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Chapter 14 Acids and Bases Chapter 14 Table of Contents 14.1 14.2 14.3 14.4 14.5 14.6 14.7 14.8 14.9 14.10 14.11 14.12 The Nature of Acids and Bases Acid Strength The pH Scale Calculating the pH of Strong Acid Solutions Calculating the pH of Weak Acid Solutions Bases Polyprotic Acids Acid–Base Properties of Salts The Effect of Structure on Acid–Base Properties Acid–Base Properties of Oxides The Lewis Acid–Base Model Strategy for Solving Acid–Base Problems: A Summary Copyright © Cengage Learning. All rights reserved 2 Section 14.1 The Nature of Acids and Bases Models of Acids and Bases • Arrhenius: Acids produce H+ ions in solution, bases produce OH- ions. • Brønsted–Lowry: Acids are proton (H+) donors, bases are proton acceptors. HCl + H2O Cl- + H3O+ acid base Return to TOC Copyright © Cengage Learning. All rights reserved 3 Section 14.1 The Nature of Acids and Bases Brønsted–Lowry Reaction Return to TOC Copyright © Cengage Learning. All rights reserved 4 Section 14.1 The Nature of Acids and Bases Acid in Water HA(aq) + acid • • H2O(l) base H3O+(aq) + A-(aq) conjugate conjugate acid base Conjugate base is everything that remains of the acid molecule after a proton is lost. Conjugate acid is formed when the proton is transferred to the base. Return to TOC Copyright © Cengage Learning. All rights reserved 5 Section 14.11 The Lewis Acid–Base Model Lewis Acids and Bases • Lewis acid: electron pair acceptor • Lewis base: electron pair donor 3+ Al3+ + 6 O Al H Lewis acid H H O H 6 Lewis base Return to TOC Copyright © Cengage Learning. All rights reserved 6 Section 14.11 The Lewis Acid–Base Model Three Models for Acids and Bases Return to TOC Copyright © Cengage Learning. All rights reserved 7 Section 14.1 The Nature of Acids and Bases Acid Ionization Equilibrium Return to TOC Copyright © Cengage Learning. All rights reserved 8 Section 14.2 Atomic Acid Strength Masses • Strong acid: Ionization equilibrium lies far to the right. Yields a weak conjugate base. See AP strong acids list! • Weak acid: Ionization equilibrium lies far to the left. Weaker the acid, stronger its conjugate base. Return to TOC Copyright © Cengage Learning. All rights reserved 9 Section 14.2 Atomic Acid Strength Masses Return to TOC Copyright © Cengage Learning. All rights reserved 10 Section 14.2 Atomic Acid Strength Masses Various Ways to Describe Acid Strength Return to TOC Copyright © Cengage Learning. All rights reserved 11 Section 14.2 Atomic Acid Strength Masses Water as an Acid and a Base • Water is amphoteric: Behaves either as an acid or as a base. • At 25°C: Kw = [H+][OH–] = 1.0 × 10–14 • No matter what the solution contains, the product of [H+] and [OH–] must always equal 1.0 × 10–14 at 25°C. Return to TOC Copyright © Cengage Learning. All rights reserved 12 Section 14.2 Atomic Acid Strength Masses Three Possible Situations • [H+] = [OH–]; neutral solution • [H+] > [OH–]; acidic solution • [H+] < [OH–]; basic solution Return to TOC Copyright © Cengage Learning. All rights reserved 13 Section 14.2 Atomic Acid Strength Masses Self-Ionization of Water Return to TOC Copyright © Cengage Learning. All rights reserved 14 Section 14.2 Atomic Acid Strength Masses Concept Check HA(aq) + H2O(l) acid base H3O+(aq) + A-(aq) conjugate conjugate acid base What is the equilibrium constant expression for an acid acting