BIOCHEMISTRY - Textbook - Ostapchenko L. I. - 2012

Chapter 2. ROLE OF WATER, MACRO- AND MICROELEMENTS IN THE VITAL ACTIVITY OF ORGANISMS

2.1. The Unique Property of Water: Dissolving Substances of Diverse Origins

The unique properties of Water are determined by its characteristic features. Compared to other liquids, water has high boiling and melting points, as well as a high heat of vaporization (the energy required to convert 1 g of liquid into vapor at boiling point and atmospheric pressure). For example, these values for water are +100 °С, 0 °С, and 540 cal, respectively, while for ethanol they are +78 °С, -117 °С, and 204 cal. These data indicate strong attraction forces between water molecules. Despite the water molecule being overall electroneutral (having an equal number of electrons and protons), the electrons are distributed asymmetrically, giving it a polar character. The oxygen atomic Nucleus slightly pulls electrons away from the hydrogen atomic nuclei, leaving them with a small net positive charge. Regions with a small net negative charge are located near the oxygen atom at two corners of an imaginary tetrahedron (Fig. 2.1, A).

Such charge Separation leads to the creation of a dipole moment. Due to this polarization, two adjacent water molecules form Hydrogen Bonds that are 25 times weaker (4.5 kcal/mol = 18.83 kJ/mol) than covalent bonds (110 kcal/mol = 460.25 kJ/mol), meaning that two neighboring water molecules interact electrostatically. This type of electrostatic attraction is called a Hydrogen bond, with a half-life of 10 ⁹ s. Since a water molecule is a tetrahedron (Fig. 2.1, A), it can bond a maximum of four neighboring water molecules. Through these hydrogen bonds, water molecules aggregate into short-lived (10^-10 s) clusters (Fig. 2.1, B), in which each molecule is connected to four neighbors. As the Temperature rises, the average cluster size decreases noticeably. While at 0 °С a cluster reaches up to 90 molecules, at 7 °С it drops to only 25; thus, increasing temperature causes a phenomenon akin to the "melting" of the cluster. The heat contribution of this "melting" accounts for a third of water's total heat capacity—18 kcal/deg·mol. This explains water's exceptionally high heat capacity compared to other liquids.

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Fig. 2.1. Cytology/cytology/92.html">SCHEMATIC Structure OF a water molecule explaining its cohesive properties:

A - localization of "charges" within an imaginary tetrahedron: 1 - electropositive region, 2 - electronegative region;

B - flickering cluster of four water molecules; C - interaction of water molecules with cations (K⁺) and anions (A^-);

D - behavior of hydrophobic molecules in water ensuring minimal contact with water

The described nature of water is called cohesive (from Lat. *cohaesus* - connected). It determines many of water's unusual properties, including high surface tension, specific heat capacity, and heat of vaporization.

Because water molecules are polar, they cluster around ions or other polar molecules (Fig. 2.1, C). Such ions and molecules that participate in forming water structures stabilized by Hydrogen bonds are called hydrophilic. They dissolve well in water; in other words, ionizable substances dissolve because bipolar water molecules interact with ions, hydrate them, and bring them into solution. If substances are polar but do not ionize (sugars, alcohols, aldehydes), they dissolve in water through The formation of H-bonds with their OH groups. Conversely, nonpolar molecules disrupt the water structure formed by H-bonds. Such compounds are called hydrophobic and are insoluble in water. Therefore, two or more such groups in water tend to approach each other, clustering as if being squeezed out by water. As a result of this hydrophobic behavior, the water structure is less disrupted (Fig. 2.1, D). However, many substances are largely hydrophobic with a hydrophilic portion. These are amphipathic compounds, such as fatty acid salts. Such substances cannot dissolve in water; instead, they disperse, forming aggregates called micelles. In such structures, the hydrophilic negatively charged carboxyl groups (COO^-) face the water, interacting with water molecules as dipoles. The nonpolar PARTS OF THE molecule are "hidden" inside the micelle (following the principles shown in Fig. 2.1, G). Micelles do not aggregate because they carry a negative charge (Fig. 2.2). Consequently, soapy water (soap, e.g., sodium oleate) is always cloudy because micelles scatter light.

