LEHNINGER PRINCIPLES OF BIOCHEMISTRY - VOL. 1. THE FOUNDATIONS OF BIOCHEMISTRY: STRUCTURE AND CATALYSIS - 2011

PART I. STRUCTURE AND CATALYSIS

Class="center">I believe that The Use of structural chemistry Methods to solve physiological problems will demonstrate that The Role of Hydrogen Bonds in physiology surpasses that of any other structural feature.

Linus Pauling, The Nature of the Chemical Bond, 1939

2. WATER

Water is the most abundant substance in living systems, accounting for 70% or more of the mass of most organisms. The first living organisms on Earth undoubtedly originated in an aqueous environment, and the course of evolution has been shaped by The properties of the watery milieu in which life began.

This chapter begins by examining the physical and Chemical properties of water, to which all structural elements and Functions of the living Cell are adapted. The forces of attraction between water molecules and its weak ionization are of paramount importance for the Structure and function of Biomolecules. In our Discussion of ionization, we will cover equilibrium constants, pH, and titration curves, and show how aqueous solutions of weak acids and bases—along with their salts—act as buffers to prevent pH fluctuations in biological systems. The water molecule and its ionization products (H⁺ and OH^-) profoundly influence the structure, assembly process, and properties of all cellular components, including Proteins, Nucleic Acids, and Lipids. The solvent properties of water, including its ability to form hydrogen bonds with both itself and solute molecules, critically influence the noncovalent interactions responsible for the strength and Specificity of biomolecular "recognition."

2.1. Weak Interactions in Aqueous Systems

The presence of hydrogen bonds results in a high internal cohesion among water molecules, keeping it liquid at room Temperature and facilitating The formation of the highly ordered crystal lattice of ice (the solid state). Polar biomolecules dissolve readily in water because they can displace water-water interactions to form energetically more favorable water-solute interactions. Conversely, nonpolar substances interfere with water-water interactions without being able to form water-solute interactions; in other words, nonpolar substances are poorly soluble in water. In aqueous solutions, nonpolar substances frequently form clusters.

Hydrogen bonds, ionic interactions, hydrophobic (from the Greek phobos, "fear") interactions, and Van der Waals interactions are individually quite weak, yet together they exert a profound influence on the three-dimensional structures of proteins, nucleic acids, Polysaccharides, and Membrane Lipids.

Table 2-1. Melting point, boiling point, and heat of vaporization of some common Solvents

*The amount of thermal energy required to convert 1.0 g of a liquid at its boiling point and atmospheric pressure into a gas at the same temperature and pressure. This value serves as a measure of the energy needed to overcome attractive forces between molecules in the liquid phase.

The unusual properties of water stem from the presence of hydrogen bonds

Compared with most other liquids, water has unusually high melting and boiling points and a high heat of vaporization (Table 2-1). These characteristics are the result of strong attractive forces between neighboring water molecules, which give liquid water a high degree of internal cohesion. THE ORIGIN OF these intermolecular interactions can be understood by examining the electronic STRUCTURE OF THE H₂O molecule.

Each hydrogen atom in a water molecule shares an electron pair with the central oxygen atom. The molecular geometry is determined by the shape of the outer electron orbitals of oxygen, which are structurally analogous to the bonding sp³ orbitals of carbon (Fig. 1-14). These orbitals are arranged in an irregular tetrahedron, with a hydrogen atom at each of two corners and unshared electron pairs of oxygen at the other two corners (Fig. 2-1a). The H-O-H bond angle is 104.5°—slightly less than the 109.5° angle of a regular tetrahedron—due to the mutual repulsion of the unshared electron pairs of oxygen.

Fig. 2-1. Structure of a water molecule. (a) The dipolar Nature of the water molecule illustrated with a ball-and-stick model. Dashed lines indicate nonbonding orbitals. Electron pairs located in the outer electron orbitals form an irregular tetrahedron around the oxygen atom. Each hydrogen atom carries a partial positive charge (δ⁺), and the oxygen atom carries a partial negative charge (2δ^-). (b) Two H₂O molecules joined by a Hydrogen bond (indicated here and throughout as three light blue bars) connecting the oxygen atom of the upper molecule to the hydrogen atom of the lower molecule. Hydrogen bonds are longer and weaker than covalent O-H bonds.

The oxygen Nucleus attracts electrons more strongly than the hydrogen nucleus (proton), meaning that oxygen is more electronegative. As a result, the shared electron pair between oxygen and hydrogen is not located at an equidistant position—the electrons spend more time closer to the oxygen atom than to the hydrogen atom. This Asymmetry results in the formation of two electric dipoles within the water molecule, oriented along each H-O bond. Consequently, a partial positive charge (δ⁺) is concentrated on the hydrogens, and a partial negative charge (equal in magnitude to the sum of the two partial positive charges, 2δ^-) is concentrated on the oxygen. This accounts for the electrostatic attraction between the oxygen atom of one water molecule and the hydrogen atom of another (Fig. 2-1b), known as a hydrogen bond. In this book, we represent hydrogen bonds by three parallel light-blue bars, as in Figure 2-1b.

