1	Chapter 16 Acids and Bases
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3	Some Definitions Arrhenius acids and bases Acid: Substance that, when dissolved in water, increases the concentration of hydrogen ions (protons, H⁺). Base: Substance that, when dissolved in water, increases the concentration of hydroxide ions.
4	Some Definitions Brønsted–Lowry: must have both 1. an Acid: Proton donor and 2. a Base: Proton acceptor
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6	"A Brønsted–Lowry acid" A Brønsted–Lowry acid… …must have a removable (acidic) proton. HCl, H₂O, H₂SO₄ A Brønsted–Lowry base… …must have a pair of nonbonding electrons. NH₃, H₂O
7	If it can be either… ...it is amphiprotic. HCO₃^– HSO₄^– H₂O
8	What Happens When an Acid Dissolves in Water? Water acts as a Brønsted–Lowry base and abstracts a proton (H⁺) from the acid. As a result, the conjugate base of the acid and a hydronium ion are formed.
9	Conjugate Acids and Bases: From the Latin word conjugare, meaning “to join together.” Reactions between acids and bases always yield their conjugate bases and acids.
10	Acid and Base Strength Strong acids are completely dissociated in water. Their conjugate bases are quite weak. Weak acids only dissociate partially in water. Their conjugate bases are strong bases.
11	Acid and Base Strength Substances with negligible acidity do not dissociate in water. Their conjugate bases are exceedingly strong.
12	Conjugate Acid-Base Pair HA(aq) + H₂O(l) ⇄ H₃O⁺(aq) + A^-(aq) A B CA CB 2 substances that are connected by a donating and accepting of H₂O really a “fight” over H+ between two bases: A- and H₂O if A- has a greater affinity, lies to left if H₂O has greater affinity, lies to right
13	Acid and Base Strength In any acid-base reaction, the equilibrium favors the reaction that moves the proton to the stronger base.
14	Acid and Base Strength
15	Acid dissociation constant equilibrium expression where H+ is removed to form conjugate base so for: HA + H₂O ⇄ H₃O⁺ + A^-
16	Strength determined by equilibrium position of dissociation reaction strong acid: lies far to right, almost all HA is dissociated large K_a values creates weak conjugate base weak acid: lies far to left, almost all HA is stays as HA small K_a values creates strong conjugate base
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18	Types of Acids diprotic: an acid with 2 acidic protons oxyacids: ex. HNO₃ most acids are this type where the acidic proton is attached to an oxygen organic acids: ex. CH₃COOH containing a carboxyl group usually weak hydrohalic acids: ex. HBr hydrogen and a halogen
19	Autoionization of Water 2 water molecules exchange H+ H₂O + H₂O ⇄ H₃O⁺ +OH^- K_w= [H₃O⁺][OH^-]= [H⁺][OH^-] dissociation constant for water at 25C, [H⁺]=[OH^-]=1.0 x 10^-7 so at 25C, K_w=1.0x10^-14 how will change it with temperature?
20	Autoionization of Water As we have seen, water is amphoteric. In pure water, a few molecules act as bases and a few act as acids. This process is called autoionization.
21	Ion-Product Constant The equilibrium expression for this process is K_c = [H₃O⁺] [OH^–] This special equilibrium constant is referred to as the ion-product constant for water, K_w. At 25°C, K_w = 1.0 ´ 10^-¹⁴
22	pH pH is defined as the negative base-10 logarithm of the hydronium ion concentration. pH = –log [H₃O⁺]
23	pH In pure water, K_w = [H₃O⁺] [OH^–] = 1.0 ´ 10^-¹⁴ Because in pure water [H₃O⁺] = [OH^-], [H₃O⁺] = (1.0 ´ 10^-¹⁴)^1/2 = 1.0 ´ 10^-⁷
24	pH Therefore, in pure water, pH = –log [H₃O⁺] = –log (1.0 ´ 10^-⁷) = 7.00 An acid has a higher [H₃O⁺] than pure water, so its pH is <7 A base has a lower [H₃O⁺] than pure water, so its pH is >7.
25	pH These are the pH values for several common substances.
26	Other “p” Scales The “p” in pH tells us to take the negative log of the quantity (in this case, hydronium ions). Some similar examples are pOH –log [OH^-] pK_w –log K_w
27	Watch This! Because [H₃O⁺] [OH⁻] = K_w = 1.0 ´ 10^-¹⁴, we know that –log [H₃O⁺] + – log [OH⁻] = – log K_w = 14.00 or, in other words, pH + pOH = pK_w = 14.00
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29	How Do We Measure pH? Litmus paper “Red” paper turns blue above ~pH = 8 “Blue” paper turns red below ~pH = 5 An indicator Compound that changes color in solution.
30	How Do We Measure pH? pH meters measure the voltage in the solution
31	Strong Acids You will recall that the seven strong acids are HCl, HBr, HI, HNO₃, H₂SO₄, HClO₃, and HClO₄. These are strong electrolytes and exist totally as ions in aqueous solution. For the monoprotic strong acids, [H₃O⁺] = [acid].
