The Properties of Mixtures - Solutions and Colloids PDF

13 The Properties of Mixtures: Solutions and Colloids 13.1 Types of Solutions: 13.3 Why Substances Dissolve: Parts of Solute by...

13 The Properties of Mixtures: Solutions and Colloids 13.1 Types of Solutions: 13.3 Why Substances Dissolve: Parts of Solute by Parts of Solution Intermolecular Forces Breaking Down the Interconverting Concentration Terms and Solubility Solution Process 13.6 Colligative Properties of Intermolecular Forces in Solution Heat of Solution and Its Components Solutions Liquid Solutions and Molecular Heat of Hydration: Dissolving Ionic Nonvolatile Nonelectrolyte Solutions Polarity Solids in Water Using Colligative Properties to Find Gas and Solid Solutions Solution Process and Entropy Change Solute Molar Mass 13.2 Intermolecular Forces and 13.4 Solubility as an Equilibrium Volatile Nonelectrolyte Solutions Biological Macromolecules Process Strong Electrolyte Solutions Structures of Proteins Effect of Temperature on Solubility Applications of Colligative Properties Dual Polarity in Soaps, Membranes, Effect of Pressure on Solubility 13.7 Structure and Properties and Antibiotics 13.5 Concentration Terms of Colloids Structure of DNA Molarity and Molality Source: © amnat11/Shutterstock.com 532  Concepts and Skills to Review Before You Study This Chapter › separation of mixtures (Section 2.9) › intermolecular forces and polarizability (Section 12.3) › calculations involving mass percent (Section 3.1) and › monomers and polymers (Section 12.7) molarity (Section 4.1) › equilibrium nature of phase changes and vapor pressure of › electrolytes; water as solvent (Sections 4.1 and 12.5) liquids; phase diagrams (Section 12.2) › mole fraction and Dalton’s law (Section 5.4) V irtually all the gases, liquids, and solids in the real world are mixtures—two or more substances mixed together physically, not combined Table 13.1 Approximate Composition of a Bacterium Mass % Number Number of chemically. Synthetic mixtures usually contain only a dozen Substance of Cell of Types Molecules or so components; for example, the soda you may be drink- Water ∼70 1 5×1010 ing while reading your chemistry text contains water, car- Ions 1 20 ? bon dioxide, sugar, caffeine, caramel color, phosphoric and Sugars* 3 200 3×108 citric acids, and other flavorings. Natural mixtures, such as Amino acids* 0.4 100 5×107 seawater and soil, often contain over 50 components. Living Lipids* 2 50 3×107 mixtures, such as trees and students, are the most complex— Nucleotides* 0.4 200 1×107 even a simple bacterial cell contains nearly 6000 different Other small molecules 0.2 ∼200 ? compounds (Table 13.1). Macromolecules 23 ∼5000 6×106 Recall from Chapter 2 that a mixture has two defining (proteins, nucleic characteristics: its composition is variable, and it retains some acids, polysaccharides) properties of its components. We focus here on two common *Includes precursors and metabolites. types of mixtures—solutions and colloids—whose main dif- ferences relate to particle size and number of phases: ∙ A solution is a homogeneous mixture; that is, it exists as one phase. In a solution, the particles are individual atoms, ions, or small molecules. ∙ A colloid is a type of heterogeneous mixture. A heterogeneous mixture has two or more phases. They may be visibly distinct, like pebbles in concrete, or not, like the much smaller particles in the colloids smoke and milk. In a colloid, the par- ticles are typically macromolecules or aggregations of small molecules that are dispersed so finely they don’t settle out. IN THIS CHAPTER... We focus on how intermolecular forces and other energy consider- ations affect a solute dissolving in a solvent, how to calculate concentration, and how solu- tions differ from pure substances. We also briefly consider the behavior and applications of colloids. › We survey intermolecular forces between solute and solvent and find that substances with similar types of forces form a solution. › We see how the dual polarity of some organic molecules gives rise to these same intermo- lecular forces, and we find that they determine the structures and functions of soaps, anti- biotics, biological macromolecules, and cell membranes. › We use a stepwise cycle to see why a substance dissolves and examine the heat involved and the dispersal of matter that occurs when a solution forms. To understand the latter factor, we introduce the concept of entropy. › We examine the equilibrium nature of solubility and see how temperature and pressure affect it. › We define various solution concentration units and see how to interconvert them mathematically. › We see why the physical properties of solutions are different from those of pure substances and learn how to apply those differences. › We investigate colloids and apply solution and colloid chemistry to the purification of water. 533 534   Chapter 13 The Properties of Mixtures: Solutions and Colloids 13.1 TYPES OF SOLUTIONS: INTERMOLECULAR FORCES AND SOLUBILITY A solute dissolves in a solvent to form a solution. In general, the solvent is the most abundant component, but in some cases, the substances are miscible—soluble in each other in any proportion—so the terms “solute” and “solvent” lose their meaning. The physical state of the solvent usually determines the physical state of the solution. Solutions can be gaseous, liquid, or solid, but we focus mostly on liquid solutions because they are by far the most important. The solubility (S) of a solute is the maximum amount that dissolves in a fixed Ion-dipole (40–600) quantity of a given solvent at a given temperature, when an excess of the solute is present. Different solutes have different solubilities: ∙ Sodium chloride (NaCl), S = 39.12 g/100. mL water at 100.°C Na+ ∙ Silver chloride (AgCl), S = 0.0021 g/100. mL water at 100.