Contents and Concepts

Chapter 16 1 Learning Objectives Acid Base Concepts Arrhenius Concept of Acids and Base a. Define acid and base according to the Arrhenius concept. Brønsted Lowry Concept of Acids and Bases a. Define acid and base according to the Brønsted Lowry concept. b. Define the term conjugate acid base pair. c. Identify acid and base species. d. Define amphiprotic species. 4 Contents and Concepts Acid Base Concepts 1. Arrhenius Concept of Acids and Bases. Brønsted Lowry Concept of Acids and Bases 3. Lewis Concept of Acids and Bases Acid and Base Strengths 4. Relative Strengths of Acids and Bases 5. Molecular Structure and Acid Strength 3. Lewis Concept of Acids and Bases a. Define Lewis acid and Lewis base. b. Identify Lewis acid and Lewis base species. Acid and Base Strengths 4. Relative Strengths of Acids and Bases a. Understand the relationship between the strength of an acid and that of its conjugate base. b. Decide whether reactants or products are favored in an acid base reaction. 5 Self-Ionization of Water and ph 6. Self-Ionization of Water 7. Solutions of a Strong Acid or Base 8.The ph of a Solution 5. Molecular Structure and Acid Strength a. Note the two factors that determine relative acid strengths. b. Understand the periodic trends in the strengths of the binary acids HX. c. Understand the rules for determining the relative strengths of oxoacids. d. Understand the relative acid strengths of a polyprotic acid and its anions. 3 6 1

Self-Ionization of Water and ph 6. Self-Ionization of Water a. Define self-ionization (or autoionization). b. Define the ion-product constant for water. 7. Solutions of a Strong Acid or Base a. Calculate the concentrations of H 3 O and OH - in solutions of a strong acid or base When gaseous hydrogen chloride meets gaseous ammonia, a smoke composed of ammonium chloride is formed. HCl(g) NH 3 (g) NH 4 Cl(s) This is an acid base reaction. 7 8. The ph of a Solution 1. Define ph.. Calculate the ph from the hydronium-ion concentration. 3. Calculate the hydronium-ion concentration from the ph. 4. Describe the determination of ph by a ph meter and by acid base indicators. We will examine three ways to explain acid base behavior: Arrhenius Concept H and OH - Brønsted Lowry Concept H = proton Lewis Concept electron pair donor acceptor donor acceptor Note: H in water is H 3 O acid base 8 11 Acid-Base Concepts Antoine Lavoisier was one of the first chemists to try to explain what makes a substance acidic. In 1777, he proposed that oxygen was an essential element in acids. 9 The actual cause of acidity and basicity was ultimately explained in terms of the effect these compounds have on water by Svante Arrhenius in 1884. 1

Acid-Base Concepts In the first part of this chapter we will look at several concepts of acid-base theory including: The Arrhenius concept The Bronsted Lowry concept The Lewis concept This chapter expands on what you learned in Chapter 3 about acids and bases. 13 Arrhenius Concept of Acids and Bases According to the Arrhenius concept of acids and bases, an acid is a substance that, when dissolved in water, increases the concentration of hydronium ion (H 3 O ). The H 3 O is shown here hydrogen bonded to three water molecules. 16 Arrhenius Concept of Acids and Bases According to the Arrhenius concept of acids and bases, an acid is a substance that, when dissolved in water, increases the concentration of hydronium ion (H 3 O ). Arrhenius Concept of Acids and Bases A base, in the Arrhenius concept, is a substance that, when dissolved in water, increases the concentration of hydroxide ion, OH -. Chemists often use the notation H for the H 3 O ion, and call it the hydrogen ion. Remember, however, that the aqueous hydrogen ion is actually chemically bonded to water, that is, H 3 O. 14 17 Arrhenius Concept of Acids and Bases The Arrhenius concept limits bases to compounds that contain a hydroxide ion. The Brønsted Lowry concept expands the compounds that can be considered acids and bases. In the Arrhenius concept, a strong acid is a substance that ionizes completely in aqueous solution to give H 3 O and an anion An example is perchloric acid, HClO 4. HClO H O(l) 4 H3O ClO4 Other strong acids include HCl, HBr, HI, HNO 3, and H SO 4. 15 18 3

