influences of first-layer water and the more distant, tetrahedrally oriented “bulk-phase” water. In concentrated salt solutions, bulk-phase water would not exist and water structure would be dominated by the ions. There is abundant evidence indicating that some ions in dilute aqueous solution have a net structure-breaking effect (solution is more fluid than pure water), whereas others have a net structure-forming effect (solution is less fluid than pure water). It should be understood that the term “net structure” refers to all kinds of structures, either normal or new types of water structure. From the standpoint of “normal” water structure, all ions are disruptive The ability of a given ion to alter net structure is related closely to its polarizing power (charge divided by radius) or simply the strength of its electric field. Ions that are small and/or multivalent (mostly positive ions, such as Li+ , Na+ , H3O+ , Ca2+, Ba2+ , Mg2+, Al3+, F- , and OH- ) have strong electric fields and are net structure formers. The structure imposed by these ions more than compensates for any loss in normal water structure. These ions strongly interact with the four to six first-layer water molecules, causing them to be less mobile and pack more densely than HOH molecules in pure water. Ions that are large and monovalent (most of the negatively charged ions and large positive ions, such as K+ , Rb+ , Cs+ , Cl- , Br- , I- , , , , and have rather weak electric fields and are net structure breakers, although the effect is very slight with K+ . These ions disrupt the normal structure of water and fail to impose a compensating amount of new structure. Ions, of course, have effects that extend well beyond their influence on water structure. Through their varying abilities to hydrate (compete for water), alter water structure, influence the permittivity of the aqueous medium, and govern the thickness of the electric double layer around colloids, ions profoundly influence the “degree of hospitality” extended to other nonaqueous solutes and to substances suspended in the medium. Thus, conformation of proteins and stability of colloids (salting-in, salting-out in accord with the Hofmeister or lyotropic series) are greatly influenced by the kinds and amounts of ions present [18,68]. 2.7.5 Interaction of Water with Neutral Groups Possessing Hydrogen-Bonding Capabilities (Hydrophilic Solutes) Interactions between water and nonionic, hydrophilic solutes are weaker than water-ion interactions and about the same strength as those of water-water hydrogen bonds. Depending on the strength of the water-solute hydrogen bonds, first-layer water may or may not exhibit reduced mobility and other altered properties as compared to bulk-phase water. Solutes capable of hydrogen bonding might be expected to enhance or at least not disrupt the normal structure of pure water. However, in some instances it is found that the distribution and orientation of the solute's hydrogen-bonding sites are geometrically incompatible with those existing in normal water. Thus, these kinds of solutes frequently have a disruptive influence on the normal structure of water. Urea is a good example of a small hydrogen-bonding solute that for geometric reasons has a marked disruptive effect on the normal structure of water. Pag e 34 It should be noted that the total number of hydrogen bonds per mole of solution may not be significantly altered by addition of a hydrogen-bonding solute that disrupts the normal structure of water. This is possible since disrupted water-water hydrogen bonds may be replaced by water-solute hydrogen bonds. Solutes that behave in this manner have little influence on “net structure” as defined in the previous section. Hydrogen bonding of water can occur with various potentially eligible groups (e.g., hydroxy1, amino, carbony1, amide, imino, etc.). This sometimes results in “water bridges” where one water molecule interacts with two eligible hydrogen-bonding sites on one or more solutes. A schematic depiction of water hydrogen bonding (dashed lines) to two kinds of functional groups found in proteins is shown: STRUCTURE 3 Hydrog en bonding (dotted lines) of water to two kinds of functional g roups occurring in proteins. A more elaborate example involving a three-HOH bridge between backbone peptide units in papain is shown in Figure 9. It has been observed that hydrophilic groups in many crystalline macromolecules are separated by distances identical to the nearest-neighbor oxygen spacing in pure water. If this spacing prevails in hydrated macromolecules this would encourage cooperative hydrogen bonding in first- and second-layer water. FIGURE 9 Examples of a three-molecule water bridg e in papain; 23, 24, and 25 are water molecules. (From Ref. 4.) Pag e 35 2.7.6 Interaction of Water with Nonpolar Substances The