Given the low energy of hydrogen bonds, their half-life, and a bond length of 0.26–0.31 nm in the liquid phase, each water molecule at elevated temperatures forms H-bonds with a smaller number of similar molecules. At room temperature, each water molecule forms bonds with 3–4 other molecules, whereas in ice (crystal lattice) it bonds with the maximum possible number of water molecules, i.e., four. Ice floats On the surface of water precisely because of its lower density.

Fig. 2.2. Schematic structure of fatty acid micelles:

"-" - Negative charge of the ionized carboxyl group

A hydrogen bond is strongest when three atoms lie in the same plane (O---H-O). If this directionality changes, the energy of the H-bond decreases, which inevitably affects the stabilization (or alteration) of the Spatial Structure of macromolecules such as Proteins, Nucleic Acids, and Polysaccharides. This is particularly important during the formation of complementary bonds.

The cooperative nature of H-bond formation between neighboring molecules plays an exceptional role in determining The properties of water and ice. In ice, each water molecule is bonded to all four of its neighbors, forming the maximum possible number of H-bonds. A continuous network of H-bonds unites all water molecules into a single system—"one giant molecule" with an exceptionally openwork structure.

Protons in such a network are not located exactly in the middle between oxygen atoms, but rather lie closer to the oxygen atom with which they share a covalent bond. Since The structure of a water molecule is symmetrical, protons have two possible states rather than one. Protons transition from one state to another via the "tunnel" of the hydrogen bond (indicated by the dashed line in Fig. 2.1, B), which facilitates this transition 70-fold. This ability of the proton to exist in two potential states within the ice structure explains its anomalously high mobility, which approaches that of electrons in metals. It is important to understand that it is not the same physical proton that moves, but merely the state of its free form; that is, protons attach to the nearest water molecule, from which another proton dissociates and attaches to the next, and so on. Such proton migration through the H-bond tunnels lowers energy barriers, meaning that ice exhibits long-range action capabilities.

True molecular solutions possess interrelated properties: boiling and freezing points, vapor pressure, and osmotic pressure. These four properties change under METABOLISM/18.html">The Influence of dissolved substances. However, this depends not on their chemical nature and size, but on the number of dissolved particles per unit volume. This is explained by the fact that 1 mole of any compound contains 6.02 × 10²³ molecules (Avogadro's number). Such properties of water are of great biological significance. For example, freshwater fish remain active at the freezing point of water because their Blood (and other fluids) has a sufficient concentration of solutes to lower its freezing point below that of pure water.

Another reason for the colligative properties of water is that dissolved substances attempt to disrupt hydrogen bonds between water molecules and "divert" a portion of water molecules to form Hydration shells (Fig. 2.1, C). This degrades the cluster structure of Water and Its properties as a solvent. Therefore, solutions of neutral salts (such as NaCl) are used to separate protein mixtures ("protein salting-out"), as different proteins precipitate from salt solutions at different rates.

Special attention is paid to H⁺ and OH^- ions, which rank first among monovalent ions in their Structuring effect on water. In water, the H⁺ ion does not exist in a free state; instead, it complexes with a water molecule to form the hydronium (hydroxonium) ion H₃O⁺. (In reality, each H⁺ is surrounded by several H2O molecules, the number of which depends on temperature). H₃O⁺ and OH^- ions coordinate three water molecules around themselves, creating the structures H₃O⁺ (Н₂О)₃ and OH^- (Н₂О)₃. At the same time, H⁺ and OH^- ions are not spatially localized but move at high speed, meaning their structuring effect is averaged across all water molecules. This is how the long-range effect in structure ordering manifests. Long-range action is also evident, albeit to a lesser extent, in the Influence of other ions on water. It consists in ions altering the sizes and lifetimes of ordered clusters.