in water? H3O A K = HA Return to TOC Copyright © Cengage Learning. All rights reserved 15 Section 14.3 The pH Mole Scale • pH = –log[H+] • A compact way to represent solution acidity. • pH decreases as [H+] increases. Return to TOC Copyright © Cengage Learning. All rights reserved 16 Section 14.3 The pH Mole Scale pH Range • pH = 7; neutral • pH > 7; basic Higher the pH, more basic. • pH < 7; acidic Lower the pH, more acidic. Return to TOC Copyright © Cengage Learning. All rights reserved 17 Section 14.3 The pH Mole Scale The pH Scale and pH Values of Some Common Substances Return to TOC Copyright © Cengage Learning. All rights reserved 18 Section 14.3 The pH Mole Scale Exercise Calculate the pH for each of the following solutions. a) 1.0 × 10–4 M H+ pH = 4.00 b) 0.040 M OH– pH = 12.60 Return to TOC Copyright © Cengage Learning. All rights reserved 19 Section 14.3 The pH Mole Scale Exercise The pH of a solution is 5.85. What is the [H+] for this solution? [H+] = 1.4 × 10–6 M Return to TOC Copyright © Cengage Learning. All rights reserved 20 Section 14.3 The pH Mole Scale pH and pOH • Recall: Kw = [H+][OH–] –log Kw = –log[H+] – log[OH–] pKw = pH + pOH 14.00 = pH + pOH Return to TOC Copyright © Cengage Learning. All rights reserved 21 Section 14.3 The pH Mole Scale Exercise Calculate the pOH for each of the following solutions. a) 1.0 × 10–4 M H+ pOH = 10.00 b) 0.040 M OH– pOH = 1.40 Return to TOC Copyright © Cengage Learning. All rights reserved 22 Section 14.3 The pH Mole Scale Exercise The pH of a solution is 5.85. What is the [OH–] for this solution? [OH–] = 7.1 × 10–9 M Return to TOC Copyright © Cengage Learning. All rights reserved 23 Section 14.4 Calculating the pH of Strong Acid Solutions Concept Check Calculate the pH of a 1.5 x 10–11 M solution of HCl. pH = 7.00 Return to TOC Copyright © Cengage Learning. All rights reserved 24 Equilibrium Constants for Weak Acids Weak acid has Ka < 1 Leads to small [H3O+] and a pH of 2 - 7 Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 1. Define equilibrium concs. in ICE table. [HOAc] [H3O+] [OAc-] initial 1.00 0 0 change -x +x +x equilib 1.00-x x x Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 2. Write Ka expression + ][OAc- ] 2 [H O x 3 Ka 1.8 x 10-5 = [HOAc] 1.00 - x This is a quadratic. Solve using quadratic formula. or you can make an approximation if x is very small! (Rule of thumb: 10-5 or smaller is ok) Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 3. Solve Ka expression + ][OAc- ] 2 [H O x 3 Ka 1.8 x 10-5 = [HOAc] 1.00 - x First assume x is very small because Ka is so small. Ka 1.8 x 10-5 = x2 1.00 Now we can more easily solve this approximate expression. Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 3. Solve Ka approximate expression Ka 1.8 x 10-5 = x2 1.00 x = [H3O+] = [OAc-] = 4.2 x 10-3 M pH = - log [H3O+] = -log (4.2 x 10-3) = 2.37 Equilibria Involving A Weak Acid Calculate the pH of a 0.0010 M solution of formic acid, HCO2H. HCO2H + H2O ↔ HCO2- + H3O+ Ka = 1.8 x 10-4 Approximate solution [H3O+] = 4.2 x 10-4 M, pH = 3.37 Exact Solution [H3O+] = [HCO2-] = 3.4 x 10-4 M [HCO2H] = 0.0010 - 3.4 x 10-4 = 0.0007 M pH = 3.47 Equilibrium Constants for Weak Bases Weak base has Kb < 1 Leads to small [OH-] and a pH of 12 - 7 Relation of Ka, Kb, [H3O+] and pH Section 14.5 Calculating the pH of Weak Acid Solutions Exercise Calculate the pH of a 0.50 M aqueous