32	Strong Bases Strong bases are the soluble hydroxides, which are the alkali metal (NaOH, KOH)and heavier alkaline earth metal hydroxides (Ca(OH)₂, Sr(OH)₂, and Ba(OH)₂). Again, these substances dissociate completely in aqueous solution. [OH^-] = [hydroxide added].
33	Dissociation Constants For a generalized acid dissociation, the equilibrium expression is This equilibrium constant is called the acid-dissociation constant, K_a.
34	Dissociation Constants The greater the value of K_a, the stronger the acid.
35	Calculating K_a from the pH The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate K_a for formic acid at this temperature. We know that
36	Calculating K_a from the pH The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate K_a for formic acid at this temperature. To calculate K_a, we need all equilibrium concentrations. We can find [H₃O⁺], which is the same as [HCOO⁻], from the pH.
37	Calculating K_a from the pH pH = –log [H₃O⁺] – 2.38 = log [H₃O⁺] 10^-^2.38 = 10^{log [H}³^O⁺^] = [H₃O⁺] 4.2 ´ 10^-³ = [H₃O⁺] = [HCOO^–]
38	Calculating K_a from pH
39	Calculating K_a from pH
40	Calculating Percent Ionization In the example: [A^-]_eq = [H₃O⁺]_eq = 4.2 ´ 10^-³ M [A^-]_eq + [HCOOH]_eq = [HCOOH]_initial = 0.10 M
41	Calculating Percent Ionization Percent Ionization = ´ 100
42	Calculating pH from K_a Calculate the pH of a 0.30 M solution of acetic acid, C₂H₃O₂H, at 25°C. K_a for acetic acid at 25°C is 1.8 ´ 10^-⁵. Is acetic acid more or less ionized than formic acid (K_a=1.8 x 10^-4)?
43	Calculating pH from K_a The equilibrium constant expression is:
44	Calculating pH from K_a
45	Calculating pH from K_a
46	Calculating pH from K_a
47	Calculating pH from K_a Now,
48	Calculating pH from K_a pH = –log [H₃O⁺] pH = – log (2.3 ´ 10⁻³) pH = 2.64
49	Polyprotic Acids Have more than one acidic proton. If the difference between the K_a for the first dissociation and subsequent K_a values is 10³ or more, the pH generally depends only on the first dissociation.
50	Weak Bases Bases react with water to produce hydroxide ion.
51	Weak Bases The equilibrium constant expression for this reaction is
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53	Weak Bases K_b can be used to find [OH^–] and, through it, pH.
54	pH of Basic Solutions What is the pH of a 0.15 M solution of NH₃?
55	pH of Basic Solutions
56	pH of Basic Solutions (1.8 ´ 10^-⁵) (0.15) = x² 2.7 ´ 10^-⁶ = x² 1.6 ´ 10^-³ = x²
57	pH of Basic Solutions Therefore, [OH^–] = 1.6 ´ 10^-³ M pOH = –log (1.6 ´ 10^-³) pOH = 2.80 pH = 14.00 – 2.80 pH = 11.20
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60	K_a and K_b K_a and K_b are related in this way: K_a ´ K_b = K_w Therefore, if you know one of them, you can calculate the other.
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62	Reactions of Anions with Water Anions are bases. As such, they can react with water in a hydrolysis reaction to form OH^– and the conjugate acid:
63	Reactions of Cations with Water Cations with acidic protons (like NH₄⁺) lower the pH of a solution by releasing H⁺. Most metal cations (like Al³⁺) that are hydrated in solution also lower the pH of the solution; they act by associating with H₂O and making it release H⁺.
64	Reactions of Cations with Water Attraction between nonbonding electrons on oxygen and the metal causes a shift of the electron density in water. This makes the O-H bond more polar and the water more acidic. Greater charge and smaller size make a cation more acidic.
65	Effect of Cations and Anions An anion that is the conjugate base of a strong acid will not affect the pH. An anion that is the conjugate base of a weak acid will increase the pH. A cation that is the conjugate acid of a weak base will decrease the pH.
66	Effect of Cations and Anions Cations of the strong Arrhenius bases will not affect the pH. Other metal ions will cause a decrease in pH. When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the affect on pH depends on the K_aand K_b values.
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68	Factors Affecting Acid Strength The more polar the H-X bond and/or the weaker the H-X bond, the more acidic the compound. Acidity increases from left to right across a row and from top to bottom down a group.
69	Factors Affecting Acid Strength In oxyacids, in which an OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid.
70	Factors Affecting Acid Strength For a series of oxyacids, acidity increases with the number of oxygens.
71	Factors Affecting Acid Strength Resonance in the conjugate bases of carboxylic acids stabilizes the base and makes the conjugate acid more acidic.
72	Lewis Acids Lewis acids are defined as electron-pair acceptors. Atoms with an empty valence orbital can be Lewis acids. A compound with no H’s can be a Lewis acid.
73	Lewis Bases Lewis bases are defined as electron-pair donors. Anything that is a Brønsted–Lowry base is also a Lewis base. (B-L bases also have a lone pair.) Lewis bases can interact with things other than protons.
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