Hydrogen bonds are relatively weak. The dissociation energy (the energy required to break the bond) of a hydrogen bond in liquid water is 23 kJ/mol, whereas the dissociation energy of a covalent O-H bond in water is 470 kJ/mol, and that of a covalent C-C bond is 348 kJ/mol. Hydrogen bond formation derives about 10% of its strength from covalent interactions (resulting from orbital overlap) and about 90% from Electrostatic Interactions. At room temperature, the thermal energy of an aqueous solution (the kinetic energy of individual atoms and molecules in motion) is of the same order of magnitude as the dissociation energy of a hydrogen bond. Raising the temperature by heating water causes individual molecules to move faster. At any given moment, most molecules in liquid water participate in hydrogen bonding, yet the lifetime of any individual hydrogen bond is only 1 to 20 picoseconds (1 ps = 10^-12 s). When one hydrogen bond breaks after 0.1 ps, another forms—either with the same partner or a new one. These short-lived clusters of water molecules held together by hydrogen bonds have been aptly termed "flickering clusters." The network of hydrogen bonds among molecules accounts for the high internal cohesion of liquid water. Extended networks of hydrogen-bonded water molecules also encompass solute molecules (such as proteins and nucleic acids), enabling large molecules to interact with one another over distances of several nanometers without direct physical contact.

Because the arrangement of electron orbitals around the oxygen atom is nearly tetrahedral (Fig. 2-1a), each water molecule can theoretically form hydrogen bonds with a maximum of four neighboring water molecules. Since molecules in the liquid state at atmospheric pressure are in continuous motion, each molecule is estimated to form hydrogen bonds with an average of 3.4 other molecules at any given instant. By contrast, in ice, molecules are spatially fixed and bonded to the maximum possible number of neighbors, forming a regular spatial lattice (Fig. 2-2). A substantial amount of thermal Energy is required to break enough hydrogen bonds to disrupt the crystal lattice of ice, which is reflected in the relatively high melting point of water (Table 2-1). Melting ice or evaporating water requires thermal energy:

H₂O (solid) —> H₂O (liquid)

 ∆Н = +5.9 kJ/mol

H₂O (liquid) —> H₂O (gas)

 ∆Н = +44.0 kJ/mol

Fig. 2-2. Hydrogen bonding between water molecules in ice. Each water molecule in ice is hydrogen-bonded to four other water molecules (the maximum possible number for water), forming a regular crystal lattice. In liquid water at room temperature and atmospheric pressure, each water molecule is hydrogen-bonded to approximately 3.4 other molecules. The crystal lattice of ice is more open (ice has a lower density than water), which is why ice floats on water.

During melting or evaporation, the Entropy of the system increases because the highly organized structure of the molecules (ice) transitions into a less organized (water) or completely disordered state (steam). At room temperature, the melting of ice and evaporation of water occur spontaneously: the tendency of water molecules to associate via hydrogen bonding is weaker than the thermodynamic advantage of transitioning to a chaotic state. Recall that for a process to occur spontaneously, The change in Free energy (∆G) must be negative: ∆G = ∆Н — T∆S, where ∆G represents the driving force of the process, ∆Н is the enthalpy change associated with the formation and breaking of bonds, and ∆S reflects the change in system order. Since ∆Н > 0 for melting and evaporation, it is precisely the increase in entropy (∆S) that makes ∆G negative and drives these transformations.

Water forms hydrogen bonds with polar solutes

Hydrogen bonds exist not only in water. They readily form between an electronegative atom (a hydrogen acceptor, typically oxygen or nitrogen) and a hydrogen atom covalently bonded to another electronegative atom (a hydrogen donor) in the same or a different molecule (Fig. 2-3). Hydrogen atoms covalently bonded to carbon do not participate in hydrogen bonding because the electronegativity of carbon is only slightly higher than that of hydrogen, making the C-H bond practically nonpolar. This difference explains the relatively high boiling point of butyl alcohol CH₃(CH₂)₂CH₂ОH (117 °C) compared to the low boiling point of butane СН₃(СН₂)₂СН₃ (-0.5 °C). The butyl alcohol molecule contains a polar hydroxyl group that can participate in intermolecular hydrogen bonding. Uncharged yet polar biomolecules (such as sugars) dissolve readily in water due to the stabilizing effect of hydrogen bonds formed between the hydroxyl groups or carbonyl oxygens of the sugar and polar water molecules. Alcohols, aldehydes, ketones, and compounds containing N-H bonds form hydrogen bonds with water molecules (Fig. 2-4) and are generally soluble in water.