solution of the weak acid HF. (Ka = 7.2 x 10–4) Return to TOC Copyright © Cengage Learning. All rights reserved 33 Section 14.5 Calculating the pH of Weak Acid Solutions Steps Toward Solving for pH H3O+(aq) + F–(aq) 0.50 M ~0 ~0 –x +x +x 0.50–x x x HF(aq) + H2O Initial Change Equilibrium Ka = 7.2 x 10–4 pH = 1.72 Return to TOC Copyright © Cengage Learning. All rights reserved 34 Section 14.5 Calculating the pH of Weak Acid Solutions Percent Dissociation (Ionization) Percent dissociation = amount dissociated (mol/L) 100% initial concentration (mol/L) • For a given weak acid, the percent dissociation increases as the acid becomes more dilute. Return to TOC Copyright © Cengage Learning. All rights reserved 35 Section 14.5 Calculating the pH of Weak Acid Solutions Exercise A solution of 8.00 M formic acid (HCHO2) is 0.47% ionized in water. Calculate the Ka value for formic acid. Ka = 1.8 x 10–4 Return to TOC Copyright © Cengage Learning. All rights reserved 36 Section 14.6 Bases Arrhenius: bases produce OH– ions. Brønsted–Lowry: bases are proton acceptors. In a basic solution at 25°C, pH > 7. Ionic compounds containing OH- are generally considered strong bases. See strong bases list! • pOH = –log[OH–] • pH = 14.00 – pOH • • • • Return to TOC Copyright © Cengage Learning. All rights reserved 37 Section 14.6 Bases Concept Check Calculate the pH of a 1.0 x 10–3 M solution of sodium hydroxide. pH = 11.00 Return to TOC Copyright © Cengage Learning. All rights reserved 38 Section 14.6 Bases Concept Check Calculate the pH of a 1.0 x 10–3 M solution of calcium hydroxide. pH = 11.30 Return to TOC Copyright © Cengage Learning. All rights reserved 39 Section 14.6 Bases • Equilibrium expression for weak bases uses Kb. CN–(aq) + H2O(l) Kb = HCN(aq) + OH–(aq) HCN OH CN Return to TOC Copyright © Cengage Learning. All rights reserved 40 Section 14.6 Bases • pH calculations for solutions of weak bases are very similar to those for weak acids. • Kw = [H+][OH–] = 1.0 × 10–14 • pOH = –log[OH–] • pH = 14.00 – pOH Return to TOC Copyright © Cengage Learning. All rights reserved 41 Equilibria Involving A Weak Base You have 0.010 M NH3. Calc. the pH. NH3 + H2O ↔ NH4+ + OHKb = 1.8 x 10-5 Step 1. Define equilibrium concs. in ICE table [NH3] [NH4+] [OH-] initial 0.010 0 0 change -x +x +x equilib 0.010 - x x x Equilibria Involving A Weak Base You have 0.010 M NH3. Calc. the pH. NH3 + H2O NH4+ + OHKb = 1.8 x 10-5 Step 2. Solve the equilibrium expression [NH4+ ][OH- ] x2 -5 Kb 1.8 x 10 = = [NH3 ] 0.010 - x Assume x is small, so x = [OH-] = [NH4+] = 4.2 x 10-4 M and [NH3] = 0.010 - 4.2 x 10-4 ≈ 0.010 M The approximation is valid ! Equilibria Involving A Weak Base You have 0.010 M NH3. Calc. the pH. NH3 + H2O NH4+ + OHKb = 1.8 x 10-5 Step 3. Calculate pH [OH-] = 4.2 x 10-4 M so pOH = - log [OH-] = 3.37 Because pH + pOH = 14, pH = 10.63 Section 14.6 Bases Concept Check Calculate the pH of a 2.0 M solution of ammonia (NH3). (Kb = 1.8 × 10–5) pH = 11.78 Return to TOC Copyright © Cengage Learning. All rights reserved 45 Ionization Constants of Conjugate Acid-Base Pairs HA (aq) A- (aq) + H2O (l) H2O (l) H+ (aq) + A- (aq) OH- (aq) + HA (aq) H+ (aq) + OH- (aq) Ka Kb Kw KaKb = Kw Weak Acid and Its Conjugate Base Kw Ka = Kb Kw Kb = Ka 15.7 Section 14.7 Polyprotic Acids • Acids that can furnish more than one proton. • Always dissociates in a stepwise manner, one proton at a