Water molecules exhibit weakly expressed electrolytic dissociation: H₂O = Н⁺ + OH^- with the Equilibrium Constant

Since the concentration of H₂O is very high (997.07 : 18.0153 = 55.35 mol/L; 997.07 being the mass of 1 L of water, and 18.0153 being the molar mass of water) and remains constant down to very low ion concentrations (10^-7 mol/L at 25 °C), then , and 55.35 K_w = [H⁺][OH^-].

Since K_w = 1.821 · 10 ¹⁶ (determined by the electrical conductivity of pure water at 25 ⁰С), this equation for the ion product of water can be written as: K_w = [H⁺ ][OH^- ] = const = 55.35 · 1.821 · 10 ¹⁶ = 100 · 10 ¹⁶ = 1 · 10 ¹⁴ . The value K_w is called the ion product of water, and at a temperature of 25 ⁰С: K_w = 1 · 10 ¹⁴ . Such a solution is neutral because the concentration of each ion, H⁺ and OH^-, is 10 ⁷ mol/L. Since K_w is always equal to 10 ¹⁴ , in acid solutions the concentration [H⁺ ] is very high and [OH^- ] is very low, meaning the solution has an acidic reaction. In alkaline solutions, conversely, [H⁺ ] < [OH^- ], yielding an alkaline reaction. Therefore, the acidity of a solution is determined by its H⁺ concentration. In scientific literature, the reaction of an aqueous medium is characterized not by the molar concentration of hydrogen ions, but by the negative decimal logarithm of this value: pH = lg 1/[H⁺ ] = - lg [H⁺] (where the symbol "p" stands for "negative logarithm"). This allows The Use of small dimensionless numbers from 0 to 14. The pH scale is logarithmic rather than arithmetic. Consequently, if the pH of two solutions differs by one pH unit, it means that their H⁺ concentrations differ by a factor of 10.

In a neutral solution, where the concentration of H+ ions is 1.0 · 10 ⁷ mol/L, the pH value at 25 ⁰С is as follows:

pH = lg 1/1.0 · 10^-7 = -lg (1 · 10^-7) = lg (1 · 10⁷) = lg 1.0 + lg10⁷ = 0 + 7 = 7.

In a similar manner, pOH can be calculated (sometimes used for the quantitative characterization of basicity, i.e., the concentration of OH ions in a solution). In all cases, pH + pOH = 14. However, in biological systems, measurements are performed at human body temperature (36.6 ⁰С), in which case pH + pOH = lg 2.325 · 10 ¹⁴ ≈ 13.6. Solutions with pH < ~ 6.8 are considered acidic, while those with pH > ~ 6.8 are alkaline. This is because the ion product of water increases with rising temperature (at 25 ⁰С, pH + pOH = 14, whereas at 100 ⁰С, pH + pOH = 12.26). The hydrogen ion exponent, pH, is widely used in biochemical research, as well as in clinical and pharmacological practice, to characterize the acid-base properties of various biological fluids and therapeutic agents.

As can be seen from Fig. 2.3, the pH of biological fluids varies over a wide range. The lowest pH value (highest H⁺ ion concentration) is characteristic of gastric juice, while the highest is found in pancreatic juice.

Fig. 2.3. H⁺ ion concentrations, their corresponding pH values, and Examples of pH values for specific fluids.

The pH values for Body Fluids span a wide range—from 1.65 for gastric juice to 8.80 for pancreatic juice. The concentration of H⁺ ions in gastric juice is approximately 10 million times greater than in pancreatic juice.

