General Chemistry I. Dr. PHAN TẠI HUÂN Faculty of Food Science and Technology Nong Lam University. Module 6: Solutions

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1 General Chemistry I Dr. PHAN TẠI HUÂN Faculty of Food Science and Technology Nong Lam University Module 6: Solutions Definitions (solutions: solvent, solute, solubility, concentration; interactions between solvent phase and solute molecules; properties of polar and non-polar solvents; properties of colloid). Chemical equilibrium for solutions (self-ionisation of water, solvation, ph; Acids and Bases; solution equilibrium with precipitation). 2 1

2 Solution A solution is defined as a homogeneous mixture of substances in which no settling occurs. A solution consists of a solvent and one or more solutes, whose proportions vary from one solution to another. The solvent is the medium in which the solutes are dissolved. The fundamental units of solutes are usually ions or molecules. Water is the most important solvent, and compounds dissolved in water are said to be in aqueous solution. 3 Solution In reality, any combination of the three states can be considered a solution. Usually a solution is formed by dissolving a solid (e.g., sugar) in a liquid (e.g., water). Air is a solution which is a mixture of various gases. Carbonated water (soda) is a mixture of a gas (CO 2 ) dissolved in a liquid (H 2 O). Even alloys such as gold-silver alloys are solutions containing two solids. A true solution is a solution which has only one solvent with one or more solutes. 4 2

3 The concept of solubility Solubility of a substance is defined as the amount of the substance that will dissolve in a particular solvent. Solubilities vary tremendously. At one extreme, some substances form solutions in all proportions and are said to be miscible. For example, acetone and water can be mixed in any proportion, from pure water to pure acetone. At the other extreme, a substance may be insoluble in another. One example is common salt, NaCl, whose solubility in gasoline is virtually zero. 5 Determinants of solubility Many combinations display solubility that is between the two extremes of miscible and insoluble. In other words, the substance dissolves, but there is a limit to the amount of solute that will dissolve in a given amount of solvent. Concentrations of solutions are expressed in terms of either the amount of solute present in a given mass or volume of solution, or the amount of solute dissolved in a given mass or volume of solvent. 6 3

4 Concentrations of solutions may be expressed in terms of percent by mass of solute, which gives the mass of solute per 100 mass units of solution. The gram is the usual mass unit. percent sulute = Percent by mass mass of solute mass of solution x 100% Thus, a solution that is 10% calcium gluconate, Ca(C 6 H 11 O 7 ) 2, by mass contains 10 grams of calcium gluconate in 100 grams of solution. This could be described as 10 grams of calcium gluconate in 90 grams of water. Unless otherwise specified, percent means percent by mass, and water is the solvent. 7 Molarity vs. Molality The Molarity, M, of a solution is defined as the number of moles of the solute per liter of solution. moles of solute Molarity = = liter of solution mol L The molality, m, of a solution is defined as the number of moles of the solute per kilogram of solvent. moles of solute Molality = kilograms of solvent 8 4

5 Normality The normality, N, of a solution is the number of equivalents of solute per liter of solution. The equivalent is usually defined in terms of a chemical reaction. For acid-base reactions, an equivalent is the amount of substance that will react or form 1 mole of hydrogen (H + ) or hydroxide (OH - ) ions. For redox (oxidation-reduction) reactions, an equivalent is the amount of substance that will react or form 1 mole of electrons. number of equivalents Normality = 1liter of solution N = nm 9 Exercise How many grams of H 2 O must be used to dissolve 50 grams of sucrose to prepare a 1.25 m solution of sucrose, C 12 H 22 O 11? Ans: 10 5

6 Hydrogen peroxide disinfectant typically contains 3.0% by mass. Assuming that the rest of the contents is water, what is the molality of this disinfectant? Ans: Exercise 11 Spontaneity of the dissolution process A process is favored by (1) a decrease in the energy of the system, which corresponds to an exothermic process, and (2) an increase in the disorder, or randomness, of the system. The energy change that accompanies a dissolution process is called the heat of solution, H solution. It depends mainly on how strongly solute and solvent particles interact. A negative value of H solution designates the release of heat. The main interactions that affect the dissolution of a solute in a solvent follow: Weak solute solute attractions favor solubility. Weak solvent solvent attractions favor solubility. Strong solvent solute attractions favor solubility. 12 6