mixing of water and hydrophobic substances, such as hydrocarbons, rare gases, and the apolar groups of fatty acids, amino acids, and proteins is, not surprisingly, a thermodynamically unfavorable event (DG>0). The free energy is positive not because DH is positive, which is typically true for low-solubility solutes, but because TDS is negative [30]. This decrease in entropy occurs because of special structures that water forms in the vicinity of these incompatible apolar entities. This process has been referred to as hydrophobic hydration (Table 3 and Fig. 10a). Because hydrophobic hydration is thermodynamically unfavorable, it is understandable that water would tend to minimize its association with apolar entities that are present. Thus, if two separated apolar groups are present, the incompatible aqueous environment will encourage them to associate, thereby lessening the water-apolar interfacial area—a process that is thermodynamically favorable (DG<0). This process, which is a partial reversal of hydrophobic hydration, is referred to as “hydrophobic interaction” and in its simplest form can be depicted as where R is an apolar group (Table 3 and Fig. 10b). Because water and apolar groups exist in an antagonistic relationship, water structures itself to minimize contact with apolar groups. The type of water structure believed to exist in the layer next to apolar groups is depicted in Figure 11. Two aspects of the antagonistic relationship between water and hydrophobic groups are worthy of elaboration: formation of clathrate hydrates, and association of water with hydrophobic groups in proteins. A clathrate hydrate is an ice-like inclusion compound wherein water, the “host” substance, forms a hydrogen-bonded, cage-like structure that physically entraps a small apolar molecule, FIGURE10 Schematic depiction of (a) hydrophobic hydration and (b) hydrophobic association. Open circles are hydrophobic g roups. Hatched areas are water. (Adapted from Ref. 28.) Pag e 36 FIGURE11 Proposed water orientation at a hydrophobic surface. (Adapted from Ref. 68.) known as the “guest.” These entities are of interest because they represent the most extreme structure-forming response of water to an apolar substance and because microstructures of a similar type may occur naturally in biological matter. Clathrate hydrates are, in fact, crystalline, they can easily be grown to visible size, and some are stable at temperatures above 0°C provided the pressure is sufficient. The guest molecules of clathrate hydrates are low-molecular-weight compounds with sizes and shapes compatible with the dimensions of host water cages comprised of 20–74 water molecules. Typical guests include low-molecular-weight hydrocarbons and halogenated hydrocarbons; rare gases; short-chain primary, secondary, and tertiary amines; and alkyl ammonium, sulfonium, and phosphonium salts. Interaction between water and guest is slight, usually involving nothing more than weak van der Waals forces. Clathrate hydrates are the extraordinary result of water's attempt to avoid contact with hydrophobic groups. There is evidence that structures similar to crystalline clathrate hydrates may exist naturally in biological matter, and if so, these structures would be of far greater importance than crystalline hydrates since they would likely influence the conformation, reactivity, and stability of molecules such as proteins. For example, it has been suggested that partial clathrate structures may exist around the exposed hydrophobic groups of proteins. It is also possible that clathrate-like structures of water have a role in the anesthetic action of inert gases such as xenon. For further information on clathrates, the reader is referred to Davidson [15]. Unavoidable association of water with hydrophobic groups of proteins has an important influence on protein functionality [5,124]. The extent of these unavoidable contacts is potentially fairly great because nonpolar side chains exist on about 40% of the amino acids in typical oligomeric food proteins. These nonpolar groups include the methy1 group of alanine, the benzyl group of phenylalanine, the isopropy1 group of the valine, the mercaptomethy1 group of cysteine, and the secondary buty1 and isobuty1 groups of the leucines. The nonpolar groups of other compounds such as alcohols, fatty acids, and free amino acids also can participate in hydrophobic interactions, but the consequences of these interactions are undoubtedly less important than those involving proteins. Because exposure of protein nonpolar groups to water is thermodynamically unfavorable, association of hydrophobic groups or “hydrophobic interaction” is encouraged, and this occurrence is depicted schematically in Figure 12. Hydrophobic interaction provides a major driving

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