time. • The conjugate base of the first dissociation equilibrium becomes the acid in the second step. • For a typical weak polyprotic acid: Ka1 > Ka2 > Ka3 • For a typical polyprotic acid in water, only the first dissociation step is important to pH. Copyright © Cengage Learning. All rights reserved Return to TOC 47 Section 14.7 Polyprotic Acids Exercise Calculate the pH of a 1.00 M solution of H3PO4. Ka1 = 7.5 x 10-3 Ka2 = 6.2 x 10-8 Ka3 = 4.8 x 10-13 pH = 1.08 Return to TOC Copyright © Cengage Learning. All rights reserved 48 Section 14.7 Polyprotic Acids Sulfuric Acid • AP always treats H2SO4 in this way: – The first H+ is STRONG. Assume 100% ionization. – The second H+ is WEAK. Use the Ka. – Add the H+ Molarities together to find the pH. Return to TOC Copyright © Cengage Learning. All rights reserved 49 Types of Acid/Base Reactions: Summary Section 14.8 Acid–Base Properties of Salts Salts • Ionic compounds. • When dissolved in water, break up into its ions (which can behave as acids or bases). Return to TOC Copyright © Cengage Learning. All rights reserved 51 Section 14.8 Acid–Base Properties of Salts Salts • The salt of a strong acid and a strong base gives a neutral solution. KCl, NaNO3 Return to TOC Copyright © Cengage Learning. All rights reserved 52 Section 14.8 Acid–Base Properties of Salts Salts • A basic solution is formed if the anion of the salt is the conjugate base of a weak acid. NaF, KC2H3O2 Kw = Ka × Kb Use Kb when starting with base. Return to TOC Copyright © Cengage Learning. All rights reserved 53 Section 14.8 Acid–Base Properties of Salts Salts • An acidic solution is formed if the cation of the salt is the conjugate acid of a weak base. NH4Cl Kw = Ka × Kb Use Ka when starting with acid. Return to TOC Copyright © Cengage Learning. All rights reserved 54 Section 14.8 Acid–Base Properties of Salts Cation neutral neutral Anion neutral conjugate base of weak acid neutral Acidic or Basic neutral basic conjugate acidic acid of weak base conjugate conjugate depends acid of base of on Ka & Kb weak base weak acid values Copyright © Cengage Learning. All rights reserved Example NaCl NaF NH4Cl Al2(SO4)3 Return to TOC 55 Section 14.8 Acid–Base Properties of Salts Qualitative Prediction of pH of Salt Solutions (from Weak Parents) Return to TOC Copyright © Cengage Learning. All rights reserved 56 Section 14.8 Acid–Base Properties of Salts Exercise HC2H3O2 HCN Ka = 1.8 x 10-5 Ka = 6.2 x 10-10 Calculate the Kb values for: C2H3O2− and CN− Kb (C2H3O2-) = 5.6 x 10-10 Kb (CN-) = 1.6 x 10-5 Return to TOC Copyright © Cengage Learning. All rights reserved 57 Section 14.8 Acid–Base Properties of Salts Concept Check Arrange the following 1.0 M solutions from lowest to highest pH. HBr NaCN NaCl NaOH NH3 HF NH4Cl HCN Justify your answer. HBr, HF, HCN, NH4Cl, NaCl, NaCN, NH3, NaOH Copyright © Cengage Learning. All rights reserved Return to TOC 58 Section 14.8 Acid–Base Properties of Salts Exercise Calculate the pH of a 0.75 M aqueous solution of NaCN. Ka for HCN is 6.2 x 10–10. Return to TOC Copyright © Cengage Learning. All rights reserved 59 Section 14.8 Acid–Base Properties of Salts Steps Toward Solving for pH CN–(aq) + H2O Initial 0.75 M 0 ~0 –x +x +x 0.75–x x x Change Equilibrium HCN(aq) + OH–(aq) Kb = 1.6 x 10–5 pH = 11.54 Return to TOC Copyright © Cengage Learning. All rights reserved 60 Section 14.9 The Effect of Structure on