Weak Interactions in Aqueous Systems: -Solutes Affect the Colligative Properties of Aqueous Solutions

Solutes of all kinds alter certain physical properties of the solvent, water: its vapor pressure, boiling point, melting point (freezing point), and osmotic pressure. These are called colligative (“tied together”) proper ties, because the effect of solutes on all four properties has the same basis: the concentration of water is lower in solutions than in pure water. The effect of solute con centration on the colligative properties of water is in dependent of the chemical properties of the solute; it depends only on the number of solute particles (molecules, ions) in a given amount of water. A compound such as NaCl, which dissociates in solution, has twice the effect on osmotic pressure, for example, as does an equal number of moles of a no dissociating solute such as glucose. Solutes alter the colligative properties of aqueous solutions by lowering the effective concentration of water. For example, when a significant fraction of the molecules at the surface of an aqueous solution are not water but solute, the tendency of water molecules to escape into the vapor phase—that is, the vapor pres sure—is lowered (Fig. 2–11). Similarly, the tendency of water molecules to move from the aqueous phase to the surface of a forming ice crystal is reduced when some of the molecules that collide with the crystal are solute, not water. In that case, the solution will freeze more slowly than pure water and at a lower temperature. For a 1.00 molal aqueous solution (1.00 mol of solute per 1,000 g of water) of an ideal, nonvolatile, and nondis sociating solute at 101 kPa (1 atm) of pressure, the freezing point is 1.86 C lower and the boiling point is 0.543 C higher than for pure water. For a 0.100 molal solution of the same solute, the changes are one-tenth as large. Water molecules tend to move from a region of higher water concentration to one of lower water con centration. When two different aqueous solutions are separated by a semipermeable membrane (one that al lows the passage of water but not solute molecules), water molecules diffusing from the region of higher water concentration to that of lower water concentration produce osmotic pressure (Fig. 2–12). This pressure, , measured as the force necessary to resist water move ment (Fig. 2–12c), is approximated by the van’t Hoff equation: II=icRT in which R is the gas constant and T is the absolute temperature. The term ic is the osmolarity of the solution, the product of the solute’s molar concentration c and the van’t Hoff factor i, which is a measure of the extent to which the solute dissociates into two or more ionic species. In dilute NaCl solutions, the solute completely

FIGURE 2–11 Solutes alter the colligative properties of aqueous solutions. (a) At 101 kPa (1 atm) pressure, pure water boils at 100 C and freezes at 0 C. (b) The presence of solute molecules reduces the probability of a water molecule leaving the solution and entering the gas phase, thereby reducing the vapor pressure of the solution and increasing the boiling point. Similarly, the probability of a water molecule colliding with and joining a forming ice crystal is reduced when some of the molecules colliding with the crystal are solute, not water, molecules. The effect is depression of the freezing point.

dissociates into Na⁺ and Cl-, doubling the number of solute particles, and thus i 2. For nonionizing solutes, i is always 1. For solutions of several (n) solutes, II is the sum of the contributions of each species:

II=RT (i₁c₁+i_2c2+ … +i_ncn)

Osmosis, water movement across a semipermeable membrane driven by differences in osmotic pressure, is an important factor in the life of most cells. Plasma membranes are more permeable to water than to most other small molecules, ions, and macromolecules. This permeability is due partly to simple diffusion of water through the lipid bilayer and partly to protein channels (aquaporins; see Fig. 11–XX) in the membrane that selectively permit the passage of water. Solutions of equal osmolarity are said to be isotonic. Surrounded by an isotonic solution, a cell neither gains nor loses water (Fig. 2–13). In a hypertonic solution, one with higher osmolarity than the cytosol, the cell shrinks as water flows out. In a hypotonic solution, with lower osmo larity than the cytosol, the cell swells as water enters. In their natural environments, cells generally contain higher concentrations of biomolecules and ions than their surroundings, so osmotic pressure tends to drive water into cells. If not somehow counterbalanced, this inward movement of water would distend the plasma membrane and eventually cause bursting of the cell (osmotic lysis). Several mechanisms have evolved to prevent this catastrophe. In bacteria and plants, the plasma mem brane is surrounded by a nonexpendable cell wall of sufficient rigidity and strength to resist osmotic pressure and prevent osmotic lysis. Certain freshwater protists that live in a highly hypotonic medium have an organelle (contractile vacuole) that pumps water out of the cell. In multicellular animals, blood plasma and interstitial fluid (the extracellular fluid of tissues) are maintained at an osmolarity close to that of the cytosol. The high concentration of albumin and other proteins in blood plasma contributes to its osmolarity. Cells also actively pump out ions such as Na+ into the interstitial fluid to stay in osmotic balance with their surroundings.