Buffer solutions are those whose pH remains practically unchanged upon The addition of small amounts of strong acids (H⁺) or bases (OH^-), as well as upon dilution. The simplest buffer solutions are:

1. A mixture of a weak acid and its salt with a common anion (acetic acid CH₃COOH and sodium acetate CH₃COONa).

2. A mixture of a weak base and its salt with a common cation (ammonium hydroxide NH₄OH and ammonium chloride NH₄Cl).

Using the "acetic acid - sodium acetate" pair as an example, let us examine The basis of a buffer system's properties—that is, its ability to counteract pH changes when a base is added, and to stabilize the hydrogen ion concentration and solution pH at a certain level. Since in the case of a pure acid [H⁺ ] = [CH₃COO^- ], and the equilibrium concentration [CH₃COOH] is practically equal to the total concentration of acetic acid (because it is a weak, poorly dissociated acid), for a 0.1 mol/L solution of it, the hydrogen ion concentration is

where 1,75·10^-5 is the electrolytic dissociation constant of this acid . The pH of such a solution will be: pH = -lg(1,32·10^-3) ≈2,88.

If sodium acetate is added to this solution to a concentration of 0,1 M, the hydrogen ion concentration (now in the buffer solution rather than the acetic acid solution) will decrease, as can be seen from the equation

and pH = - lg(1,75 · 10^{- 5} ) ≈ 4,76 . This pH value corresponds to the pK' value, as seen in Table 2.1 and from the Henderson-Hasselbalch equation analyzed below. The symbol "p" in pK', as in the case of pH, denotes the "negative logarithm".

Now let us consider the titration of an acetic acid solution with a NaOH solution. The OH^- ions formed during the dissociation of the alkali

combine with H⁺ ions to form H₂O. Since K_w = [H⁺][OH^-] = 1 ·10^-14 at all times, a decrease in [H⁺] due to H₂O formation causes some of the undissociated CH₃COOH molecules to dissociate, thereby restoring the concentration of H⁺ ions. In other words, the more NaOH is added, the more acetic acid molecules dissociate. This means that during the titration of an acetic acid solution with an alkali, the number of undissociated CH₃COOH molecules decreases, while the number of CH₃COO^- ions increases. When equality [CH₃COOH] = [CH₃COO^-] is reached (the midpoint in Fig. 2.4), the pH of the solution equals the pK' of acetic acid. In this case, pH = pK' = 4,76.

Table 2.1

Examples of electrolytic dissociation constants (K', M) of acids and their pK' values at 25 °C

Upon further addition of alkali, the remaining acetic acid molecules dissociate, H⁺ ions are consumed in the formation of water, and the concentration of acetate ions increases. The process continues until the acetic acid is completely dissociated (Fig. 2.4). This titration process can also be performed in reverse: added H⁺ ions (resulting from the dissociation of the added acid) bind with CH₃COO^-, leading to the formation of undissociated CH₃COOH molecules. Accounting for volume changes, the reverse titration curve completely coincides with The titration curve.

Figure 2.4 shows three titration curves for three weak acids of different origins. As can be seen, they have the same shape but are positioned at different pH levels due to having different dissociation constants (Table 2.1), i.e., different strengths. The greater the strength of an acid, the more readily it donates a proton. These titration curves also feature a relatively linear region, indicating small changes in pH despite the constant addition of alkali. This region is called the buffer region (zone) of the conjugate acid-base pair. At the midpoint of the buffer zone, where the concentration of the proton donor equals the concentration of the proton acceptor, the buffer capacity of the system is maximal, and pH = pK'. The figure also shows that the three buffer systems are effective near pH values close to their pK' values. Therefore, using the data presented in Figs. 2.3 and 2.4, one can select conjugate acid-Base Pairs as effective buffer systems: acetic acid–acetate for urine, and the H₂PO₄^-–HPO₄^2- pair for bile and blood. The NH₄⁺–NH₃ acid-base pair is rarely used for biological fluids. Hence, biochemical studies most frequently employ buffer systems containing Na₂CO₃, NaHCO₃, NaH₂PO₄, Na₂HPO₄, or their potassium counterparts.