°C Solubility is a quantitative term, but dilute and concentrated are qualitative, referring to the relative amounts of dissolved solute: ∙ The NaCl solution above is concentrated (a relatively large amount of solute dis- solved in a given quantity of solvent). H bond H2O ∙ The AgCl solution is dilute (a relatively small amount of solute in a given quantity (10–40) of solvent). Methanol (C H3OH) A given solute may dissolve in one solvent and not another. The explanation lies in the relative strengths of the intermolecular forces within both solute and solvent and between them. The useful rule-of-thumb “like dissolves like” says that substances Dipole -dipole (5–25) with similar types of intermolecular forces dissolve in each other. Thus, by knowing the forces, we can often predict whether a solute will dissolve in a solvent. Ethanal (C H3CHO) Chloroform Intermolecular Forces in Solution (CHCl3) All the intermolecular forces we discussed for pure substances also occur in solutions (Figure 13.1; also see Section 12.3): Ion–induced dipole (3–15) 1. Ion-dipole forces (attractions between ions and polar molecules) are the principal force involved when an ionic compound dissolves in water. Two events occur simultaneously: Cl– ∙ Forces compete. When a soluble salt is added to water, each type of ion attracts the oppositely charged pole of a water molecule. These attractions between ions Hexane (C6H14) and water compete with and overcome attractions between the ions, and the crystal structure breaks down. Dipole–induced dipole ∙ Hydration shells form. As an ion separates from the crystal structure, water (2–10) molecules cluster around it in hydration shells. The number of water molecules in the innermost shell depends on the ion’s size: four fit tetrahedrally around H2O Xenon small ions like Li+, while the larger Na+ and F− have six water molecules sur- rounding them octahedrally (Figure 13.2). In the innermost shell, normal hydro- Dispersion gen bonding between water molecules is disrupted to form the ion-dipole forces. (0.05–40) But these water molecules are H bonded to others in the next shell, and those are H bonded to others still farther away. 2. Hydrogen bonding (attractions between molecules with an H atom bonded to N, O, or F) is the principal force in solutions of polar, O- and N-containing organic and biological compounds, such as alcohols, amines, and amino acids. C6H14 3. Dipole-dipole forces (attractions between polar molecules), in the absence of Octane H bonding, allow polar molecules like propanal (CH3CH2CHO) to dissolve in polar (C8H18) solvents like dichloromethane (CH2Cl2). Figure 13.1 Types of intermolecular 4. Ion–induced dipole forces, one type of charge-induced dipole force, rely on polariz- forces in solutions. Forces are listed in ability. They arise when an ion’s charge distorts the electron cloud of a nearby non- order of decreasing strength (values are polar molecule, giving it a temporary dipole moment. This type of force initiates the in kJ/mol), with an example of each. binding of the Fe2+ ion in hemoglobin to an O2 molecule that enters a red blood cell.  13.1 Types of Solutions: Intermolecular Forces and Solubility    535 Ion-dipole Hydration Figure 13.2 Hydration shells around an forces shells Na+ ion. Ion-dipole forces orient water molecules around an ion. In the innermost shell here, six water molecules surround the cation octahedrally. + Hydrogen bonds 5. Dipole–induced dipole forces, also based on polarizability, arise when a polar mol- ecule distorts the electron cloud of a nonpolar molecule. They are weaker than ion– induced dipole forces because the charge of each pole is less than an ion’s (Coulomb’s law). The solubility in water of atmospheric O2, N2, and noble gases, while limited, is due in part to these forces. Paint thinners and grease solvents also rely on them. 6. Dispersion forces contribute to the solubility of all solutes in all solvents, but they are the principal intermolecular force in solutions of nonpolar substances, such as petroleum and gasoline. The same forces maintain the shapes of biological macromolecules (Section 13.2). Liquid Solutions and the Role of Molecular Polarity From cytoplasm to tree sap, gasoline to cleaning fluid, iced tea to urine, liquid solu- tions are very familiar. Water is the most prominent solvent, but there are many other liquid solvents, with polarities from very polar to nonpolar. Applying the Like-Dissolves-Like Rule The like-dissolves-like rule says that when the forces within the solute are similar to those within the solvent, the forces can replace each other and a solution forms. Thus, ∙ Salts are soluble in water because the ion-dipole attractions between ion and water are similar in strength to the strong attractions between the ions and the strong H bonds between water molecules, so they can replace each other. ∙ Salts are insoluble in hexane (C6H14) because the ion–induced dipole forces between ion and nonpolar hexane are very weak and cannot replace the strong attractions between the ions. ∙ Oil is insoluble in water because the weak dipole–induced dipole forces between oil and water molecules cannot replace the strong H bonds between water mole- cules or the extensive dispersion forces within the oil. ∙ Oil is soluble in hexane because dispersion forces in one can replace the similar dispersion forces in the other. Dual Polarity and Effects on Solubility To examine these ideas further, let’s compare the solubilities of a series of alcohols in water and in hexane (CH3CH2CH2CH2CH2CH3), two solvents with very different intermolecular forces; polar water molecules exhibit H bonds, while nonpolar hexane molecules have dispersion forces. Alcohols are organic compounds that have a dual polarity, a polar hydroxyl (OH) group bonded to a nonpolar hydrocarbon group: ∙ The OH portion interacts through strong H bonds with water and through weak dipole–induced dipole forces with hexane. ∙ The hydrocarbon portion interacts through dispersion forces with hexane and through very weak dipole–induced dipole forces with water. 536   Chapter 13 The Properties of Mixtures: Solutions and Colloids Table 13.2 Solubility* of a Series of Alcohols in Water and in Hexane Solubility Solubility in Alcohol Model in Water Hexane CH3OH ∞ 1.2 (methanol) CH3CH2OH ∞ ∞ (ethanol) CH3(CH2)2OH ∞ ∞ (1-propanol) CH3(CH2)3OH 1.1 ∞ (1-butanol) CH3(CH2)4OH 0.30 ∞ (1-pentanol) CH3(CH2)5OH 0.058 ∞ (1-hexanol) *Expressed in mol alcohol/1000 g solvent at 20°C. The general formula for an alcohol is CH3(CH2)nOH, and we’ll look at straight-chain In pure water, H bonds link the molecules. examples with one to six carbons (n = 0 to 5): 1. Solubility in water is high for smaller alcohols. From the models in Table 13.2, we see that the OH group is a relatively large portion of the alcohols with one to three carbons (n = 0 to 2). These molecules interact with each other through H bond- ing, just as water molecules do. When they mix with water, H bonding within solute and within solvent is replaced by H bonding between solute and solvent (Figure 13.3). As a result, these smaller alcohols are miscible with water. 