Acid–Base Properties Models of Acids and Bases • Two factors for acidity in binary compounds: Bond Polarity (high is good) Bond Strength (low is good) Return to TOC Copyright © Cengage Learning. All rights reserved 61 Section 14.9 The Effect of Structure on Acid–Base Properties Bond Strengths and Acid Strengths for Hydrogen Halides Return to TOC Copyright © Cengage Learning. All rights reserved 62 Section 14.9 The Effect of Structure on Acid–Base Properties Oxyacids • Contains the group H–O–X. • For a given series the acid strength increases with an increase in the number of oxygen atoms attached to the central atom. • The greater the ability of X to draw electrons toward itself, the greater the acidity of the molecule. Return to TOC Copyright © Cengage Learning. All rights reserved 63 Section 14.9 The Effect of Structure on Acid–Base Properties Several Series of Oxyacids and Their Ka Values Return to TOC Copyright © Cengage Learning. All rights reserved 64 Section 14.10 Acid–Base Properties of Oxides Oxides • Acidic Oxides (Acid Anhydrides): O-X bond is strong and covalent. SO2, NO2, CO2 • When H-O-X grouping is dissolved in water, the O-X bond will remain intact. It will be the polar and relatively weak H-O bond that will tend to break, releasing a proton. Return to TOC Copyright © Cengage Learning. All rights reserved 65 Section 14.10 Acid–Base Properties of Oxides Oxides • Basic Oxides (Basic Anhydrides): O-X bond is ionic. K2O, CaO • If X has a very low electronegativity, the O-X bond will be ionic and subject to being broken in polar water, producing a basic solution. Return to TOC Copyright © Cengage Learning. All rights reserved 66 Chapter 15 Acid–Base Equilibria Chapter 15 Table of Contents 15.1 15.2 15.3 15.4 15.5 Solutions of Acids or Bases Containing a Common Ion Buffered Solutions Buffering Capacity Titrations and pH Curves Acid–Base Indicators Copyright © Cengage Learning. All rights reserved 68 Section 15.1 Solutions of Acids or Bases Containing a Common Ion Common Ion Effect • Shift in equilibrium position that occurs because of the addition of an ion already involved in the equilibrium reaction. • An application of Le Châtelier’s principle. Return to TOC Copyright © Cengage Learning. All rights reserved 69 Section 15.1 Solutions of Acids or Bases Containing a Common Ion Example HCN(aq) + H2O(l) H3O+(aq) + CN-(aq) • Addition of NaCN will shift the equilibrium to the left because of the addition of CN-, which is already involved in the equilibrium reaction. • A solution of HCN and NaCN is less acidic than a solution of HCN alone. Return to TOC Copyright © Cengage Learning. All rights reserved 70 Section 15.2 Atomic Masses Buffered Solutions Key Points about Buffered Solutions • Buffered Solution – resists a change in pH. • They are weak acids or bases containing a salt with common ion. • After addition of strong acid or base, deal with stoichiometry first, then the equilibrium. Return to TOC Copyright © Cengage Learning. All rights reserved 71 Which of the following are buffer systems? (a) KF/HF (b) KCl/HCl, (c) Na2CO3/NaHCO3 (a) KF is a weak acid and F- is its conjugate base buffer solution (b) HCl is a strong acid not a buffer solution (c) CO32- is a weak base and HCO3- is it conjugate acid buffer solution 16.3 Section 15.2 Atomic Masses Buffered Solutions Adding an Acid to a Buffer Return to TOC Copyright © Cengage