Fig. 2-3. Common types of hydrogen Bonds Found in biological systems. The hydrogen acceptor is typically an oxygen or nitrogen atom; the hydrogen donor is another electronegative atom.

Fig. 2-4. Selected hydrogen bonds of importance in biological systems.

Hydrogen bonds are strongest when the relative orientation of the bonded molecules maximizes electrostatic interaction energy. This condition is met when the hydrogen atom and the two atoms sharing it lie in a straight line—that is, when the acceptor lies on the axis of the covalent bond between the donor atom and the hydrogen atom (Fig. 2-5)—and when the positively charged hydrogen ion is positioned exactly between the partial negative charges on the oxygen atoms. In other words, hydrogen bonds are directional and, as a result, can hold both bonded molecules or groups in a specific orientation. As we will see later, this property of hydrogen bonds helps stabilize the precise three-dimensional structures characteristic of proteins and nucleic acids, which contain numerous intramolecular hydrogen bonds.

Fig. 2-5. Directionality of the hydrogen bond. Maximum attraction between partial electrical charges (Fig. 2-1) occurs when the three atoms involved in the hydrogen bond (O, H, and O in this case) lie on a straight line. If the groups participating in the hydrogen bond are rigidly fixed (e.g., as PARTS OF THE same protein molecule) and such ideal alignment is impossible, a weaker bond is formed.

Electrostatic interactions exist between water and charged species

Water is a polar solvent. Most biomolecules, which are typically charged or polar substances, dissolve readily in it (Table 2-2). Substances that dissolve easily in water are termed hydrophilic (from Greek, meaning water-loving). Conversely, nonpolar solvents (such as chloroform and benzene) are poor solvents for polar biomolecules, but they readily dissolve substances with hydrophobic properties—that is, nonpolar molecules like lipids and Waxes.

Water dissolves salts such as NaCl by hydrating and stabilizing Na⁺ and Cl^- ions, weakening the electrostatic interactions between them and thereby preventing their reassociation into a crystal lattice (Fig. 2-6). The same principles apply to the dissolution of charged biomolecules containing functional groups such as an ionized carboxyl group (-СОО-), a protonated amino group (-NH₃⁺), or phosphoric acid esters and anhydrides. Water readily dissolves such substances by replacing solute-solute hydrogen bonds with solute-water hydrogen bonds, while also screening the electrostatic interactions between solute molecules.

Water is very convenient for evaluating electrostatic interactions between dissolved ions because it has a high dielectric constant (a physical parameter reflecting the number of dipoles in a solvent). The strength of ionic interactions in solution (F) depends on the magnitude of the charges (Q), the distance between the charged groups (r), and the Dielectric Constant of the solvent (ε, a dimensionless quantity) in which the interaction occurs:

Fig. 2-6. Water as a solvent. The high solubility of many crystalline salts in water is due to the Hydration of the ions composing those salts. The NaCl crystal lattice breaks down as water molecules surround Na⁺ and Cl^- ions. The ionic charges are partially neutralized, and the electrostatic attraction required to maintain the lattice is weakened.

The dielectric constant of water at 25 °C is 78.5, whereas for a nonpolar solvent like benzene, ε = 4.6. Thus, ionic interactions between dissolved ions are much stronger in nonpolar than in polar environments. Because ionic interaction strength depends inversely on distance (1/r²), ionic attraction or repulsion in water operates only over very short ranges—typically between 10 and 40 nm, depending on the electrolyte concentration.

Table 2-2. Examples of polar, nonpolar, and amphiphilic biomolecules (in ionic form at pH 7)

Entropy increases upon the dissolution of crystalline substances

When NaCl dissolves, Na⁺ and Cl^- ions leave the crystal lattice and gain greater translational freedom (Fig. 2-6). The high solubility of salts like NaCl in water is largely driven by an increase in the entropy of the system (greater disorder). In thermodynamic terms, solution formation is favored by the change in free energy: ∆G = ∆Н — T∆S, where ∆Н has a small positive value and T∆S has a large positive value, resulting in ∆G < 0.