2. Solubility in water is low for larger alcohols. Solubility decreases dramatically for alcohols larger than three carbons (n > 2); in fact, those with chains longer than six carbons (n > 5) are insoluble in water. For larger alcohols to dissolve, the nonpo- In pure methanol, H bonds lar chains have to move among the water molecules, replacing the strong H-bond link the molecules. attractions between water molecules with the weak attractions of the chains for water. While the OH portion of such an alcohol forms H bonds to water, these cannot make up for all the other H bonds between water molecules that have to break to make room for the long hydrocarbon portion. Table 13.2 shows that the opposite trend occurs with hexane: 1. Solubility in hexane is low for the smallest alcohol. For alcohols in hexane, in addition to dispersion forces, weak dipole–induced dipole forces exist between the OH of methanol (CH3OH) and hexane. These cannot replace the strong H bonding between CH3OH molecules, so solubility is relatively low. In a solution of water and 2. Solubility in hexane is high for larger alcohols. In any larger alcohol (n > 0), methanol, H bonds link dispersion forces between the hydrocarbon portion and hexane can replace dispersion the two types of molecules. forces between hexane molecules. With only weak forces within the solvent to be Figure 13.3 Like dissolves like: solubility replaced, even ethanol, with a two-carbon chain, has enough dispersion forces between of methanol in water. it and hexane to be miscible.  13.1 Types of Solutions: Intermolecular Forces and Solubility    537 Many organic molecules have polar and nonpolar portions, which determine their solubility. For example, carboxylic acids and amines behave like alcohols: methanoic acid (HCOOH, formic acid) and methanamine (CH3NH2) are miscible with water and slightly soluble in hexane, whereas hexanoic acid [CH3(CH2)4COOH] and 1-hexanamine [CH3(CH2)5NH2] are slightly soluble in water and very soluble in hexane. SAMPLE PROBLEM 13.1 Predicting Relative Solubilities Problem Predict which solvent will dissolve more of the given solute: (a) Sodium chloride in methanol (CH3OH) or in 1-propanol (CH3CH2CH2OH) (b) Ethylene glycol (HOCH2CH2OH) in hexane (CH3CH2CH2CH2CH2CH3) or in water (c) Diethyl ether (CH3CH2OCH2CH3) in water or in ethanol (CH3CH2OH) Plan We examine the formulas of solute and solvent to determine the forces in and between solute and solvent. A solute is more soluble in a solvent whose intermolecular forces are similar to, and therefore can replace, its own. Solution (a) Methanol. NaCl is ionic, so it dissolves through ion-dipole forces. Both methanol and 1-propanol have a polar OH group, but the hydrocarbon portion of each alcohol interacts only weakly with the ions and 1-propanol has a longer hydrocarbon portion than methanol. (b) Water. Ethylene glycol molecules have two OH groups, so they interact with each other through H bonding. H bonds formed with H2O can replace these H bonds between solute molecules better than dipole–induced dipole forces with hexane can. (c) Ethanol. Diethyl ether molecules interact through dipole-dipole and dispersion forces. They can form H bonds to H2O or to ethanol. But ethanol can also interact with the ether effectively through dispersion forces because it has a hydrocarbon chain. FOLLOW-UP PROBLEMS Brief Solutions for all Follow-up Problems appear at the end of the chapter. 13.1A State which solute is more soluble in the given solvent and which forces are most important: (a) 1-butanol (CH3CH2CH2CH2OH) or 1,4-butanediol (HOCH2CH2CH2CH2OH) in water; (b) chloroform (CHCl3) or carbon tetrachloride (CCl4) in water. 13.1B State which solvent can dissolve more of the given solute and which forces are most important: (a) chloromethane (CH3Cl) in chloroform (CHCl3) or in methanol (CH3OH); (b) pentanol (C5H11OH) in water or in hexane (C6H14). SOME SIMILAR PROBLEMS 13.13 and 13.14 Gas-Liquid Solutions A substance with very weak intermolecular attractions should Correlation Between have a low boiling point and, thus, would be a gas under ordinary conditions. Also, Boiling Point and it would not be very soluble in water because of weak solute-solvent forces. Thus, Table 13.3  Solubility in Water for nonpolar or slightly polar gases, boiling point generally correlates with solubility in water (Table 13.3). A higher boiling point is an indication of stronger intermolec- Gas Solubility (mol/L )* bp (K) ular forces, which result in a greater solubility in water. −4 He 4.2×10   4.2 The small amount of a nonpolar gas that does dissolve may be vital. At 25°C and 1 atm, the solubility of O2 is only 3.2 mL/100. mL of water, but aquatic animal life Ne 6.6×10−4 27.1 requires it. At times, the solubility of a nonpolar gas may seem high because it is also N2 10.4×10−4 77.4 reacting with the solvent. Oxygen seems more soluble in blood than in water because CO 15.6×10−4 81.6 it bonds chemically to hemoglobin in red blood cells. Carbon dioxide, which is essen- O2 21.8×10−4 90.2 tial for aquatic plants and coral-reef growth, seems very soluble in water (∼81 mL of NO 32.7×10−4 121.4 CO2 /100. mL of H2O at 25°C and 1 atm) because it is dissolving and reacting: *At 273 K and 1 atm. CO2(g) + H2O(l) ⥫⥬ H+(aq) + HCO −3 (aq) Gas Solutions and Solid Solutions Gas solutions and solid solutions also have vital importance and numerous applications. Gas-Gas Solutions All gases are miscible with each other. Air is the classic exam- ple of a gaseous solution, consisting of about 18 gases in widely differing proportions. 538   Chapter 13 The Properties of Mixtures: Solutions and Colloids Zinc Carbon Copper Iron A Brass, a substitutional alloy B Carbon steel, an interstitial alloy Figure 13.4 The arrangement of atoms in two types of alloys. Anesthetic gas proportions are finely adjusted to the needs of the patient and the Source: (A) © Ruth Melnick; (B) © Ingram length of the surgical procedure. The proportions of many industrial gas mixtures, Publishing/age fotostock RF such as CO/H2 in syngas production or N2/H2 in ammonia production, are controlled to optimize product yield under varying conditions of temperature and pressure. Gas-Solid Solutions When a gas dissolves in a solid, it occupies the spaces between the closely packed particles. Hydrogen gas can be purified by passing an impure sample through palladium. Only H2 molecules are small enough to fit between the Pd atoms, where they form PdH bonds. The H atoms move from one Pd atom to another and emerge from the metal as H2 molecules (see Figure 14.2). The ability of gases to dissolve in a solid also has disadvantages. The electrical conductivity of copper is greatly reduced by the presence of O2, which dissolves into the crystal structure and reacts to form copper(I) oxide. High-conductivity copper is prepared by melting and recasting the metal in an O2-free atmosphere. Solid-Solid Solutions Solids diffuse so little that their mixtures are usually hetero- geneous. Some solid-solid solutions can be formed by melting the solids and then mixing them and allowing them to freeze. Many alloys, mixtures of elements that have a metallic character, are solid-solid solutions (although several have microscopic heterogeneous regions). Alloys generally fall into one of two categories: ∙ In a substitutional alloy like