Learning. All rights reserved 73 Section 15.2 Atomic Masses Buffered Solutions Buffers Return to TOC Copyright © Cengage Learning. All rights reserved 74 Section 15.2 Atomic Masses Buffered Solutions Solving Problems with Buffered Solutions Return to TOC Copyright © Cengage Learning. All rights reserved 75 Section 15.2 Atomic Masses Buffered Solutions Henderson–Hasselbalch Equation A pH = pK a + log HA • For a particular buffering system (conjugate acid–base pair), all solutions that have the same ratio [A–] / [HA] will have the same pH. Return to TOC Copyright © Cengage Learning. All rights reserved 76 What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK? Mixture of weak acid and conjugate base! HCOOH (aq) Initial (M) Change (M) Equilibrium (M) 0.30 0.00 0.52 -x +x +x 0.30 - x x 0.52 + x Ka for HCOOH = 1.8 x 10 -4 [H+] [HCOO-] Ka = H+ (aq) + HCOO- (aq) x = 1.038 X 10 -4 pH = 3.98 [HCOOH] 16.2 OR…… Use the Henderson-Hasselbach equation Consider mixture of salt NaA and weak acid HA. NaA (s) Na+ (aq) + A- (aq) HA (aq) H+ (aq) + A- (aq) [H+] Ka [HA] = [A-] -log [H+] = -log Ka - log [HA] [A-] -] [A -log [H+] = -log Ka + log [HA] [H+][A-] pK = -log K Ka = a a [HA] Henderson-Hasselbach equation [conjugate base] pH = pKa + log [acid] [A-] pH = pKa + log [HA] 16.2 What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK? Mixture of weak acid and conjugate base! HCOOH (aq) Initial (M) Change (M) Equilibrium (M) Common ion effect 0.30 – x 0.30 0.52 + x 0.52 H+ (aq) + HCOO- (aq) 0.30 0.00 0.52 -x +x +x 0.30 - x x 0.52 + x [HCOO-] pH = pKa + log [HCOOH] [0.52] = 4.01 pH = 3.77 + log [0.30] HCOOH pKa = 3.77 16.2 Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer system. What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the buffer solution? NH3 (aq) + H2O (l) Kb = [NH4+] [OH-] [NH3] Initial Change End NH4+ (aq) + OH- (aq) = 1.8 X 10-5 0.30 -x 0.30 - x 0.36 +x 0 +x 0.36 + x x (.36 + x)(x) 1.8 X 10-5 = 1.8 X 10-5 (.30 – x) 0.36x 0.30 x = 1.5 X 10-5 pOH = 4.82 pH= 9.18 16.3 Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer system. What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the buffer solution? final volume = 80.0 mL + 20.0 mL = 100 mL NH4+ 0.36 M x 0.080 L = 0.029 mol / .1 L = 0.29 M OH- 0.050 x 0.020 L = 0.001 mol / .1 L = 0.01M NH3 0.30 M x 0.080 = 0.024 mol / .1 L = 0.24M start (M) end (M) Ka= 0.01 0.29 NH4+ (aq) + OH- (aq) 0.28 0.0 [H+] [NH3] = 5.6 X 10-10 [NH4+] [H+] 0.25 0.28 0.24 H2O (l) + NH3 (aq) 0.25 [H+] = 6.27 X 10 -10 = 5.6 X 10-10 pH = 9.20 16.3 Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer system. What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the buffer solution? NH4+ (aq) [NH3] pH = pKa + log [NH4+] H+ (aq) + NH3 (aq) pKa = 9.25 [0.30] pH = 9.25 + log = 9.17 [0.36] final volume = 80.0 mL + 20.0 mL = 100 mL start (M) end (M) 0.01 0.29 NH4+ (aq) + OH- (aq) 0.28 0.0 0.24 H2O (l) + NH3 (aq) 0.25 [0.25] pH = 9.25 + log = 9.20 [0.28] 16.3 Section 15.2 Atomic Masses Buffered Solutions Buffered Solution Characteristics • Buffers contain relatively large amounts of weak acid and corresponding conjugate base. • Added H+ reacts to completion with the conjugate base. • Added OH- reacts to completion with the weak acid. Return to TOC Copyright © Cengage Learning. All rights reserved 83 Section 15.2 Atomic Masses Buffered Solutions Buffered Solution Characteristics • The pH in the