7 Spontaneity of the dissolution process 13 Like dissolves Like Solubility is a complex phenomenon that depends on the balance of several properties. The general features of solubility are summarized by the expression like dissolves like. Substances that dissolve in each other usually have similar types of intermolecular interactions. One substance dissolves in another if the forces of attraction between the solute and the solvent are similar to the solvent solvent and solute solute interactions. 14 7

8 Dissolution of liquids Water and methanol are alike in that both substances contain O-H groups that form hydrogen bonds readily. When these liquids are mixed, H 2 O... H 2 O hydrogen bonds and CH 3 OH... CH 3 OH hydrogen bonds break, but H 2 O...CH 3 OH hydrogen bonds form. The net result is that the degree of hydrogen bonding in the solution is about the same as in either of the pure liquids, making these two liquids miscible. 15 Dissolution of liquids The intermolecular interactions of octane and cyclohexane are alike. Octane and cyclohexane have low polarities, so these molecules in the pure liquids are held together by the dispersion forces caused by their polarizable electron clouds. Dispersion forces in solutions of octane and cyclohexane are about the same as in the pure liquids. So these two liquids are miscible. 16 8

9 Dissolution of liquids Water and octane are not alike and nearly insoluble in each other. Octane does not form hydrogen bonds, so the only forces of attraction between water molecules and octane molecules are dispersion forces. Because hydrogen bonds are stronger than dispersion forces, the cost of disrupting the hydrogen-bonding network in water is far greater than the stability gained from octane water dispersion forces. 17 Dissolution of liquids Some liquids can interact with other substances in multiple ways. Acetone, for instance, has a polar CO bond and a three-carbon bonding framework. The bonding framework is similar to that of a hydrocarbon, so acetone mixes with cyclohexane and octane. The polar CO group makes acetone miscible with other polar molecules such as acetonitrile. The polar oxygen atom in acetone has lone pairs of electrons that can form hydrogen bonds with hydrogen atoms of ammonia or water. 18 9

10 Exercise Give a molecular explanation for the following trend in alcohol solubilities in water: n-propanol CH 3 CH 2 CH 2 OH Completely miscible n-butanol CH 3 CH 2 CH 2 CH 2 OH 1.1 M n-pentanol CH 3 CH 2 CH 2 CH 2 CH 2 OH 0.30 M n-hexanol CH 3 CH 2 CH 2 CH 2 CH 2 CH 2 OH M Strategy Solubility limits depend on the stabilization generated by solute solvent interactions balanced against the destabilization that occurs when solvent solvent interactions are disrupted by solute. Intermolecular interactions involving water and alcohol molecules must be examined. 19 Solubility of solids: network solids Network solids such as diamond, graphite, or silica cannot dissolve without breaking covalent chemical bonds. Because intermolecular forces of attraction are always much weaker than covalent bonds, solvent solute interactions are never strong enough to offset the energy cost of breaking bonds. Covalent solids are insoluble in all solvents, although they may react with specific liquids or vapors

11 Dissolution of solids: metalic solids Metals do not dissolve in water, because they contain extensive delocalized bonding networks that must be disrupted before the metal can dissolve. A few metals react with water, and several react with aqueous acids, but no metal will simply dissolve in water. Likewise, metals do not dissolve in nonpolar liquid solvents. Zn(s) + 2H 3 O + (aq) Zn 2+ (aq) + H 2 (g) + 2H 2 O(l) The aqueous medium dissolves the metal by a chemical reaction that converts the insoluble metal into soluble cations. The solution produced when zinc reacts with aqueous HCl is an aqueous solution of ions, not a solution of Zn metal in water. 21 Dissolution of solids: molecular solids At the opposite extreme, molecular solids contain individual molecules bound together by various combinations of dispersion forces, dipole forces, and hydrogen bonds. Conforming to like dissolves like, molecular solids dissolve readily in solvents with similar types of intermolecular forces. Nonpolar, for instance, is soluble in nonpolar liquids such as carbon tetrachloride. Many organic compounds are molecular solids that dissolve in organic liquids such as cyclohexane and acetone

12 Dissolution of solids: molecular solids Hydrogen bonding makes sugars such as sucrose and glucose highly soluble in water. When glucose dissolves, hydrogen bonds between water and glucose replace the hydrogen bonds lost by the water molecules of the solvent. This balance means that the energy requirements for solution formation are small, and glucose is quite soluble in water. 23 Dissolution of solids: molecular solids Hydrogen bonding allows water to dissolve materials that form hydrogen bonds. On the other hand, naphthalene, a similarly sized solid hydrocarbon limited to dispersion forces, is nearly insoluble in water. The best solvent for a molecular solid is one whose intermolecular forces match the forces holding the molecules in the crystal. For a solid held together by dispersion forces, good solvents are nonpolar liquids such as carbon tetrachloride (CCl 4 ) and cyclohexane (C 6 H 12 ). For polar solids, a polar solvent such as acetone works well