brass (Figure 13.4A), atoms of zinc replace atoms of the main element, copper, at some sites in the cubic closest packed array. This occurs when the atoms of the elements in the alloy are similar in size. ∙ In an interstitial alloy like carbon steel (Figure 13.4B), atoms of carbon (a nonmetal is typical in this type of alloy) fill some spaces (interstices) between atoms of the main element, iron, in the body-centered array. Waxes are also solid-solid solutions. Most are amorphous solids with some small regions of crystalline regularity. A wax is defined as a solid of biological origin that is insoluble in water and soluble in nonpolar solvents. Beeswax, which bees secrete to build their combs, is a homogeneous mixture of fatty acids, long-chain carboxylic acids, and hydrocarbons in which some of the molecules are more than 40 carbon atoms long. Carnauba wax, from a South American palm, is a mixture of compounds, each consisting of a fatty acid bound to a long-chain alcohol. It is hard but forms a thick gel in nonpolar solvents, so it is perfect for waxing cars. › Summary of Section 13.1 › A solution is a homogeneous mixture of a solute dissolved in a solvent through the action of intermolecular forces. › Ion-dipole, ion–induced dipole, and dipole–induced dipole forces occur in solutions, in addition to all the intermolecular forces that also occur in pure substances. › If similar intermolecular forces occur in solute and solvent, they replace each other when the substances mix and a solution is likely to form (the like-dissolves-like rule). › When ionic compounds dissolve in water, the ions become surrounded by hydration shells of H-bonded water molecules.  13.2 Intermolecular Forces and Biological Macromolecules    539 › Solubility of organic molecules in various solvents depends on the relative sizes of their polar and nonpolar portions. › The solubility of nonpolar gases in water is low because of weak intermolecular forces. Gases are miscible with one another and dissolve in solids by fitting into spaces in the crystal structure. › Solid-solid solutions include alloys (some of which are formed by mixing molten components) and waxes. 13.2 INTERMOLECULAR FORCES AND BIOLOGICAL MACROMOLECULES We discuss the shapes of proteins, nucleic acids, and cell membranes, as well as the functions of soaps and antibiotics, in this chapter on solutions because they depend on intermolecular forces too. These shapes and functions are explained by two ideas: ∙ Polar and ionic groups attract water, but nonpolar groups do not. ∙ Just as separate molecules attract each other, so do distant groups on the same molecule. The Structures of Proteins Proteins are very large molecules (called polymers) formed by linking together many smaller molecules called amino acids; about 20 different amino acids occur in pro- teins, which range in size from about 50 amino acids (ℳ ≈ 5×103 g/mol) to several thousand (ℳ ≈ 5×105 g/mol). Proteins with a few types of amino acids in repeating patterns have extended helical or sheetlike shapes and give structure to hair, skin, and so forth. Proteins with many types of amino acids have complex, globular shapes and function as antibodies, enzymes, and so forth. Let’s see how intermolecular forces among amino acids influence the shapes of globular proteins. The Polarity of Amino-Acid Side Chains In a cell, a free amino acid has four groups bonded to one C, which is called the α-carbon (Figure 13.5): charged carboxyl (COO−) one of 20 and amine (NH+3 ) groups, an H atom, and a side chain, represented by an R, which different ranges from another H atom, to a few C atoms, to a two-ringed C9H8N group. In a R side chains protein, the carboxyl group of one amino acid is linked covalently to the amine group + O of the next by a peptide bond. (We discuss amino acid structures and peptide bond H3N C C formation in Chapter 15.) Thus, as Figure 13.6 shows, the backbone of a protein is a α-carbon O− H polypeptide chain: an α-carbon connected through a peptide bond (orange screen) to the next α-carbon, and so forth. The various side chains (gray screens) dangle off the Figure 13.5 The charged form of an α-carbons on alternate sides of the chain. amino acid under physiological conditions. Glutamic acid Carbon Nitrogen Oxygen Hydrogen – Serine Glycine Lysine α α α α Peptide bond Figure 13.6 A portion of a polypeptide chain. Three peptide bonds (orange + screens) join four amino acids in this chain portion. 540   Chapter 13 The Properties of Mixtures: Solutions and Colloids Amino acids can be classified by the polarity or charge of their side chains: nonpolar, polar, and ionic. A few examples are + NH3 NONPOLAR POLAR IONIC CH2 CH3 CH2 COO – CH3 CH OH SH CH2 CH2 CH2 CH2 CH2 CH2 CH2 + + + + + H3N C COO – H3N C COO – H3N C COO – H3N C COO – H3N C COO – H H H H H Leucine Serine Cysteine Lysine Glutamic acid Intermolecular Forces and Protein Shape The same forces that act between sep- arate molecules are responsible for a protein’s shape, because distant groups on the protein chain end up near each other as the chain bends. Figure 13.7 depicts the forces within a small portion of a protein and between the protein and the aqueous medium of the cell. In general order of importance, they are ∙ Covalent peptide bonds create the backbone (polypeptide chain). ∙ Helical and sheetlike segments arise from H bonds between the CO of one pep- tide bond and the N—H of another. ∙ Polar and ionic side chains protrude into the surrounding cell fluid, interacting with water through ion-dipole forces and H bonds. ∙ Nonpolar side chains interact through dispersion forces within the nonaqueous protein interior. Figure 13.7 The forces that maintain protein structure. Covalent, ionic, and in- termolecular forces act between the parts H bonds within of a portion of a protein and between the polypeptide chain protein and surrounding H2O molecules to determine the protein’s shape. (Water molecules and some amino-acid side chains are shown as ball-and-stick models within space-filling contours.) + Region where dispersion forces predominate Ion-dipole forces – – + H bond Salt link H bonds within polypeptide chain H bond Disulfide Disulfide bond bridge  13.2 Intermolecular Forces and Biological Macromolecules    541 ∙ The SH ends of two cysteine side chains form a covalent SS bond, a disulfide bridge, between distant parts of the chain that creates a loop. ∙ Oppositely charged ends of ionic side chains, COO− and NH 3+ groups, that lie near each other form an electrostatic salt link (or ion pair) that creates bend in the protein chain. ∙ Other H bonds between side chains keep distant chain portions near each other. Thus, soluble proteins have polar-ionic exteriors and nonpolar interiors. As we empha- size in Chapter 15, the