buffered solution is determined by the ratio of the concentrations of the weak acid and weak base. As long as this ratio remains virtually constant, the pH will remain virtually constant. This will be the case as long as the concentrations of the buffering materials (HA and A– or B and BH+) are large compared with amounts of H+ or OH– added. Return to TOC Copyright © Cengage Learning. All rights reserved 84 Section 15.3 The Mole Capacity Buffering • The amount of protons or hydroxide ions the buffer can absorb without a significant change in pH. • Determined by the magnitudes of [HA] and [A–]. • A buffer with large capacity contains large concentrations of the buffering components. Return to TOC Copyright © Cengage Learning. All rights reserved 85 Section 15.3 The Mole Capacity Buffering • Optimal buffering occurs when [HA] is equal to [A–]. • It is for this condition that the ratio [A–] / [HA] is most resistant to change when H+ or OH– is added to the buffered solution. Return to TOC Copyright © Cengage Learning. All rights reserved 86 Section 15.4 Titrations and pH Curves Titration Curve • Plotting the pH of the solution being analyzed as a function of the amount of titrant added. • Equivalence (Stoichiometric) Point – point in the titration when enough titrant has been added to react exactly with the substance in solution being titrated. Return to TOC Copyright © Cengage Learning. All rights reserved 87 Section 15.4 Titrations and pH Curves Neutralization of a Strong Acid with a Strong Base Return to TOC Copyright © Cengage Learning. All rights reserved 88 Section 15.4 Titrations and pH Curves The pH Curve for the Titration of 50.0 mL of 0.200 M HNO3 with 0.100 M NaOH Return to TOC Copyright © Cengage Learning. All rights reserved 89 Section 15.4 Titrations and pH Curves The pH Curve for the Titration of 100.0 mL of 0.50 M NaOH with 1.0 M HCI Return to TOC Copyright © Cengage Learning. All rights reserved 90 Section 15.4 Titrations and pH Curves Weak Acid–Strong Base Titration Step 1: Step 2: A stoichiometry problem (reaction is assumed to run to completion) then determine remaining species. An equilibrium problem (determine position of weak acid equilibrium and calculate pH). Return to TOC Copyright © Cengage Learning. All rights reserved 91 Section 15.4 Titrations and pH Curves Concept Check Calculate the pH of a solution made by mixing 0.20 mol HC2H3O2 (Ka = 1.8 x 10–5) with 0.030 mol NaOH in 1.0 L of aqueous solution. Return to TOC Copyright © Cengage Learning. All rights reserved 92 Section 15.4 Titrations and pH Curves Let’s Think About It… HC2H3O2(aq) + OH– Before Change After 0.20 mol 0.030 mol –0.030 mol –0.030 mol 0.17 mol 0 C2H3O2–(aq) + H2O 0 +0.030 mol 0.030 mol K = 1.8 x 109 Return to TOC Copyright © Cengage Learning. All rights reserved 93 Section 15.4 Titrations and pH Curves Steps Toward Solving for pH H3O+ + C2H3O2-(aq) 0.170 M ~0 0.030 M –x +x +x 0.170 – x x 0.030 + x HC2H3O2(aq) + H2O Initial Change Equilibrium Ka = 1.8 x 10–5 pH = 3.99 Copyright © Cengage Learning. All rights reserved Return to TOC 94 Section 15.4 Titrations and pH Curves Exercise Calculate the pH of a 100.0 mL solution of 0.100 M acetic acid (HC2H3O2), which has a Ka value of 1.8 x 10–5. pH = 2.87 Return to TOC Copyright © Cengage Learning. All rights reserved 95 Section 15.4 Titrations and pH Curves Concept Check Calculate the pH of a