Arrhenius Concept of Acids and Bases In the Arrhenius concept, a strong base is a substance that ionizes completely in aqueous solution to give OH - and a cation. An example is sodium hydroxide, NaOH. NaOH(s) H O Na OH Other strong bases include LiOH, KOH, Ca(OH), Sr(OH), and Ba(OH). Brønsted-Lowry Concept of Acids and Bases According to the Brønsted-Lowry concept, an acid is the species donating the proton in a proton-transfer reaction. A base is the species accepting the proton in a proton-transfer reaction. In any reversible acid-base reaction, both forward and reverse reactions involve proton transfer. 19 Arrhenius Concept of Acids and Bases Most other acids and bases that you encounter are weak. They are not completely ionized and exist in reversible reaction with the corresponding ions. An example is acetic acid, HC H 3 O. Consider the reaction of NH 3 and H 0. NH 3 HO(l) NH4 base acid NH 3 HO(l) NH4 OH OH HC H3O HO(l) H3O CH3O Ammonium hydroxide, NH 4 OH, is a weak base. NH4OH NH4 OH 0 H In the forward reaction, NH 3 accepts a proton from H O. Thus, NH 3 is a base and H O is an acid. 3 Brønsted Lowry Concept of Acids and Bases An acid base reaction is considered a proton (H ) transfer reaction. H H H H 1 4 4

Consider the reaction of NH 3 and H O. acid NH 3 HO(l) NH4 base OH What is the conjugate acid of H O? What is the conjugate base of H O? base H acid NH 3 HO(l) NH4 OH The species NH 4 and NH 3 are a conjugate acid-base pair. A conjugate acid-base pair consists of two species in an acid-base reaction, one acid and one base, that differ by the loss or gain of a proton. 5 The conjugate acid of H O has gained a proton. It is H 3 O. The conjugate base of H O has lost a proton. It is OH -. 8 Brønsted-Lowry Concept of Acids and Bases Consider the reaction of NH 3 and H O. base acid NH 3 HO(l) NH4 OH Here NH 4 is the conjugate acid of NH 3 and NH 3 is the conjugate base of NH 4. The Brønsted-Lowry concept defines a species as an acid or a base according to its function in the proton-transfer 6 reaction. Label each species as an acid or base. Identify the conjugate acid-base pairs. a. HCO 3- HF H CO 3 F - Base Acid Conjugate Conjugate acid base b. HCO 3- OH - CO - 3 H O(l) Acid Base Conjugate Conjugate base acid 9 Substances in the acid base reaction that differ by the gain or loss of a proton, H, are called a conjugate acid base pair. The acid is called the conjugate acid; the base is called a conjugate base. A Brønsted Lowry acid is the species donating a proton in a proton-transfer reaction; it is a proton donor. Acid Base Conjugate Conjugate base acid 7 A Brønsted Lowry base is the species accepting a proton in a proton-transfer reaction; it is a proton acceptor. 30 5

Some species can act as an acid or a base. An amphoteric species is a species that can act either as an acid or a base (it can gain or lose a proton). For example, HCO 3- acts as a proton donor (an acid) in the presence of OH - HCO OH CO 3 3 H O(l) In the Brønsted-Lowry concept: 1. A base is a species that accepts protons; OH - is only one example of a base.. Acids and bases can be ions as well as molecular substances. 3. Acid-base reactions are not restricted to aqueous solution. 4. Some species can act as either acids or bases depending on what the other reactant is. H 31 Practice exercise 16.7 and 16.8 34 An amphoteric species is a species that can act either as an acid or a base (it can gain or lose a proton). Alternatively, HCO 3- can act as a proton acceptor (a base) in the presence of HF. HCO HF 3 HCO3 F H The amphoteric characteristic of water is important in the acid-base properties of aqueous solutions. 3 Lewis Concept of Acids and Bases The Lewis concept defines an acid as an electron pair acceptor and a base as an electron pair donor. This concept broadened the scope of acidbase theory to include reactions that did not involve H. The Lewis concept embraces many reactions that we might not think of as acid-base reactions. 35 Water reacts as an acid with the base NH 3. NH H O(l) 3 NH 4 H OH Water can also react as a base with the acid HF. The reaction of boron trifluoride with ammonia is an example. : : F B : : : F: : F: : H : N H H : : F B : : : F: : F: : H N H H HF H O(l) F H3O H Boron trifluoride accepts the electron pair, so it is a Lewis acid. Ammonia donates the electron pair, so it is the Lewis base. 33 36 6