amino acid sequence of a protein determines its shape, which determines its function. Dual Polarity in Soaps, Membranes, and Antibiotics Dual polarity, which we noted as a key factor in the solubility of alcohols, also helps explain how soaps, cell membranes, and antibiotics function. Action of Soaps A soap is the salt formed when a strong base (a metal hydroxide) reacts with a fatty acid, a carboxylic acid with a long hydrocarbon chain. A typical soap molecule is made up of a nonpolar “tail” 15–19 carbons long and a polar-ionic “head” consisting of a COO− group and the cation of the strong base. The cation greatly influences a soap’s properties. Lithium soaps are hard and high melting and used in car lubricants. Potassium soaps are low melting and used in liquid form. Several sodium soaps, including sodium stearate, CH3(CH2)16COONa, are components of common bar soaps: – Na+ Sodium stearate, C17H35COONa When grease on your hands or clothes is immersed in soapy water, the soap mol- ecules’ nonpolar tails interact with the nonpolar grease molecules through dispersion forces, while the polar-ionic heads attract water molecules through ion-dipole forces and H bonds. Tiny aggregates of grease molecules, embedded with soap molecules whose polar-ionic heads stick into the water, are flushed away by added water (Figure 13.8). Figure 13.8 The cleaning ability of a soap depends on the dual polarity of its molecules. Polar heads of soap molecules interact with H2O molecules. Nonpolar tails of soap molecules interact with hydrocarbons in a grease stain. Nonpolar molecules make up a grease stain. 542   Chapter 13 The Properties of Mixtures: Solutions and Colloids + Lipid Bilayers and the Structure of the Cell Membrane The most abundant mol- Polar-ionic ecules in cell membranes are phospholipids. Like soaps, they have a dual polarity—a head nonpolar tail consists of two fatty acid chains, and an organophosphate group is the – polar-ionic head (Figure 13.9). Remarkably, phospholipids self-assemble in water into a sheetlike double layer called a lipid bilayer, with the tails of the two layers touching and the heads in the water. In the laboratory, bilayers form spherical vesicles that trap water inside. These structures are favored energetically because of their intermolecular forces: ∙ Ion-dipole forces occur between polar heads and water inside and outside. Nonpolar ∙ Dispersion forces occur between nonpolar tails within the bilayer interior. tail ∙ Minimal contact exists between nonpolar tails and water. A typical animal cell membrane consists of a phospholipid bilayer with proteins partially embedded in it; a small portion of such a membrane appears in Figure 13.10. Membrane proteins, which play countless essential roles, differ fundamentally from soluble proteins in terms of their dual polarity: ∙ Soluble proteins have polar exteriors and nonpolar interiors. They form ion-dipole Figure 13.9 A membrane phospholipid. and H-bonding forces between water and polar groups on the exterior and disper- Lecithin (phosphatidylcholine), a phospho- sion forces between nonpolar groups in the interior (see Figure 13.7). lipid, is shown as a space-filling model and ∙ Membrane proteins have exteriors that are partially polar (red) and partially nonpo- as a simplified purple-and-gray shape. lar (blue). They have polar groups on the exterior portion that juts into the aqueous surroundings and nonpolar groups on the exterior portion embedded in the mem- brane. These nonpolar groups form dispersion forces with the phospholipid tails of the bilayer. Channel proteins also have polar groups lining the aqueous channel. Polar-ionic heads of Other proteins interact phospholipids interact with Polar regions of membrane with membrane proteins through water through ion-dipole proteins interact with water ion-dipole forces and H bonds. forces and H bonds. through ion-dipole forces and H bonds. AQUEOUS CELL EXTERIOR CUTAWAY OF CHANNEL PROTEIN Nonpolar region MEMBRANE PROTEIN Nonpolar regions of membrane Polar proteins interact with nonpolar region Nonpolar tails of phospholipids tails through dispersion forces. AQUEOUS CELL interact through dispersion forces. INTERIOR Figure 13.10 Intermolecular forces and cell membrane structure. Action of Antibiotics A key function of a cell membrane is to balance internal and external ion concentrations: Na+ is excluded from the cell, and K+ is kept inside. Gramicidin A and similar antibiotics act by forming channels in the cell membrane of a bacterium through which ions flow (Figure 13.11). Two helical gramicidin A mole- cules, their nonpolar groups outside and polar groups inside, lie end to end to form a channel through the membrane. The nonpolar outside stabilizes the molecule in the cell membrane through dispersion forces, and the polar inside passes the ions along using ion-dipole forces, like a “bucket brigade.” Over 107 ions pass through each of these channels per second, which disrupts the vital ion balance, and the bacterium dies. The Structure of DNA The chemical information that guides the design and construction, and therefore the function, of all proteins is contained in nucleic acids, unbranched polymers made up  13.2 Intermolecular Forces and Biological Macromolecules    543 K + flows out of the cell; Na+ flows in. Figure 13.11 The mode of action of the antibiotic gramicidin A. The K+ ions CELL Polar interior of EXTERIOR K+ gramicidin A are shown leaving the cell. At the same Na+ time (not shown), Na+ ions enter the cell. Na+ K+ Polar outer surface of membrane Nonpolar interior of membrane Polar inner surface of membrane K+ K+ Bacterial CELL K+ Nonpolar membrane INTERIOR exterior of K+ gramicidin A – of smaller units (monomers) called mononucleotides. Each mononucleotide consists of an N-containing base, a sugar, and a phosphate group (Figure 13.12). In DNA (deoxyribonucleic acid), the sugar is 2-deoxyribose, in which H substitutes for OH on the second C atom of the five-C sugar ribose. The repeating pattern of the DNA chain is sugar linked to phosphate linked to sugar linked to phosphate, and so on. Attached to each sugar is one of four N-containing Phosphate se group ba bases, flat ring structures that dangle off the polynucleotide chain, similar to the way amino-acid side chains dangle off the polypeptide chain. – Intermolecular Forces and the Double Helix DNA exists as two chains wrapped around each other in a double helix that is stabilized by intermolecular forces One of four ­(Figure 13.13): possible bases Sugar ∙ On the more polar exterior, negatively charged sugar-phosphate groups interact with (2-deoxyribose) the aqueous surroundings via ion-dipole forces and H bonds. se ∙ In the less polar interior, flat, N-containing bases stack above each other and interact ba 2 by dispersion forces. ∙ Bases form specific interchain H bonds; that is, each base in one chain is always H bonded with its complementary base in the other chain. Thus, the base sequence Figure 13.12 A short portion of the of one chain is the H-bonded complement of the base sequence of the other. polynucleotide chain of DNA. H bonds – between bases Cytosine (C) – Guanine (G) Polar-ionic exterior Ion-dipole – – force – – – 2-Deoxyribose 2-Deoxyribose H bond – G always forms an – H-bonded pair with C. Nonpolar interior – – Carbon Nitrogen Oxygen Hydrogen Figure 13.13 The double helix of DNA. A segment of DNA (left) has its polar-ionic sugar-phosphate portion (pink) facing the water and the non- polar bases (gray) stacking in the interior. The expanded portion (right) shows an H-bonded pair of the bases guanine and cytosine. 