Nonpolar gases are sparingly soluble in water

The molecules CO₂, O₂, and N₂, which are critical to biological systems, are nonpolar. In oxygen and nitrogen molecules, electrons are shared equally between the two atoms. In the CO₂ molecule, each C=O bond is polar, but the two dipoles have opposite directions and neutralize each other (Table 2–3). The transition of molecules from a disordered state in the gas phase to an aqueous solution disrupts their freedom of movement as well as the motion of water molecules, i.e., it is accompanied by a decrease in entropy. The nonpolar nature of these molecules and the decrease in entropy during the phase transition from gas to liquid account for the very poor solubility of these compounds (Table 2–3). Some organisms possess water-soluble carrier proteins (such as Hemoglobin and Myoglobin) that facilitate Oxygen transport. Carbon dioxide forms carbonic acid (H₂CO₃) in water and is transported either as the bicarbonate ion (HCO₃^-) in a free state (bicarbonate is highly soluble in water: ~100 g/L at 25 °C) or in a complex with hemoglobin. Gases such as NH₃, NO, and H₂S are also important for the vital activity of certain organisms; they are polar, highly soluble in water, and ionized in aqueous solutions.

Table 2–3. Solubility of some gases in water

* Arrows indicate electric dipoles: the direction from the region with a partial positive charge (δ⁺, not shown here) to the region with a partial negative charge (δ^-).

** Note that polar molecules dissolve in water better even at low temperatures than nonpolar molecules do at elevated temperatures.

Nonpolar substances cause energetically unfavorable changes in water structure upon dissolution

When water is mixed with benzene or hexane, two phases are formed. Nonpolar compounds (benzene and hexane) are hydrophobic, meaning they are unable to engage in energetically favorable interactions with water molecules and instead disrupt the hydrogen bonds existing between water molecules. All molecules or ions in an aqueous solution disrupt these bonds, but polar or charged substances, such as NaCl, compensate for the disruption of water-water bonds by forming new water-solute bonds. The overall enthalpy change (∆H) upon dissolution of such molecules is usually small. In contrast, hydrophobic substances do not compensate for the disruption of the hydrogen-bonding network between water molecules; consequently, their dissolution is accompanied by a small increase in enthalpy. Breaking the hydrogen bonds between water molecules requires a certain amount of energy, which must be drawn from the surroundings. Furthermore, dissolving hydrophobic substances in water leads to a significant decrease in entropy. Water molecules in the immediate vicinity of a nonpolar molecule are constrained into a specific orientation, forming a highly ordered shell around each solute molecule. Water molecules organized in this manner do not have as regular a structure as clathrates (crystals of nonpolar substances and water), but the effect in both cases is the same—the ordering of water molecules reduces entropy. The number of ordered water molecules, and consequently the magnitude of the entropy decrease, is proportional to the surface area of the hydrophobic substance encased within the shell of water molecules. Thus, the free-energy change accompanying the dissolution of a nonpolar substance in water is unfavorable to the process: ∆G = ∆H — T∆S, where ∆H is a positive number and ∆S is a negative number, hence ∆G > 0.

Amphiphilic compounds possess both polar (charged) and nonpolar groups (Table 2–2). When an amphiphilic compound is mixed with water, its polar hydrophilic regions tend to dissolve, whereas its nonpolar hydrophobic regions attempt to avoid contact with the solvent (Fig. 2–7a). The nonpolar groups aggregate to form clusters with minimal surface area, while the polar groups orient themselves to maximize their contact with water (Fig. 2–7b). As a result, stable structures of amphiphilic substances called micelles are formed in water, which may consist of hundreds or thousands of molecules. The forces holding the nonpolar groups of solute molecules together are termed hydrophobic interactions. This force is not related to any intrinsic attraction between nonpolar residues. Rather, the interaction is the result of the system's tendency to achieve maximum thermodynamic stability by minimizing the number of ordered water molecules surrounding the hydrophobic portions of the solute molecules.

Fig. 2–7. Behavior of amphiphilic compounds in aqueous solution, a) Long-chain Fatty acids have hydrophobic alkyl chains surrounded by a layer of ordered water molecules. b) Micelle formation minimizes the contact of the Hydrophobic surface of fatty acids with water; moreover, in such a structure, the minimum number of ordered water molecules is involved in forming the shell around the nonpolar molecules. Micelle stabilization occurs due to an energy gain resulting from the release of Immobilized Water molecules.

Many biomolecules are amphiphilic: proteins, pigments, certain Vitamins, as well as membrane sterols and Phospholipids, all have both polar and nonpolar regions. The structures formed by such molecules are held together by hydrophobic interactions between nonpolar groups. Hydrophobic interactions among lipids, as well as between lipids and proteins, represent a fundamental principle of biological membrane architecture. Hydrophobic interactions between nonpolar Amino Acids also play a role in stabilizing the three-dimensional structure of proteins.