Fig. 2.4. Titration curves for three weak acids:

on the left — predominant forms of these compounds at given pH values; on the right — the buffering zones of these buffer systems

In All living organisms, extracellular and intracellular fluids maintain a constant pH value regulated by buffer systems. In warm-blooded animals, the most vital buffer systems are the oxyhemoglobin-Hemoglobin system, the phosphate system, and the bicarbonate system. The latter serves as the primary buffer system in blood plasma. It consists of a conjugate acid-base pair comprising the proton donor H₂СО₃ and the proton acceptor bicarbonate ion HCO₃^-: Н₂СО₃ ↔ H⁺ + HCO₃^-. A unique feature of this system is that one of its components, H₂СО₃, is formed through the interaction of dissolved carbon dioxide CO₂ (aq) with water via the reversible reaction: СО₂ (aq) + Н₂О ↔ Н₂СО₃. The concentration of dissolved СО₂ is determined by equilibrium with the gaseous phase of carbon dioxide: СО₂ (g) ↔ СО₂ (aq).

The bicarbonate buffer system operates at a pH of approximately 7.4 because the proton donor H₂СО₃ in blood plasma remains in equilibrium with a large reserve volume of gaseous CO₂ in the Lungs. If the blood is "forced" to absorb OH^- ions, causing the pH to rise, The amount of H₂СО₃ (which has partially converted into HCO₃^- due to the reaction with OH^-) is rapidly restored by the large reservoir of gaseous СО₂ in the lungs. Gaseous СО₂ dissolves in the blood to form dissolved СО₂ (aq), which then reacts with water to yield H₂СО₃. Conversely, when the blood pH drops, some of the HCO₃^- ions bind with excess H⁺ ions, producing an excess of H₂СО₃. Carbonic acid then dissociates to release СО₂ (aq), which subsequently transitions into the gaseous phase СО₂ (g) in the lungs and is exhaled from the body.

The phosphate buffer system is characterized by the equilibrium between hydrogen phosphate and dihydrogen phosphate ions:

The oxyhemoglobin-hemoglobin buffer system accounts for approximately 75% of the blood's buffer capacity. This system is characterized by the equilibrium between hemoglobin ions Hb^- and undissociated hemoglobin HHb, which is a very weak acid: Hb^- + H⁺ ↔ HHb; Hb^- + Н₂О ↔ HHb + OH^-, as well as between oxyhemoglobin ions HbO₂^- and undissociated oxyhemoglobin HHbO₂ (which is a slightly stronger acid than hemoglobin): HbO₂^- + H⁺ ↔ HHbO₂; HbO₂^- + Н₂О ↔ HHbO₂ + OH^-

The ions HCO₃^-, HPO₄^2-, Hb^-, and HbO₂^- are anions of weak acids and function as efficient acceptors of H⁺ ions. Consequently, if strong acids enter the bloodstream, their H⁺ ions bind with these anions to form undissociated molecules of carbonic acid, dihydrogen phosphate ions, hemoglobin, and oxyhemoglobin:

This means that thanks to the buffering action of these systems, only a slight decrease in blood pH occurs. Importantly, this process reduces the content of HCO₃^-, HPO₄^2-, Hb^-, and HbO₂^- ions in the blood, thereby diminishing the buffer capacity of these systems, also known as the alkaline reserve of the blood. It is measured by the chemically bound volume of СО₂ (in the form of bicarbonates) per 100 mL of blood plasma saturated with a gas at a partial pressure of СО₂ equal to that of alveolar air (53.3 hPa). The alkaline reserve of the blood is expressed in volume fractions of chemically bound СО₂ in the blood. Under normal conditions, it corresponds to 50–70% (25–30 mmol/L).

Similar processes occur in the blood upon the Introduction of alkalis. In this case, hydroxyl ions generated during the Hydrolysis of corresponding salts interact with the free acids H₂СО₃, HHb, HHbO₂, and dihydrogen phosphate ions, releasing water:

The Blood Buffer Systems counteract shifts in pH toward lower values, given that nutrient Metabolism in the Body produces substantial amounts of СО₂ (550–775 g/day). The interaction between СО₂ and Н₂О yields carbonic acid in amounts equivalent to the introduction of 25–35 mol/day of H⁺ ions. A decrease in pH is also promoted by The conversion of hemoglobin to oxyhemoglobin in the lungs, since HHbO₂ is a stronger acid than HHb. In tissue capillaries, oxyhemoglobin is converted back into hemoglobin (the reverse of the process described above), which increases the alkaline reserve of the blood.