544   Chapter 13 The Properties of Mixtures: Solutions and Colloids A DNA molecule contains millions of H bonds linking bases in these prescribed pairs. The total energy of the H bonds keeps the chains together, but each H bond is weak enough (around 5% of a typical covalent single bond) that a few at a time can break as the chains separate during crucial cellular processes. (In Chapter 15, we’ll see how H-bonded base pairs are essential for protein synthesis and DNA replication.) › Summary of Section 13.2 › In soluble proteins, polar and ionic amino-acid side chains on the exterior interact with surrounding water, and nonpolar side chains in the interior interact with each other. › With polar-ionic heads and nonpolar tails, a soap dissolves grease and interacts with water. › Like soaps, phospholipids have dual polarity. They assemble into a water-impermeable lipid bilayer. In a cell membrane, the embedded portions of membrane proteins have exterior nonpolar side chains that interact with the nonpolar tails in the lipid bilayer through dispersion forces. Some antibiotics form channels with nonpolar exteriors and polar interiors that shuttle ions through the cell membrane. › DNA forms a double helix with a sugar-phosphate, polar-ionic exterior. In the interior, N-containing bases H bond in specific pairs and stack through dispersion forces. 13.3 WHY SUBSTANCES DISSOLVE: BREAKING DOWN THE SOLUTION PROCESS The qualitative macroscopic rule “like dissolves like” is based on molecular interactions between the solute and the solvent. To see why like dissolves like, we’ll break down the solution process conceptually into steps and examine each of them quantitatively. The Heat of Solution and Its Components Before a solution forms, solute particles (ions or molecules) are attracting each other, as are solvent particles (molecules). For one to dissolve in the other, three steps must take place, each accompanied by an enthalpy change: Step 1. S olute particles separate from each other. This step involves overcoming inter- molecular (or ionic) attractions, so it is endothermic: Solute (aggregated) + heat ⟶ solute (separated)   ΔHsolute > 0 Step 2. S olvent particles separate from each other. This step also involves overcoming attractions, so it is endothermic, too: Solvent (aggregated) + heat ⟶ solvent (separated)   ΔHsolvent > 0 Step 3. S olute and solvent particles mix and form a solution. The different particles attract each other and come together, so this step is exothermic: Solute (separated) + solvent (separated) ⟶ solution + heat   ΔHmix < 0 The overall process is called a thermochemical solution cycle, and in yet another application of Hess’s law, we combine the three individual enthalpy changes to find the heat (or enthalpy) of solution (ΔHsoln), the total enthalpy change that occurs when solute and solvent form a solution: ΔHsoln = ΔHsolute + ΔHsolvent + ΔHmix (13.1) Overall solution formation is exothermic or endothermic, and ΔHsoln is either negative or positive, depending on the relative sizes of the individual ΔH values: ∙ Exothermic process: ΔHsoln < 0. If the sum of the endothermic terms (ΔHsolute + Student Hot Spot ΔHsolvent) is smaller than the exothermic term (ΔHmix), the process is exothermic and ΔHsoln is negative (Figure 13.14A). Student data indicate that you may struggle with the concept of energy changes and the solution ∙ Endothermic process: ΔHsoln > 0. If the sum of the endothermic terms is larger than process. Access the Smartbook to view additional the exothermic term, the process is endothermic and ΔHsoln is positive (Figure 13.14B). Learning Resources on this topic. If ΔHsoln is highly positive, the solute may not dissolve significantly in that solvent.  13.3 Why Substances Dissolve: Breaking Down the Solution Process    545 Solvent separated Solute Solute Solvent ΔHmix separated separated separated Enthalpy, H ΔHsolute Solution + ΔHsolvent Enthalpy, H Solute ΔHsolute Solvent ΔHsolvent + Hfinal Solvent aggregated ΔHsolvent aggregated Solute ΔHsolute ΔHsolute ΔHsolvent aggregated aggregated ΔH mix ΔHsoln > 0 H initial Solution H initial ΔHsoln < 0 H final A Exothermic solution process B Endothermic solution process Figure 13.14 Enthalpy components of the heat of solution. A, ΔHmix is larger than the sum of ΔHsolute and ΔHsolvent, so ΔHsoln is negative. B, ΔHmix is smaller than the sum of ΔHsolute and ΔHsolvent, so ΔHsoln is positive. The Heat of Hydration: Dissolving Ionic Solids in Water The ΔHsolvent and ΔHmix components of the solution cycle are difficult to measure individually. Combined, they equal the enthalpy change for solvation, the process of surrounding a solute particle with solvent particles: ΔHsolvation = ΔHsolvent + ΔHmix Solvation in water is called hydration. Thus, enthalpy changes for separating the water molecules (ΔHsolvent) and mixing the separated solute with them (ΔHmix) are combined into the heat (or enthalpy) of hydration (ΔHhydr). In water, Equation 13.1 becomes ΔHsoln = ΔHsolute + ΔHhydr The heat of hydration is a key factor in dissolving an ionic solid. Breaking the H bonds in water is more than compensated for by forming the stronger ion-dipole forces, so hydration of an ion is always exothermic. The ΔHhydr of an ion is defined as the enthalpy change for the hydration of 1 mol of separated (gaseous) ions: HO M+(g) [or X−(g)] ---⟶ 2 M+(aq) [or X−(aq)]  ΔHhydr of the ion (always 0) = ΔHlattice F− 133 −431 Thus, for ionic compounds in water, the heat of solution is the lattice energy (always Cl− 181 −313 positive) plus the combined heats of hydration of the ions (always negative): Br− 196 −284 I− 220 −247 ΔHsoln = ΔHlattice + ΔHhydr of the ions (13.2) 546   Chapter 13 The Properties of Mixtures: Solutions and Colloids Na + (g) NH4+ (g) Cl –(g) NO–3 (g) ΔHhydr ΔH solute ΔHhydr NH4+ (aq) Enthalp y, H Enthalp y, H (ΔH lattice) NO –3(aq) Hfinal Na + (g) OH –(g) ΔH solute (ΔH lattice) ΔHsoln ΔHsoln Na + (aq) = 