The formation of hydrogen bonds between water and polar compounds also leads to the ordering of water molecules, but the energetic effect is smaller than in the dissolution of nonpolar substances. Part of the driving force directed toward binding a polar substrate (reactant) to the complementary polar surface of an enzyme is expended on increasing entropy, as the enzyme displaces ordered water molecules from the substrate surface and the substrate displaces ordered water molecules from the enzyme surface (Fig. 2–8).

Fig. 2–8. The release of ordered water molecules promotes the Formation of the enzyme-substrate complex. A layer of ordered water molecules exists around both the enzyme molecule and the substrate molecule. The binding of the enzyme to the substrate results in the release of a certain number of ordered water molecules. The resulting increase in entropy provides the thermodynamic driving force for the formation of the enzyme-substrate complex (see p. 282).

Van der Waals interactions are weak forces of interatomic attraction

When two uncharged atoms approach each other too closely, their surrounding electron clouds interact. Random fluctuations in THE POSITION OF electrons around one nucleus can create an electric dipole, which is capable of inducing an opposing electric dipole in a neighboring atom. The two dipoles weakly attract each other, causing the nuclei to draw closer. Such weak attractive forces are called Van der Waals attraction (or London forces). As the two nuclei approach further, a counteracting repulsion force between their electron clouds arises. At the point where the net attraction is maximal, the nuclei are said to be in Van der Waals contact. Each atom is characterized by a Van der Waals radius that determines how closely that atom allows another atom to approach (Table 2–4). In the space-filling molecular models presented in this book, the sizes of all atoms are scaled in proportion to their Van der Waals radii.

Table 2–4. Van der Waals radii and covalent radii (for single bonds) of selected elements

Sources: Values for Van der Waals radii are taken from Chauvin, R. (1992) Explicit periodic trend of van der Waals radii. J. Phys. Chem. 96, 9194–9197. Values for covalent radii are taken from Pauling, L. (1960) The Nature of the Chemical Bond, 3rd ed., Cornell University Press, Ithaca, NY.

Note: Van der Waals radii describe the spatial dimensions of an atom. If two atoms are linked by a covalent bond, the atomic radii in the bonding region are smaller than the Van der Waals radii because the bonded atoms share a pair of electrons. The distance between nuclei in Van der Waals interactions or covalent bond formation is approximately equal to the sum of the Van der Waals or covalent radii of the two atoms, respectively. Thus, the length of a single C–C bond is approximately 0.077 nm + 0.077 nm = 0.154 nm.

Weak interactions play a critically important role in the structure and function of macromolecules

The noncovalent interactions described above (hydrogen bonds, as well as ionic, hydrophobic, and Van der Waals interactions) (Table 2–5) are much weaker than covalent bonds. The bond dissociation energy is 350 kJ/mol for a single C–C bond and 410 kJ/mol for a C–H bond, whereas for Van der Waals forces it is only about 4 kJ/mol. Hydrophobic interactions are also much weaker than covalent bonds, but they are significantly enhanced in a polar solvent (e.g., in a concentrated salt solution). The strength of ionic interactions and hydrogen bonds depends on solvent polarity and the orientation of the hydrogen-bonded atoms, but it is always weaker than that of covalent bonds. The thermal energy of molecules in an aqueous solution at 25 °C is typically of the same order of magnitude as the strength of weak interactions; consequently, the probability of solute-solvent (water) interactions is practically the same as that of solute-solute interactions. For this reason, hydrogen bonds as well as ionic, hydrophobic, and Van der Waals interactions are constantly forming and breaking.

Table 2–5. Four types of noncovalent ("weak") interactions between biomolecules in aqueous solution

Although these four types of interactions are individually weaker than a covalent bond, the cumulative effect of multiple interactions of this kind can be quite substantial. For example, the noncovalent binding of an enzyme to its substrate may involve a combination of hydrogen bonds, ionic interactions, hydrophobic interactions, and Van der Waals forces. The formation of each of these weak bonds contributes to the overall decrease in the Free energy of the system. The stability of a noncovalent interaction, such as a hydrogen bond between a small molecule and a macromolecule, can be calculated from the bond energy. The stability, described by the binding Equilibrium Constant (see below), depends exponentially on the bond energy. For the dissociation of a complex of two biomolecules (e.g., an enzyme and a substrate) held together by multiple weak forces, all these interactions must be disrupted simultaneously. Because these interactions are continually forming and breaking, the simultaneous rupture of all bonds is highly improbable. Thus, a molecular structure held together by 5 to 20 weak interactions turns out to be far more stable than might be expected from a simple summation of small binding energies.