Relative Strength of Acids and Bases The stronger acids are those that lose their hydrogen ions more easily than other acids. Similarly, the stronger bases are those that hold onto hydrogen ions more strongly than other bases. If an acid loses its H, the resulting anion is now in a position to reaccept a proton, making it a Brønsted-Lowry base. It is logical to assume that if an acid is considered strong, its conjugate base (that is, its anion) would be weak, since it is unlikely to accept a hydrogen ion. 37 40 Name and define the three acid base concepts we have discussed. Give examples of each type of acid-base. Relative Strength of Acids and Bases Consider the equilibrium below. HC H3O HO(l) acid base H3O CH3O acid base conjugate acid-base pairs In this system we have two opposing Brønsted- Lowry acid-base reactions. In this example, H 3 O is the stronger of the two acids. Consequently, the equilibrium is skewed toward reactants. 38 41 Relative Strength of Acids and Bases The Brønsted-Lowry concept introduced the idea of conjugate acid-base pairs and proton-transfer reactions. We consider such acid-base reactions to be a competition between species for hydrogen ions. From this point of view, we can order acids by their relative strength as hydrogen ion donors. 39 Consider the equilibrium below. HC H3O HO(l) acid base conjugate acid-base pairs H O 3 CH3O acid base Table 15. outlines the relative strength of some common acids and their conjugate bases. This concept of conjugate pairs is fundamental to understanding why certain salts can act as acids or bases. See Exercises 11.4-7 Problems 11.46-65 4 7

Table 11.3 and 11.4 in Text Practice Exercises 16.8 16.19 43 46 1. What is the conjugate acid of NH 3? a) H 3 O b) HCl c) NH 4 d) H O Ans : c. Which of the following is the weakest Brønsted acid? a) NH 4 K a = 5.6 - b) HSO - 4 K a = 1. - c) HF K a = 7. -4 d) HPO - 4 K a = 4. -13 44 Ans: d 47 Molecular Structure and Acid Strength Two factors are important in determining the relative acid strengths. One is the polarity of the bond to which the hydrogen atom is attached. a. Acetate ion b. Formate ion Acetate Ion 45 The H atom should have a partial positive charge: δ δ H X The more polarized the bond, the more easily the proton is removed and the greater the acid strength. 48 8

The second factor is the strength of the bond. Or, in other words, how tightly the proton is held. This depends on the size of atom X. H O and H S These are binary acids from the same group, so we compare the size of O and S. Because S is larger, H S is the stronger acid. δ δ- H X The larger atom X, the weaker the bond and the greater the acid strength. HCl and H S These are binary acids from the same period, but different groups, so we compare the electronegativity of Cl and S. Because Cl is more electronegative, HCl is the stronger acid. 49 5 For a binary acid, as the size of X in HX increases, going down a group, acid strength increases. For a binary acid, going across a period, as the electronegativity increases, acid strength increases. Molecular Structure and Acid Strength Consider a series of binary acids from a given column of elements. As you go down the column of elements, the radius increases markedly and the H-X bond strength decreases. You can predict the following order of acidic strength. HF < HCl < HBr < HI 50 53 Which is a stronger acid: HF or HCl? Which is a stronger acid: H O or H S? Which is a stronger acid: HCl or H S? HF and HCl These are binary acids from the same group, so we compare the size of F and Cl. Because Cl is larger, HCl is the stronger acid. 51 As you go across a row of elements, the polarity of the H-X bond becomes the dominant factor. N < 3 H O < HF H As electronegativity increases going to the right, the polarity of the H-X bond increases and the acid strength increases. You can predict the following order of acidic strength. 3 N < H O < HF H 54 9

Consider the oxoacids. An oxoacid has the structure: H O Y The acidic H atom is always attached to an O atom, which in turn is attached to another atom Y. Bond polarity is the dominant factor in the relative strength of oxoacids. As a result, the H atom becomes more acidic. The acid strengths of the oxoacids of chlorine increase in the following order. HClO < HClO < < HClO3 HClO4 This, in turn, depends on the electronegativity of the atom Y. 55 58 Consider the oxoacids. An oxoacid has the structure: H O Y If the electronegativity of Y is large, then the O-H bond is relatively polar and the acid strength is greater. You can predict the following order of acidic strength. HOCl > HOBr > HOI Other groups, such as O atoms or O-H groups, may be attached to Y. With each additional O atom, Y becomes effectively more electronegative. 56 For a series of oxoacids differing only in the central atom Y, the acid strength increases with the electronegativity of Y. Stronger Weaker 59 For oxoacids, several factors are relevant: the number and bonding of oxygens, the central element, and the charge on the species. For a series of oxoacids, (OH) m YO n, acid strength increases as n increases. (OH)Cl n = 0 (OH)ClO n = 1 (OH)ClO n = (OH)ClO 3 n = 3 Molecular Structure and Acid Strength Consider polyprotic acids and their corresponding anions. Each successive H atom becomes more difficult to remove. Therefore the acid strength of a polyprotic acid and its anions decreases with increasing negative charge. Weakest Strongest 57 < 4 HPO4 H3PO4 HPO < 60