3.9 kJ/mol ΔH solute = 25.7 Cl –(aq) Hfinal (ΔH lattice) kJ/mol NaCl( s) NaOH( s) NH4NO3(s) Enthalp y, H Hinitial Hinitial Hinitial A NaCl. ΔHlattice is slightly larger than ΔHhydr : ΔHhydr C NH 4 NO3. ΔHlattice dominates: ΔHsoln is ΔHsoln is small and positive. ΔHsoln large and positive. = – 44.5 kJ/mol Na + (aq) OH –(aq) Hfinal B NaOH. ΔHhydr dominates: ΔHsoln is large and negative. Figure 13.15 Enthalpy diagrams for three ionic compounds dissolving in water. Once again, the sizes of the individual terms determine the sign of ΔHsoln. Figure 13.15 shows qualitative enthalpy diagrams for three ionic solutes dissolv- ing in water: ∙ NaCl. Sodium chloride has a small positive ΔHsoln (3.9 kJ/mol) because its lattice energy is only slightly greater than the combined ionic heats of hydration: if you dissolve NaCl in water in a flask, you don’t feel the small temperature change. ∙ NaOH. Sodium hydroxide has a large negative ΔHsoln (–44.5 kJ/mol) because its lattice energy is much smaller than the combined ionic heats of hydration: if you dissolve NaOH in water, the flask feels hot. ∙ NH4NO3. Ammonium nitrate has a large positive ΔHsoln (25.7 kJ/mol) because its lattice energy is much larger than the combined ionic heats of hydration: if you dissolve NH4NO3 in water, the flask feels cold. Hot and cold “packs” consist of a thick outer pouch of water and a thin inner pouch of a salt. A squeeze breaks the inner pouch, and the salt dissolves. Most hot packs use anhydrous CaCl2 (ΔHsoln = −82.8 kJ/mol). In Japan, some soup is sold in double- walled cans with a salt in a packet immersed in water between the walls. Open the can and the packet breaks, the salt dissolves, and the soup quickly warms to about 90°C. Cold packs use NH4NO3 (ΔHsoln = 25.7 kJ/mol). A cold pack can keep the solution at 0°C for about half an hour, long enough to soothe a sprain. SAMPLE PROBLEM 13.2 Calculating an Aqueous Ionic Heat of Solution Problem With secondary applications ranging from sedative to fire retardant, calcium bromide is used primarily in concentrated solution as an industrial drilling fluid. (a) Use Table 13.4 and the lattice energy (2132 kJ/mol) to find the heat of solution (kJ/mol) of calcium bromide. (b) Draw an enthalpy diagram for this solution process. (a) Calculating the heat of solution of CaBr2. Plan We are given, or can look up, the individual enthalpy components for a salt dissolving in water and have to determine their signs to calculate the overall heat of solution (ΔHsoln). The components are the lattice energy (the heat absorbed when the solid separates into gaseous ions) and the heat of hydration for each ion (the heat  13.3 Why Substances Dissolve: Breaking Down the Solution Process    547 released when the ion becomes hydrated). The lattice energy is always positive, so ΔHlattice = 2132 kJ/mol. Heats of hydration are always negative, so from Table 13.4, ΔHhydr of Ca2+ = −1591 kJ/mol and ΔHhydr of Br− = −284 kJ/mol. We use Equation 13.2, noting that there are 2 mol of Br−, to obtain ΔHsoln. Solution ΔHsoln = ΔHlattice + ΔHhydr of the ions = ΔHlattice + ΔHhydr of Ca2+ + 2(ΔHhydr of Br− ) = 2132 kJ/mol + (−1591 kJ/mol) + 2(−284 kJ/mol) = −27 kJ/mol Check Rounding to check the math gives 2100 kJ/mol – 1600 kJ/mol – 560 kJ/mol = –60 kJ/mol This small negative sum indicates that our answer is correct. (b) Drawing an enthalpy diagram for the process of dissolving CaBr2. Plan Along a vertical enthalpy axis, the lattice energy (endothermic) is represented by Ca 2+ (g) an upward arrow leading from solid salt to gaseous ions. The hydration of the ions 2Br – (g) (exothermic) is represented by a downward arrow from gaseous to hydrated ions. From part (a), ΔHsoln is small and negative, so the downward arrowhead is slightly below the tail of the upward arrow. ΔHsolute (ΔHlattice) Enthalpy, H Solution The enthalpy diagram is shown in the margin. FOLLOW-UP PROBLEMS 13.2A Use the following data to find the combined heat of hydration for the ions in CaBr2(s) ΔHhydr Hinitial KNO3: ΔHsoln = 34.89 kJ/mol and ΔHlattice = 685 kJ/mol. ΔHsoln = 13.2B Use the following data to find the heat of hydration of CN−: ΔHsoln of NaCN = Ca 2+ (aq) −27 kJ/mol 1.21 kJ/mol, ΔHlattice of NaCN = 766 kJ/mol, and ΔHhydr of Na+ = −410. kJ/mol. 2Br – (aq) Hfinal SOME SIMILAR PROBLEMS 13.30, 13.31, 13.36, and 13.37 The Solution Process and the Change in Entropy The heat of solution (ΔHsoln) is one of two factors that determine whether a solute dissolves. The other factor is the natural tendency of a system of particles to spread out, which results in the system’s kinetic energy becoming more dispersed or more widely distributed. A thermodynamic variable called entropy (S) is directly related to the number of ways a system can distribute its energy, which involves the freedom of motion of the particles. Let’s see what it means for a system to “distribute its energy.” We’ll first compare the three physical states and then compare solute and solvent with solution. Entropy and the Three Physical States The states of matter differ significantly in their entropy. ∙ In a solid, the particles are fixed in their positions with little freedom of motion. In a liquid, they can move around each other and so have greater freedom of motion. And in a gas, the particles have little restriction and much more freedom of motion. ∙ The more freedom of motion the particles have, the more ways they can distribute their kinetic energy; thus, a liquid has higher entropy than a solid, and a gas has higher entropy than a liquid: Sgas > Sliquid > Ssolid ∙ Thus, there is a change in entropy (ΔS) associated with a phase change, and it can be positive or negative. For example, when a liquid vaporizes, the change in entropy (ΔSvap = Sgas − Sliquid) is positive (ΔSvap > 0, increase in entropy) since Sgas > Sliquid; when a liquid freezes (fusion), the change in entropy (ΔSfus = Ssolid − Sliquid) is negative (ΔSfus < 0, decrease in entropy) since Ssolid < Sliquid. 