In macromolecules such as proteins, DNA, and RNA, there are numerous sites for hydrogen bonding, as well as hydrophobic, ionic, and van der Waals interactions, meaning that the cumulative effect of these weak forces can be immense. The most stable (i.e., native) structure of macromolecules is typically characterized by the maximum number of realized weak interactions. It is in accordance with this principle that polypeptide or polynucleotide chains fold into their characteristic spatial Conformations. The binding of Antigens to their specific Antibodies similarly relies on the cumulative effect of numerous weak interactions. As noted above, The energy released during the noncovalent binding of an enzyme to its substrate serves as the driving force for the catalytic reaction. The binding of a hormone or neurotransmitter to a cellular receptor is also the result of weak interactions. Due to the large size of enzyme and receptor molecules (compared to their substrates or ligands), their surfaces feature abundant sites for weak interactions to take place. Molecular complementarity between interacting macromolecules implies the potential for weak interactions among polar, charged, and hydrophobic groups on their surfaces.

When studying Protein Structure—such as hemoglobin (Fig. 2-9), for instance—using X-ray crystallography (see Box 4-5, p. 196), researchers frequently find that water molecules are bound to the protein so tightly that they behave as part of the crystal lattice. An analogous phenomenon is observed in RNA or DNA crystals. These bound water molecules, which can be detected in aqueous solution by nuclear magnetic Resonance spectroscopy, exhibit properties entirely distinct from those of bulk water. For example, they lack osmotic activity (see below). Bound water molecules play a vital role in the function of many proteins. Thus, in a key photosynthetic reaction driven by light, protons traverse a biological membrane as electrons are transferred along a series of specialized carrier proteins (Fig. 19-60). One of these proteins—cytochrome f—features a chain of five bound water molecules (Fig. 2-10) that may facilitate proton translocation across the membrane via a mechanism termed "proton hopping" (see below). Another such light-driven proton pump, Bacteriorhodopsin, apparently utilizes a specifically oriented chain of bound water molecules for transmembrane proton transfer (Fig. 19-67).

Fig. 2-9. Water binding to a hemoglobin molecule (PDB ID 1A3N). The crystal structure of the hemoglobin molecule (a) with bound water molecules (red spheres) and (b) without water molecules. The water molecules are bound to the protein so tightly that, in X-ray crystallographic analysis, they behave as though they are part of the crystal lattice. The two α-subunits of hemoglobin are shown in gray, and the two β-subunits in blue. Each subunit is bound to a heme group (depicted as red sticks; from this perspective, the heme is visible only in the β-subunits). The structure and function of hemoglobin are discussed in Chapter 5.

Fig. 2-10. A chain of water molecules in cytochrome f. Water molecules are bound within the proton channel of the membrane protein cytochrome f, which is a component of the photosynthetic apparatus in METABOLISM/14.html">Chloroplasts (Fig. 19-64). Five water molecules are linked to one another and to Functional groups of the protein via hydrogen bonds; this interaction involves backbone atoms of valine, Proline, Arginine, and Alanine residues, as well as side-chain atoms from three asparagine residues and two glutamine residues. The protein molecule is bound to a heme group (Fig. 5-1), in which the iron ion facilitates electron transfer during Photosynthesis. The flow of electrons is coupled to proton translocation across the membrane, potentially via a "hopping" mechanism of proton transfer along the chain of bound water molecules (Fig. 2-13).

Solutes alter the Properties of Water

Any solute alters Certain physical properties of the solvent (water), notably its vapor pressure, boiling point, freezing (melting) point, and osmotic pressure. These properties are referred to as colligative properties because the underlying cause for all these alterations is the same: the concentration of water in a solution is lower than that of pure water. Changes in colligative properties are independent of the Chemical Nature of the solute and depend solely on the concentration of solute particles (molecules or ions) per unit volume of solvent. For instance, NaCl dissociating in an aqueous solution alters the osmotic pressure twice as much as an equivalent molar amount of glucose, which does not dissociate in water.

Water molecules tend to migrate from a region of higher water concentration to a region of lower water concentration, in accordance with the drive of any natural system toward increased disorder. When two aqueous solutions of differing solute concentrations are separated by a semipermeable membrane (which allows water molecules to pass while blocking solutes), water molecules diffuse from the region of higher water concentration to that of lower concentration, thereby generating osmotic pressure (Fig. 2-11). Osmotic pressure (П) is defined as the force that must be applied to counteract the Movement of water (Fig. 2-11c) and is calculated using the van 't Hoff equation:

П = icRT

where R is the universal gas constant and T is the absolute temperature. The term ic, which is the product of the molar concentration of the solution (c) and the van 't Hoff factor (i), is called the osmolarity of the solution and reflects the degree to which the solute dissociates into ions. In a dilute NaCl solution, the salt dissociates completely into Na⁺ and Cl^- ions, which doubles the number of solute particles, i.e., i = 2. For all non-dissociating substances, i = 1. The Osmotic Pressure of a solution containing multiple (n) solutes is the sum of the osmotic pressures exerted by each individual solute:

П = RT (i₁c₁ + i₂c₂ + ... ⁺i_nc_n)

Fig. 2-11. Osmosis and the definition of osmotic pressure. (a) Initial state. The dialysis tubing contains an aqueous solution, and the beaker contains pure water. The semipermeable membrane permits the passage of water but not solute. Water flows from the beaker into the tubing in an effort to equalize water concentrations on both sides of the membrane. (b) Final state. Water has entered the solution of molecules that cannot cross the membrane, resulting in dilution of the solution. At equilibrium, the hydrostatic pressure of the solution Column of height h balances the osmotic pressure, which reflects the tendency of water to enter the region of lower water concentration. (c) Osmotic pressure (П) is defined as the force that must be applied to the piston to lower the solution level in the tubing back to the water level in the beaker. This force is proportional to the height of the liquid column (h) in the tubing in state (b).

The phenomenon of osmosis—manifested as the movement of water across a semipermeable membrane driven by an osmotic pressure gradient—plays a critical role in the life of most Cells. Plasma Membranes are far more permeable to water than to most other small molecules, ions, and macromolecules. Water entry into cells occurs partly through simple Diffusion of Water molecules across The Lipid Bilayer and partly via membrane protein channels known as Aquaporins (see Fig. 11-46), which facilitate the selective transport of water. Solutions with an osmolarity equal to that of the Cell Cytoplasm are termed isotonic to The Cell. The water content of a cell placed in an isotonic solution neither decreases nor increases (Fig. 2-12).

Fig. 2-12. Effect of External osmolarity on water movement across The Plasma Membrane. When a cell in osmotic equilibrium with its surroundings (i.e., in an isotonic solution) (a) is transferred to a hypertonic (b) or hypotonic solution (c), water fluxes across the plasma membrane in an attempt to equilibrate intracellular and extracellular osmolarity.

In a hypertonic solution, whose osmolarity exceeds that of the Cytosol, cells shrink as water flows out of them. Conversely, in a hypotonic solution, whose osmolarity is lower than that of the cytosol, cells swell as water rushes inward. Under physiological conditions, the concentration of biomolecules and ions inside cells is typically higher than in their external environment, meaning that osmotic pressure drives water into cells. Unless this pressure is counterbalanced in some way, the incoming water will stretch the plasma membrane and eventually rupture the cell (osmotic lysis).

Cells possess several mechanisms to prevent such a catastrophe. The plasma membranes of bacterial and plant cells are encased in a rigid, stretch-resistant Cell wall strong enough to withstand osmotic pressure and prevent osmotic lysis. Some freshwater Protozoa inhabiting hypotonic environments feature specialized Organelles called contractile vacuoles that actively pump water out of the cell. In Multicellular animals, Blood Plasma and interstitial fluid (the extracellular fluid bathing Tissues) have an osmolarity very close to that of the cytosol. High concentrations of albumin and other proteins make a significant contribution to blood plasma osmolarity. Furthermore, to maintain osmotic equilibrium with their environment, cells actively pump Na⁺ and certain other ions out into the interstitial fluid.

Because The Effect of solutes on osmolarity depends on the number of dissolved particles rather than their mass, macromolecules (proteins, nucleic acids, polysaccharides) exert a much smaller effect than an equivalent mass of their constituent monomeric units. For example, one gram of a polysaccharide consisting of 1,000 glucose residues has the exact same osmotic impact as one milligram of glucose. Cellular energy reserves are stored as polysaccharides (starch or Glycogen) rather than as glucose or Other simple sugars, thereby avoiding an astronomical rise in osmotic pressure within the storage cell.

In plants, osmotic pressure serves to maintain mechanical rigidity. Water enters plant cells thanks to the very high concentration of solutes within the plant vacuole (Fig. 2-12). The resulting osmotic pressure exerted against The cell wall (turgor pressure) increases, conferring structural rigidity to cells, tissues, and the plant as a whole. Salad greens wilt because water loss leads to a drop in turgor. The phenomenon of osmosis is also crucial to consider in laboratory practice. For instance, Mitochondria, chloroplasts, and Lysosomes are bounded by semipermeable membranes. The isolation of organelles from disrupted Cells must be carried out in an isotonic medium (see Fig. 1-8) to prevent the influx of excess water into the organelles, which would lead to Swelling and lysis. Buffers used for Cell Fractionation typically contain a sufficient concentration of sucrose or other inert solutes to protect organelles from osmotic lysis.

Example 2-1. Osmotic Pressure in an Organelle. I

Assume that the major solutes within intact lysosomes are KCl (~0.1 M) and NaCl (~0.03 M). What must the sucrose concentration of the extraction medium be at room temperature (25 °C) to prevent lysosomal swelling and lysis?