FIGURE 2–12 Osmosis and the measurement of osmotic pressure. (a) The initial state. The tube contains an aqueous solution, the beaker contains pure water, and the semipermeable membrane allows the passage of water but not solute. Water flows from the beaker into the tube to equalize its concentration across the membrane. (b) The final state. Water has moved into the solution of the nonpermeant com pound, diluting it and raising the column of water within the tube. At equilibrium, the force of gravity operating on the solution in the tube exactly balances the tendency of water to move into the tube, where its concentration is lower. (c) Osmotic pressure (II) is measured as the force that must be applied to return the solution in the tube to the level of that in the beaker. This force is proportional to the height, h, of the column in (b).

FIGURE 2–13 Effect of extracellular osmolarity on water movement across a plasma membrane. When a cell in osmotic balance with its surrounding medium (that is, in an isotonic medium) (a) is transferred into a hypertonic solution (b) or hypotonic solution (c), water moves across the plasma membrane in the direction that tends to equalize osmolarity outside and inside the cell.

Because the effect of solutes on osmolarity depends on the number of dissolved particles, not their mass, macromolecules (proteins, nucleic acids, polysaccha rides) have far less effect on the osmolarity of a solu tion than would an equal mass of their monomeric com ponents. For example, a gram of a polysaccharide composed of 1,000 glucose units has the same effect on osmolarity as a milligram of glucose. One effect of storing fuel as polysaccharides (starch or glycogen) rather than as glucose or other simple sugars is prevention of an enormous increase in osmotic pressure within the storage cell. Plants use osmotic pressure to achieve mechanical rigidity. The very high solute concentration in the plant cell vacuole draws water into the cell (Fig. 2–13). The resulting osmotic pressure against the cell wall (turgor pressure) stiffens the cell, the tissue, and the plant body. When the lettuce in your salad wilts, it is because loss of water has reduced turgor pressure. Sudden alterations in turgor pressure produce the movement of plant parts seen in touch-sensitive plants such as the Venus flytrap and mimosa (Box 2–1). Osmosis also has consequences for laboratory pro tocols. Mitochondria, chloroplasts, and lysosomes, for ex ample, are bounded by semipermeable membranes. In isolating these organelles from broken cells, biochemists must perform the fractionations in isotonic solutions (see Fig. 1–8). Buffers used in cellular fractionations commonly contain sufficient concentrations (about 0.2 M) of sucrose or some other inert solute to protect the organelles from osmotic lysis.

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مواضيع ذات صلة

Ionization of Water, Weak Acids, and Weak Bases: -The pH Scale Designates the H+ and OH- Concentrations

Ionization of Water, Weak Acids, and Weak Bases: -The Ionization of Water Is Expressed by an Equilibrium Constant

Ionization of Water, Weak Acids, and Weak Bases: - Pure Water Is Slightly Ionized

Weak Interactions in Aqueous Systems: -Weak Interactions Are Crucial to Macromolecular Structure and Function

Weak Interactions in Aqueous Systems: - van der Waals Interactions Are Weak Interatomic Attractions

Weak Interactions in Aqueous Systems: -Nonpolar Compounds Force Energetically Unfavorable Changes in the Structure of Water

Weak Interactions in Aqueous Systems: -Entropy Increases as Crystalline Substances Dissolve

Weak Interactions in Aqueous Systems: - Water Interacts Electrostatically with Charged Solutes

Weak Interactions in Aqueous Systems: -Water Forms Hydrogen Bonds with Polar Solutes

Weak Interactions in Aqueous Systems: - Hydrogen Bonding Gives Water Its Unusual Properties

Evolutionary Foundations: -Molecular Anatomy Reveals Evolutionary Relationships

Evolutionary Foundations: -Eukaryotic Cells Evolved from Prokaryotes in Several Stages