548   Chapter 13 The Properties of Mixtures: Solutions and Colloids Entropy and the Formation of Solutions The formation of solutions also involves a change in entropy. A solution usually has higher entropy than the pure solute and pure solvent because the number of ways to distribute the energy is related to the number of interactions between different molecules. There are far more interactions possible when solute and solvent are mixed than when they are pure; thus, Ssoln > (Ssolute + Ssolvent)   or   ΔSsoln > 0 You know from everyday experience that solutions form naturally, but pure solutes and solvents don’t: you’ve seen sugar dissolve in water, but you’ve never seen a sugar solution separate into pure sugar and water. In Chapter 20, we’ll see that energy is needed to reverse the natural tendency of systems to distribute their energy—to get “mixed up.” Water treatment plants, oil refineries, steel mills, and many other industrial facilities expend a lot of energy to separate mixtures into pure components. Enthalpy vs. Entropy Changes in Solution Formation Solution formation involves the interplay of two factors: systems change toward a state of lower enthalpy and higher entropy, so the relative sizes of ΔHsoln and ΔSsoln determine whether a solu- tion forms. Let’s consider three solute-solvent pairs to see which factor dominates in each case: 1. NaCl in hexane. Given their very different intermolecular forces, we predict that sodium chloride does not dissolve in hexane (C6H14). An enthalpy diagram (Figure 13.16A) shows that separating the nonpolar solvent is easy because the dispersion forces are weak (ΔHsolvent ≥ 0), but separating the solute requires sup- plying the very large ΔHlattice (ΔHsolute ≫ 0). Mixing releases little heat because ion–induced dipole forces between Na+ (or Cl−) and hexane are weak (ΔHmix ≤ 0). Because the sum of the endothermic terms is much larger than the exothermic term, ΔHsoln ≫ 0. A solution does not form because the entropy increase from mixing solute and solvent would be much smaller than the enthalpy increase required to separate the solute: ΔSmix ≪ ΔHsolute. 2. Octane in hexane. We predict that octane (C8H18) is soluble in hexane because both are held together by dispersion forces of similar strength; in fact, these two substances are miscible. That is, both ΔHsolute and ΔHsolvent are around zero. The similar forces mean that ΔHmix is also around zero. And a lot of heat is not released; in fact, ΔHsoln is around zero (Figure 13.16B). So why does a solution form so readily? With no enthalpy change driving the process, octane dissolves in hexane because the entropy increases greatly when the pure substances mix: ΔSmix ≫ ΔHsoln. ΔHmix Figure 13.16 Enthalpy diagrams for dissolving (A) NaCl and Solution (B) octane in hexane. Na + (g)+Cl –(g) H final separated ΔHsolute + ΔHsoln >> 0 ΔH solute ΔHsolvent Enthalp y, H (ΔH lattice) Enthalp y, H Hexane Hexane Octane ΔHsolute separated separated separated + ΔHsolvent ΔHsolvent ΔHsolvent ΔHsolute Hexane NaCl( s) Hexane Octane ΔHmix aggregated aggregated aggregated aggregated H initial H initial Solution ΔHsoln ≈ 0 H final A NaCl. ΔHmix is much smaller than ΔHsolute: ΔHsoln is so much larger B Octane. ΔHsoln is very small, but the entropy increase due to mixing than the entropy increase due to mixing that NaCl does not dissolve. is large, so octane does dissolve.  13.4 Solubility as an Equilibrium Process    549 3. NH4NO3 in water. A large enough increase in entropy can sometimes cause a – solution to form even when the enthalpy increase is large (ΔHsoln ≫ 0). As we saw + + previously, in Figure 13.15C, when ammonium nitrate dissolves in water, the process is highly endothermic; that is, ΔHlattice ≫ ΔHhydr of the ions. Nevertheless, the increase – + – in entropy that occurs when the crystal breaks down and the ions mix with water – – + + molecules is greater than the increase in enthalpy: ΔSsoln > ΔHsoln. + – – + + – – In Chapter 20, we’ll return in depth to the relation between enthalpy and entropy to – + – – – + understand physical and chemical systems. + – + + – + › Summary of Section 13.3 – + – + › In a thermochemical solution cycle, the heat of solution is the sum of the endothermic Figure 13.17 Equilibrium in a saturated separations of solute and of solvent and the exothermic mixing of their particles. solution. At some temperature, the num- › In water, solvation (surrounding solute particles with solvent) is called hydration. For ions, ber of solute particles dissolving (white heats of hydration depend on the ion’s charge density but are always negative because ion- ­arrows) per unit time equals the number dipole forces are strong. Charge density exhibits periodic trends. recrystallizing (black arrows). › Systems naturally increase their entropy (distribute their energy in more ways). A gas has higher entropy than a liquid, which has higher entropy than a solid, and a solution has higher entropy than the pure solute and solvent. Seed crystal › Relative sizes of the enthalpy and entropy changes determine solution formation. A substance with a positive ΔHsoln dissolves only if ΔSsoln is larger than ΔHsoln. 13.4 SOLUBILITY AS AN EQUILIBRIUM PROCESS When an excess amount of solid is added to a solvent, particles leave the crystal, are surrounded by solvent, and move away. Some dissolved solute particles collide with undissolved solute and recrystallize, but, as long as the rate of dissolving is greater than A the rate of recrystallizing, the concentration rises. At a given temperature, when solid is dissolving at the same rate as dissolved particles are recrystallizing, the concentration remains constant and undissolved solute is in equilibrium with dissolved solute: Solute (undissolved) ⥫⥬ solute (dissolved) Figure 13.17 shows an ionic solid in equilibrium with dissolved cations and anions. (We learned about the concept of equilibrium in Section 4.7 and saw how it occurs between phases in Section 12.2.) Three terms express the extent of this solution process: B ∙ A saturated solution is at equilibrium and contains the maximum amount of dis- solved solute at a given temperature in the presence of undissolved solute. There- fore, if you filter off the solution and add more solute, the added solute doesn’t dissolve. ∙ An unsaturated solution contains less than the equilibrium concentration of dis- solved solute; add more solute, and more will dissolve until the solution is saturated. ∙ A supersaturated solution contains more than the equilibrium concentration and is unstable relative to the saturated solution. You can often prepare a supersaturated solution if the solute is more soluble at higher temperature. While heating, dissolve more than the amount required for saturation at some lower temperature, and then C slowly cool the solution. If the excess solute remains dissolved, the solution is supersaturated. Add a “seed” crystal of solute or tap the container, and the excess Figure 13.18 Sodium acetate crystalliz- solute crystallizes immediately, leaving behind a saturated solution (Figure 13.18). ing from a supersaturated solution. When a seed crystal of sodium acetate is added to a supersaturated solution of the com- Effect of Temperature on Solubility pound (A), solute begins to crystallize (B) and continues until the remaining You know that more sugar dissolves in hot tea than in iced tea; in fact, temperature ­solution is saturated (C). affects the solubility of most substances. Let’s examine the effects of temperature on Source: A–C: © McGraw-Hill Education/ the solubility of solids and gases.