The acid strength of a polyprotic acid and its anions decreases with increasing negative charge. H CO 3 is a stronger acid than HCO 3-. H SO 4 is a stronger acid than HSO 4-. H 3 PO 4 is a stronger acid than H PO 4-. H PO 4- is a stronger acid than HPO 4 -. A reaction will always go in the direction from stronger acid to weaker acid, and from stronger base to weaker base. 61 Self-Ionization of Water H O(l) H O(l) H 3 O OH - Base Acid Conjugate acid Conjugate base 64 Decide which species are favored at the completion of the following reaction: HCN HSO 3- CN - H SO 3 We first identify the acid on each side of the reaction: HCN and H SO 3. Next, we compare their acid strength: H SO 3 is stronger. This reaction will go from right to left ( ), and the reactants are favored. 6 Self-ionization of Water Self-ionization is a reaction in which two like molecules react to give ions. In the case of water, the following equilibrium is established. H O(l) H O(l) H O (aq ) 3 OH (aq ) The equilibrium-constant expression for this system is: [H3O ][OH ] K = c [H O] 65 Self-ionization of Water Self-ionization is a reaction in which two like molecules react to give ions.. The concentration of ions is extremely small, so the concentration of H O remains essentially constant. This gives: H O] K = [H O ][OH [ c 3 constant ] 63 66 11

We call the equilibrium value for the ion product [H 3 O ][OH - ] the ion-product constant for water, which is written K w. = [H O K w 3 ][OH At 5 o C, the value of K w is 1.0 x -14. Like any equilibrium constant, K w varies with temperature. Because we often write H 3 O as H, the ionproduct constant expression for water can be written: K w = [H ][OH ] ] 67 Solutions of Strong Acid or Base In a solution of a strong acid you can normally ignore the self-ionization of water as a source of H. The H concentration is usually determined by the strong acid concentration. However, the self-ionization still exists and is responsible for a small concentration of OH - ion. 70 Self-ionization of Water These ions are produced in equal numbers in pure water, so if we let x = [H ] = [OH - ] 1.0 x = 14 = 1.0 (x)(x) 14 at 5 = 1.0 7 Thus, the concentrations of H and OH - in pure water are both 1.0 x -7 M. If you add acid or base to water they are no longer equal but the K w expression still holds. o C 68 As an example, calculate the concentration of OH - ion in 0. M HCl. Because you started with 0. M HCl (a strong acid) the reaction will produce 0. M H. HCl H Cl Substituting [H ]=0. into the ion-product expression, we get: 1.0 = (0.)[OH 14 ] 71 H O(l) H O(l) H 3 O OH - We call the equilibrium constant the ionproduct constant, K w. K w = [H 3 O ][OH - ] Solutions of Strong Acid or Base As an example, calculate the concentration of OH - ion in 0. M HCl. Because you started with 0. M HCl (a strong acid) the reaction will produce 0. M H. HCl H Cl At 5 C, K w = 1.0-14 Substituting [H ]=0. into the ion-product expression, we get: As temperature increases, the value of K w increases. 69-14 1.0 [ OH ] = = 1.0 0. -13 M 7 1