13 Example: solubility of vitamin 25 Solubility of solids: ionic solids The ability of an ionic solid to go into solution depends most strongly on its crystal lattice energy, or the strength of attractions among the particles making up the solid. Crystal lattice energies are always negative: M(g) + X(g) MX(s) + energy If the solvent is water, the energy that must be supplied to expand the solvent includes that required to break up some of the hydrogen bonding between water molecules. The third major factor contributing to the heat of solution is the extent to which solvent molecules interact with particles of the solid. The process in which solvent molecules surround and interact with solute ions or molecules is called solvation. When the solvent is water, the more specific term is hydration

14 27 Solubility of solids: ionic solids A cluster of water molecules surrounds each ion in solution. Notice how the water molecules are oriented so that their dipole moments align with charges of the ions. The partially negative oxygen atoms of water molecules point toward cations, whereas the partially positive hydrogen atoms of water molecules point toward anions

15 Solubility of solids: ionic solids Hydration energy is defined as the energy change involved in the (exothermic) hydration of one mole of gaseous ions. M n+ (g) + xh 2 O M(H 2 O) x n+ + energy (for cation) X y- (g) + rh 2 O X(H 2 O) r y- + energy (for anion) Hydration is usually highly exothermic for ionic or polar covalent compounds, because the polar water molecules interact very strongly with ions and polar molecules. The overall heat of solution for a solid dissolving in a liquid is equal to the heat of solvation minus the crystal lattic energy. H solution = (heat of solvation) - (crystal lattice energy) 29 Solubility of solids: ionic solids Magnitudes of crystal lattice and hydration energies generally increase with increasing charge and decreasing size of ions (ionic charge densities increase). Hydration energies and lattice energies are usually of about the same magnitude for low-charge species, so the dissolution process is slightly endothermic for many ionic substances. As the charge-to-size ratio (charge density) increases for ions in ionic solids, the magnitude of the crystal lattice energy usually increases more than the hydration energy. For ex. aluminum fluoride, AlF 3 ; magnesium oxide, MgO; and chromium(iii) oxide, Cr 2 O 3 are very endothermic and not very soluble in water

16 Ionic radii, charge/radius ratios, and hydration energies for some cations 31 Dissolution of gases The only gases that dissolve appreciably in water are: (1) those that are capable of hydrogen bonding (such as HF), (2) those that ionize (such as HCl, HBr, and HI), (3) those that react with water (such as CO 2 )

17 Rates of dissolution and saturation When a solid is placed in water, some of its particles solvate and dissolve. The rate of this process slows as time passes because the surface area of the crystals gets smaller and smaller. At the same time, the number of solute particles in solution increases, so they collide with the solid more frequently. Some of these collisions result in recrystallization. The rates of the two opposing processes become equal after some time. The solid and dissolved ions are then in equilibrium with each other. After equilibrium is established, no more solid dissolves without the simultaneous crystallization of an equal mass of dissolved ions. 33 Rates of dissolution and saturation The solubilities of many solids increase at higher temperatures. Supersaturated solutions contain higher-than-saturated concentrations of solute. The saturated solution is cooled slowly, without agitation, to a temperature at which the solute is less soluble. At this point, the resulting supersaturated solution is metastable (temporarily stable). A supersaturated solution produces crystals rapidly if it is slightly disturbed or if it is seeded with a dust particle or a tiny crystal

18 Effect of temperature on solubility LeChatelier s Principle: A system at equilibrium, or changing toward equilibrium, responds in the way that tends to relieve or undo any stress placed on it. Many ionic solids dissolve by endothermic processes. Their solubilities in water usually increase as heat is added and the temperature increases. Endothermic: reactants + heat products For example:

19 Effect of pressure on solubility Changing the pressure has no appreciable effect on the solubilities of either solids or liquids in liquids. The solubilities of gases in all solvents increase, however, as the partial pressures of the gases increase. Henry s Law: The pressure of a gas above the surface of a solution is proportional to the concentration of the gas in the solution. P gas = kc gas The relationship is valid at low concentrations and low pressures. 37 Colloids Particles whose dimensions are between 1 nanometer and 1 micrometer, called colloids, are larger than the typical molecule but smaller than can be seen under an optical microscope. When a colloid is mixed with a second substance, the colloid can become uniformly spread out, or dispersed, throughout the dispersing medium. Such a dispersion is a colloidal suspension that has properties intermediate between those of a true solution and those of a heterogeneous mixture. Mixture Example Approximate Particle Size suspension sand in water larger than 10,000 Å colloidal dispersion starch in water 10 10,000 Å solution sugar in water 1 10 Å 38 19

20 Types of colloids 39 Aqueous solutions: an introduction Approximately 3/4 of the earth s surface is covered with water. The body fluids of all plants and animals are mainly water. Many important chemical reactions occur in aqueous (water) solutions, or in contact with water

21 Electrolytes and extent of ionization Solutes that are water-soluble can be classified as either electrolytes or nonelectrolytes. Electrolytes are substances whose aqueous solutions conduct electric current. Strong electrolytes are substances that conduct electricity well in dilute aqueous solution. Weak electrolytes conduct electricity poorly in dilute aqueous solution. Aqueous solutions of nonelectrolytes do not conduct electricity. Electric current is carried through aqueous solution by the movement of ions. The strength of an electrolyte depends on the number of ions in solution and also on the charges on these ions. 41 Electrolytes and extent of ionization Three major classes of solutes are strong electrolytes: (1) strong acids, (2) strong bases, and (3) most soluble salts. These compounds are completely or nearly completely ionized (or dissociated) in dilute aqueous solutions, and therefore are strong electrolytes. Dissociation refers to the process in which a solid ionic compound, such as NaCl, separates into its ions in solution: Ionization refers to the process in which a molecular compound separates or reacts with water to form ions in solution: 42 21

22 Acid Base The Arrhenius theory (1884) An acid is a substance that contains hydrogen and produces H in aqueous solution. A base is a substance that contains the OH (hydroxyl) group and produces hydroxide ions, OH, in aqueous solution. Neutralization is defined as the combination of H ions with OH ions to form H 2 O molecules. H + (aq) + OH - (aq) H 2 O (l) (neutralization) The hydronium ion (hydrated hydrogen ion) The hydrated hydrogen ion is the species that gives aqueous solutions of acids their characteristic acidic properties. 43 Acid Base The Brønsted Lowry theory (1923) An acid is defined as a proton donor and a base is defined as a proton acceptor. An acid base reaction is the transfer of a proton from an acid to a base. The ionization of hydrogen chloride, HCl, in water is an acid base reaction in which water acts as a base or proton acceptor. We can describe Brønsted Lowry acid base reactions in terms of conjugate acid base pairs. These are two species that differ by a proton

23 45 Acid Base The Lewis theory (1923) An acid is any species that can accept a share in an electron pair. A base is any species that can make available, or donate, a share in an electron pair. These definitions do not specify that an electron pair must be transferred from one atom to another only that an electron pair, residing originally on one atom, must be shared between two atoms. Neutralization is defined as coordinate covalent bond formation

24 The autoionization of water Water is said to be amphiprotic; that is, H 2 O molecules can both donate and accept protons. 47 Chemical equilibrium Reactions that do not go to completion and that can occur in either direction are called reversible reactions. rate k [ C][ c D] d f = [ ] a [ ] b f rate r = k r A B At equilibrium, rate f = rate r k k f = r c d [ C] eq [ D ] eq a [ A ] [ ] b eq B eq K c = c d [ C] eq [ D ] eq a [ A ] [ ] b eq B eq (For any pure liquid or pure solid, the activity is taken as 1.) Both k f and k r are constant, so k f /k r is also a constant and given a special name and symbol the equilibrium constant, K c or simply K. ΔG o = RTlnK 48 24

25 The autoionization of water The equilibrium constant is known as the ion product for water and is usually represented as K w. K w = [H 3 O + ][OH - ] Careful measurements show that, in pure water at 25 C, [H 3 O + ]=[OH - ]= 1.0 x 10-7 mol/l K w = [H 3 O + ][OH - ]= 1.0 x (at 25 C) Although this expression was obtained for pure water, it is also valid for dilute aqueous solutions at 25 C. This is one of the most useful relationships chemists have discovered. It gives a simple relationship between H 3 O + and OH - concentrations in all dilute aqueous solutions. 49 Exercise Calculate the concentrations of H 3 O + and OH - ions in a 0.05 M HNO 3 solution. Ans: 50 25