A shift in the acid-base balance of the blood toward an increased concentration of H⁺ ions (a drop in pH) represents a decrease in the blood's alkaline reserve. This condition is termed acidosis. For example, in diabetes, the concentration of metabolic acids increases, and blood pH drops to 6.5–6.8. An increase in pH driven by a lower concentration of H⁺ ions — which corresponds to an elevated alkaline reserve — is termed alkalosis. Acidosis and alkalosis occur either as a result of the direct intake of excess acidic or alkaline substances (food, water, beverages, medications, polluted air) or as a consequence of abnormal clearance of such substances from the body during various pathological states, such as Metabolic Disorders, respiratory dysfunction, or circulatory failure.

In clinical practice, the body's acid-base balance is determined by blood gas analysis using the Astrup method and is expressed in units of BE (base excess). The normal value (pH = 7.40) corresponds to BE = 0. BE values ranging from 0 to ±3 are likewise considered normal; BE = ±(6–9) indicates a moderate disturbance; BE = ±(10–14) is a serious threat; and BE > 14 is critical.

Titration curves of weak acids, including those described above, demonstrate a general underlying pattern. The shape of titration curves is described by the Henderson-Hasselbalch equation (one of the mathematical expressions for the acid dissociation constant), the analysis of which helps elucidate the buffering properties of blood that maintain its acid-base equilibrium. The acid dissociation constant $K'$ (often denoted as $K_a$, where "a" stands for acid) is expressed as follows:

Taking the negative Logarithms of the terms in the equation yields: -lg[H⁺] = -lg$K'$ - lg[HA]/[A^-] or pH = p$K'$ - lg[HA]/[A^-], which (upon inverting the numerator and denominator and changing the "-" sign to "+") corresponds to pH = p$K'$ + lg[A^-]/[HA], meaning pH = p$K'$ + lg([proton acceptor]/[proton donor]). This is the Henderson-Hasselbalch equation, which explains why the p$K'$ value of a weak acid is numerically equal to the pH of its solution at the midpoint of titration: at this point, [HA] = [A^-]. Consequently, pH = p$K'$ + lg 1.0 = p$K'$ + 0 = p$K'$.

Furthermore, this equation makes it possible to calculate the p$K'$ value for a given pH or to determine the ratio between the molar concentrations of the proton donor and proton acceptor at any pH value. The equation shows that the concentration of H⁺ ions in a buffer solution depends not only on the dissociation constant $K'$ of the weak acid or weak base, but also on the concentration of a salt that shares a common anion with the acid or a common cation with the base. The higher the salt concentration in "weak acid - salt" buffer solutions, the lower the concentration of H⁺ ions within them. At equal concentrations of the acid and its salt, the concentration of H⁺ ions in such solutions approaches a value equal to the acid dissociation constant:

From the foregoing, it is clear that the ability of a buffer solution to maintain a constant pH during titration is limited by its buffer capacity. The unit of buffer capacity is defined as the amount of strong acid or strong base (in moles of equivalents per 1 liter of solution) required to change the pH by one unit. Buffer capacity ($B$) is determined as follows:

where pH₁ and pH₂ are the extreme pH Limits of the buffer, and $C$ is the concentration of the acid (or alkali). The buffer capacity of a solution increases as the concentration of its components rises and as The ratio of the acid to its salt (or the alkali to its salt) approaches unity. The total buffer capacity of arterial blood is 25.3 mmol/L, while in venous blood it is slightly lower at 24.3 mmol/L.

Maintaining the constancy of body fluid pH is of paramount importance for physiological processes for at least three reasons:

1) H⁺ ions exert a catalytic effect on numerous biochemical transformations;

2) Enzymes and Hormones exhibit biological activity only within a specific pH range;

3) even minor changes in H⁺ concentration in blood and intercellular fluid significantly affect the magnitude of osmotic pressure in these fluids.