Similarly, in a solution of a strong base you can normally ignore the self-ionization of water as a source of OH -. The OH - concentration is usually determined by the strong base concentration. However, the self-ionization still exists and is responsible for a small concentration of H ion. 73 Solutions of Strong Acid or Base By dissolving substances in water, you can alter the concentrations of H and OH -. In a neutral solution, the concentrations of H and OH - are equal, as they are in pure water. In an acidic solution, the concentration of H is greater than that of OH -. In a basic solution, the concentration of OH - is greater than that of H. 76 As an example, calculate the concentration of H ion in 0.0 M NaOH. Because you started with 0.0 M NaOH (a strong base) the reaction will produce 0.0 M OH -. H NaOH(s) O Na OH At 5 C, you observe the following conditions. In an acidic solution, [H ] > 1.0 x -7 M. In a neutral solution, [H ] = 1.0 x -7 M. In a basic solution, [H ] < 1.0 x -7 M. Substituting [OH - ]=0.0 into the ionproduct expression, we get: 1.0 = [H 14 ](0.0) 74 77 Because you started with 0.0 M NaOH (a strong base) the reaction will produce 0.0 M OH -. H NaOH(s) O Na OH Substituting [OH - ]=0.0 into the ionproduct expression, we get: -14 1.0 [ H ] = = 1.0 0.0-1 M The ph of a Solution Although you can quantitatively describe the acidity of a solution by its [H ], it is often more convenient to give acidity in terms of ph. The ph of a solution is defined as the negative logarithm of the molar hydrogenion concentration. ph = log[h ] 75 78 13

For a solution in which the hydrogen-ion concentration is 1.0 x -3, the ph is: ph = log(1.0 3 = 3.00 Note that the number of decimal places in the ph equals the number of significant figures in the hydrogen-ion concentration. ) Calculate the hydronium and hydroxide ion concentration at 5 C in a. 0. M HCl b. 1.4-4 M Mg(OH) a. When HCl ionizes, it gives H and Cl -. So [H ] = [Cl - ] = [HCl] = 0. M. a. When Mg(OH) ionizes, it gives Mg and OH -. So [OH - ] = [Mg ] = [Mg(OH) ] =.8-4 M. 79 8 The ph of a Solution In a neutral solution, whose hydrogen-ion concentration is 1.0 x -7, the ph = 7.00. For acidic solutions, the hydrogen-ion concentration is greater than 1.0 x -7, so the ph is less than 7.00. Similarly, a basic solution has a ph greater than 7.00. Figure 15.8 shows a diagram of the ph scale and the ph values of some common solutions. A Problem to Consider A sample of orange juice has a hydrogen-ion concentration of.9 x -4 M. What is the ph? ph = log[h ph = log(.9 ph = 3.54 ] 4 ) 80 83 A Problem to Consider Figure 15.8: The ph Scale The ph of human arterial blood is 7.40. What is the hydrogen-ion concentration? [ H ] = anti log( ph) [ H ] = anti log( 7.40) [H ] = = 4.0 7.40 8 M 81 See Exercises 11.8- Problems 74-9 84 14

The ph of a Solution A measurement of the hydroxide ion concentration, similar to ph, is the poh. Then because K w = [H ][OH - ] = 1.0 x -14 at 5 o C, you can show that A has 5 H 3 O and 5 OH -. It is neutral. B has 7 H 3 O and 3 OH -. It is acidic. C has 3 H 3 O and 7 OH -. It is basic. ph poh = 14.00 Listed from most acidic to most basic: B, A, C 85 See Exercise 11.8- and Problems 11.75-9 88 The poh of a Solution A measurement of the hydroxide ion concentration, similar to ph, is the poh. http://www.quia.com/rr/4051.html The poh of a solution is defined as the negative logarithm of the molar hydroxideion concentration. poh = log[oh ] 86 89 The poh of a Solution The ph of a Solution A measurement of the hydroxide ion concentration, similar to ph, is the poh. Then because K w = [H ][OH - ] = 1.0 x -14 at 5 o C, you can show that ph poh = 14.00 The ph of a solution can accurately be measured using a ph meter (see Figure 15.9). Although less precise, acid-base indicators are often used to measure ph because they usually change color within a narrow ph range. Figure 15.8 shows the color changes of various acid-base indicators. 87 90 15

Figure 15.9: A digital ph meter. Photo courtesy of American Color. 91 94 Operational Skills 1 Identify acid and base species. Identify Lewis acids and bases. 3 Decide whether reactants or products are favored in an Acid-base reaction. 4 Calculate Concentrations of H 3 O and OH -. 5 Calculate the ph from the hydronium concentration and vise versa. 9 95 End of This part of Chapter 16 93 96 16

Figure 15.1: Preparation of Sodium Hydroxide by Hydrolysis Problem 15.37 97 0 Problem 15.7 Problem 15.38 98 1 Problem 15.8 Operational Skills Identifying acid and base species Identifying Lewis acid and base species Deciding whether reactants or products are favored in an acid-base reaction Calculating the concentration of H and OH - in solutions of strong acid or base Calculating the ph from the hydrogen-ion concentration, and vice versa 99 17

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