26 The ph and poh scales In common chemical applications, the concentration of hydronium ion, the measure of a solution s acidity, ranges from fairly large (1 M) to very small (10-14 M). The ph and poh scales provide a convenient way to express the acidity and basicity of dilute aqueous solutions that avoids using very small numbers. The ph and poh of a solution are defined as ph = - log [H 3 O + ] or [H 3 O + ] = 10 -ph poh= - log [OH - ] or [OH - ] = 10 -poh It is convenient to describe the autoionization of water in terms of pkw. pk w = -log K w 51 The ph and poh scales [H 3 O + ][OH - ]= 1.0 x (at 25 C) ph + poh =

27 Calculate [H 3 O + ], ph, [OH - ], and poh for a M Ca(OH) 2 solution. Ans: Exercise 53 The ph and poh scales 54 27

28 Strong and weak acids Strong acids ionize (separate into hydrogen ions and stable anions) completely, or very nearly completely, in dilute aqueous solution. 55 Strong and weak acids Weak acids cannot dissociate completely. They undergo the same type of dissociation as that of strong acids and bases, but the extent of dissociation is very little compared to strong acid or strong base dissociations (usually less than 5%). Acid-ionization constant (K a ): K a = + - [ H O ][ A ] 3 [ HA] 56 28

29 Strong and weak acids 57 Find the degree of ionization of 0.1 M solution of acetic acid (CH,COOH). Also find the ph of the solution. (The acid-ionization constant of acetic acid is 1.8 x ) Ans: Exercise 58 29

30 Strong bases, insoluble bases, and weak bases Most common bases are ionic metal hydroxides in the solid state. Strong bases are soluble in water and are dissociated completely in dilute aqueous solution. Common strong bases 59 Strong bases, insoluble bases, and weak bases Other metals form ionic hydroxides, but these are so sparingly soluble in water that they cannot produce strongly basic solutions. They are called insoluble bases or sometimes sparingly soluble bases. Typical examples include Cu(OH) 2, Zn(OH) 2, Fe(OH) 2, and Fe(OH) 3. Common weak bases are molecular substances that are soluble in water but form only low concentrations of ions in solution. The most common weak base is ammonia, NH 3. Just like the acid-ionization constant, there is also the baseionization (K b ) + - [ NH ][ ] 4 OH K b = NH [ ]

31 Dissociation of polyprotic acids Some acids have two or more protons that can be released upon dissociation. Such acids are called polyprotic acids. Sulfuric acid is a polyprotic acid that can lose two protons in solution. The first ionization is complete because sulfuric acid is a strong acid. Dissociation constant, K a l = very large In the second ionization, an equilibrium exists because hydrogen sulfate ion (HSO 4- ) is not as strong as H 2 SO 4. Dissociation constant, K a 2 = 1.2 x l Differentiating acidic and basic salts Salt solutions can be acidic, neutral, or basic. A salt of a weak acid and a strong base gives a basic aqueous solution. NaCN is a salt of a weak acid (HCN) and a strong base (NaOH). A salt of a weak base and a strong acid gives an acidic aqueous solution. Zn(NO) 3 is salt of a weak base (Zn(OH) 3 ) and a strong acid (HNO 3 ). A salt of a strong base and a strong acid gives a neutral aqueous solution. NaCl is a salt of a strong acid (HC1) and a strong base (NaOH)

32 Common-ion effect When a salt is added to a solution containing either the same cation or anion, there will be changes in the solubility because of what is commonly known as common-ion effect. The phenomenon can be best explained in terrns of Le Chatelier's principle. For ex. adding magnesium fluoride (MgF 2 ) to a solution of sodium fluoride (NaF). 63 Solubility guidelines for common ionic compounds in water Compounds whose solubility in water is less than about 0.02 mol/l are usually classified as insoluble compounds. No gaseous or solid substances are infinitely soluble in water

33 Solubility guidelines for common ionic compounds in water 65 Bonding, solubility, electrolyte characteristics, and predominant forms of solutes in contact with water 66 33

34 Solubility product principle In equilibria that involve slightly soluble compounds in water, the equilibrium constant is called a solubility product constant, K sp,or solubility product. K sp = [M z+ ] y [X y- ] z In general, the solubility product expression for a compound is the product of the concentrations of its constituent ions, each raised to the power that corresponds to the number of ions in one formula unit of the compound. The quantity is constant at constant temperature for a saturated solution of the compound. 67 Ion product Ion product (reaction quotient, Q sp ) is the product of the concentrations of the ions from the compound (solute) in a solution, each concentration raised to a power equal to its coefficient in the balanced equation. In other words, the expression for the ion product is the same as that of K sp. By comparing the ion product of a compound against its K sp, we can predict whether or not precipitation is likely to occur