Solution. We need to find the sucrose concentration that provides an osmotic pressure equal to that generated by the KCl and NaCl present inside the lysosomes. To calculate the osmotic pressure, we use the van 't Hoff equation:

П = RT (i₁c₁ + i₂c₂ + i₃c₃ + ... + i_nc_n)

where R is the universal gas constant (8.315 J/(mol·K)), T is the absolute temperature (in kelvins), c₁, c₂, and c₃ are the molar concentrations of the solutes, and i₁, i₂, and i₃ are the number of particles formed in solution upon dissociation of each substance (for KCl and NaCl, i = 2).

Let us calculate the osmotic pressure in the lysosome:

П_lysosome = RT(i_KClC_KCl + i_NaClC_NaCl) = RT [(2)(0.03 mol/L) + (2)(0.1 mol/L)] = RT (0.26 mol/L)

Since the solute concentrations are known only to one decimal place, we obtain П_lysosome = RT (0.3 mol/L).

The osmotic pressure of a sucrose solution is determined by the formula:

П_sucrose = RT (i_sucrose C_sucrose)

In this case, i_sucrose = 1, since sucrose is not ionized in water. Thus,

П_sucrose = RT C_sucrose

Equating the osmotic pressure of the lysosomal contents to the osmotic pressure of the sucrose solution:

П_sucrose = П_lysosome

RT (c_sucrose) = RT (0.3 mol/L)

c_sucrose = 0.3 mol/L

Thus, the required concentration of sucrose (mol. mass 342) is (0.3 mol/L)(342 g/mol) = 102.6 g/L. Or, considering only significant figures, c_sucrose = 0.1 kg/L.

Example 2-2. Osmotic Pressure in an Organelle. II

Suppose that to balance the osmotic pressure inside the lysosomes (see Example 2-1) we decide to use a polysaccharide solution, such as glycogen. Assuming that the linear polymer consists of 100 glucose units, calculate the amount of polymer required to achieve the same osmotic pressure as in Example 2-1. The M_r of the polymer consisting of glucose residues is ~18,000, and, like sucrose, it is not ionized in water.

Solution. As we found in Example 2-1:

П_sucrose = RT (0.3 mol/L)

Similarly,

П_glycogen = RT (i_glycogen c_glycogen) = RT (c_glycogen)

For a glycogen solution with the same osmotic pressure as the sucrose solution, we can write:

П_glycogen = П_sucrose

RT (c_glycogen) = RT (0.3 mol/L)

c_glycogen = 0.3 mol/L = (0.3 mol/L) (18,000 g/mol) = 5.4 kg/L

Taking into account only significant figures, we get the answer: c_glycogen = 5 kg/L — an incredibly high concentration!

As we will see later (p. 353), Liver and Muscle cells store CARBOHYDRATES not as low-molecular-weight sugars such as glucose or sucrose, but as the high-molecular-weight polymer glycogen, which minimizes the effect on cytosolic osmolarity.

Summary of Section 2.1. Weak Interactions in Aqueous Media

■ The significant electronegativity difference between H and O atoms accounts for the high polarity of the water molecule, enabling it to form hydrogen bonds with other water molecules and with solute molecules. These hydrogen bonds are transient, predominantly electrostatic in nature, and weaker than covalent bonds. Water serves as an excellent solvent for polar (hydrophilic) substances, with which it forms hydrogen bonds, and for charged substances, with which it engages in electrostatic interactions.

■ Nonpolar (hydrophobic) substances dissolve poorly in water and are incapable of forming hydrogen bonds with the solvent. When hydrophobic substances are present in an aqueous solution, an energetically unfavorable ordering of water molecules occurs at their surface. To minimize the surface area exposed to water, nonpolar substances such as lipids aggregate into micelles, orienting their hydrophobic residues toward the interior while their polar groups interact with the surrounding water. The forces that drive hydrophobic particles to associate in a polar solvent are referred to as hydrophobic interactions.

■ Numerous weak noncovalent interactions play a crucial role in maintaining the native Conformation of Macromolecules such as proteins and nucleic acids. The most stable conformation of a biomolecule is characterized by the maximization of internal hydrogen bonds as well as those formed between the molecule and the solvent. Furthermore, hydrophobic residues cluster within the interior of the biomolecule, thereby minimizing their contact with water.

■ The Physical Properties of aqueous solutions strongly depend on solute concentration. When two aqueous solutions are separated by a semipermeable membrane (such as the plasma membrane separating a cell from its external environment), water tends to cross the membrane to equalize the osmolarity of both solutions. The pressure exerted by water in this process is known as osmotic pressure.