35 If 100 ml of M sodium sulfate, Na 2 SO 4, is mixed with 50 ml of M barium chloride, BaCl 2, will a precipitate form? (K sp for BaSO 4 = 1.1 x ). Ans: Exercise 69 Buffers A solution which resists changes in ph when small amounts of acid or base are added to it is called a buffer solution. A buffer solution is usually a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid. A buffer solution contains equilibrium amounts of acid and base species. If a strong acid is added to this buffer solution, the H + concentration increases and they react with the base (HCO 3 ) - thereby decreasing the H + concentration and maintaining the initial ph. On the other hand, if a strong base is added, the OH - ions supplied by it will react with the H + so that more of H 2 CO 3, will dissociate thereby restoring the initial ph

36 Buffers do not have unlimited capacity to resist phchange. The buffer capacity of a buffer depends on the nature of the buffer and the amount of acid and conjugate base present in the solution. The Henderson-Hasselbalch equation can be used to relate the ph of a buffer and the concentrations of base and acid. For a weak acid-conjugate base buffer: ph = pk + log a [ conjugate base] [ weak acid] For a weak base-conjugate acid buffer. poh = pk + log b Buffers [ conjugate acid] [ weak base] 71 Calculate the concentration of H 3 O + and the ph of a buffer solution that is 0.10 M in CH 3 COOH and 0.20 M in NaCH 3 COO. Ans: Exercise 72 36

37 Titration curves of acids and bases Consider the reaction involving 50 ml of a 0.1 M solution of HC1, titrated with 0.1 M solution of NaOH. 73 Titration curves of acids and bases Acid-base titrations are reactions by which we can determine the amount of acid or base present in a solution. This is done by reacting the solution with a base or acid (of known concentration), and by measuring the volume of the known acid or base used up in the process. The equivalence point denotes the point at which equivalent amounts of acid and base have reacted. To know the equivalence point, we usually add an indicator which will change its color close to the equivalence point. An indicator is an organic dye; its color depends on the concentration of H 3 O + ions, or ph, in the solution. By the color an indicator displays, it indicates the acidity or basicity of a solution

38 Indicator Methyl red is red at ph 4 and below; it is yellow at ph 7 and above. Between ph 4 and ph 7 it changes from red to redorange, to orange, to yellow. 75 Indicator Bromthymol blue is yellow at ph 6 and below; it is blue at ph 8 and above. Between ph 6 and 8 it changes from yellow to yellowgreen, to green, to blue-green, to blue

39 Indicator Phenolphthalein is colorless below ph 8 and bright pink above ph 10. It changes from colorless to pale pink, to pink, to bright pink in the ph range 8 to Indicator The ph at which the color change occurs is characteristic of each indicator. For an acid-base reaction, the indicator is chosen based on the ph at which the equivalence point is expected to occur. When the indicator is in an acidic solution, the equilibrium shifts to the left (LeChatelier's principle), and the predominant species is HIn making the indicator show yellow color. In a basic solution, the equilibrium shifts favoring the forward reaction and the predominant species is In - (blue color)

40 Some Common Indicators 79 Summary After you have studied this module 6, you should be able to Describe the solution. Express concentrations of solutions in terms of molality and mole fractions. Describe the factors that favor the dissolution process. Describe the dissolution of solids in liquids, liquids in liquids, and gases in liquids. Describe how temperature and pressure affect solubility. Recognize and describe colloids: the Tyndall effect, the adsorption phenomenon, hydrophilic and hydrophobic colloids

41 Summary Recognize strong electrolytes and calculate concentrations of their ions. Recognize and classify acids (strong, weak), bases (strong, weak, insoluble), and salts (soluble, insoluble); use the solubility guidelines. Understand the autoionization of water. Understand the ph and poh scales and how they are used. Use ionization constants. Describe how polyprotic acids ionize in steps. 81 Summary Explain the common ion effect. Understand the solution equilibrium with precipitation. Recognize buffer solutions and describe their chemistry. Describe how to prepare a buffer solution of a specified ph. Carry out calculations related to buffer solutions and their action. Explain what acid base indicators are and how they function