solution made by mixing 100.0 mL of a 0.100 M solution of acetic acid (HC2H3O2), which has a Ka value of 1.8 x 10–5, and 50.0 mL of a 0.10 M NaOH solution. pH = 4.74 Return to TOC Copyright © Cengage Learning. All rights reserved 96 Finding the Equivalence Point (calculation method) • Strong Acid vs. Strong Base – 100 % ionized! pH = 7 No equilibrium! • Weak Acid vs. Strong Base – Acid is neutralized; Need Kb for conjugate base equilibrium • Strong Acid vs. Weak Base – Base is neutralized; Need Ka for conjugate acid equilibrium • Weak Acid vs. Weak Base – Depends on the strength of both; could be conjugate acid, conjugate base, or pH 7 Exactly 100 mL of 0.10 M HNO2 are titrated with 100 mL of a 0.10 M NaOH solution. What is the pH at the equivalence point ? start (moles) 0.01 0.01 HNO2 (aq) + OH- (aq) NO2- (aq) + H2O (l) end (moles) 0.0 0.0 0.01 Final volume = 200 mL NO2- (aq) + H2O (l) Initial (M) Change (M) 0.01 = 0.05 M 0.200 OH- (aq) + HNO2 (aq) [NO2-] = 0.05 0.00 0.00 -x +x +x x x Equilibrium (M) 0.05 - x [OH-][HNO2] x2 -11 = 2.2 x 10 Kb = = [NO2-] 0.05-x pOH = 5.98 0.05 – x 0.05 x 1.05 x 10-6 = [OH-] pH = 14 – pOH = 8.02 Section 15.4 Titrations and pH Curves The pH Curve for the Titration of 50.0 mL of 0.100 M HC2H3O2 with 0.100 M NaOH Return to TOC Copyright © Cengage Learning. All rights reserved 99 Section 15.4 Titrations and pH Curves The pH Curve for the Titration of 100.0mL of 0.050 M NH3 with 0.10 M HCl Return to TOC Copyright © Cengage Learning. All rights reserved 100 Section 15.5 Acid–Base Indicators • Marks the end point of a titration by changing color. • The equivalence point is not necessarily the same as the end point (but they are ideally as close as possible). Return to TOC Copyright © Cengage Learning. All rights reserved 101 Section 15.5 Acid–Base Indicators The Acid and Base Forms of the Indicator Phenolphthalein Return to TOC Copyright © Cengage Learning. All rights reserved 102 Section 15.5 Acid–Base Indicators The Methyl Orange Indicator is Yellow in Basic Solution and Red in Acidic Solution Return to TOC Copyright © Cengage Learning. All rights reserved 103 Section 15.5 Acid–Base Indicators Useful pH Ranges for Several Common Indicators Return to TOC Copyright © Cengage Learning. All rights reserved 104 Chapter 16 Solubility and Complex Ion Equilibria Chapter 16 Table of Contents 16.3 Equilibria Involving Complex Ions Copyright © Cengage Learning. All rights reserved 106 Section 16.3 The Mole Involving Complex Ions Equilibria Complex Ion Equilibria • Charged species consisting of a metal ion surrounded by ligands. Ligand: Lewis base • Formation (stability) constant. Equilibrium constant for each step of the formation of a complex ion by the addition of an individual ligand to a metal ion or complex ion in aqueous solution. Return to TOC Copyright © Cengage Learning. All rights reserved 107 Chapter 21 Transition Metals and Coordination Chemistry Chapter 21 Table of Contents 21.1 21.2 21.3 The Transition Metals: A Survey The First-Row Transition Metals Coordination Compounds Copyright © Cengage Learning. All rights reserved 109 Section 21.1 The Transition Metals: A Survey Forming Ionic Compounds • More than one oxidation state is often found. • Cations are often complex ions – species where the transition metal ion is surrounded by a certain number of ligands (Lewis bases). Return to TOC Copyright © Cengage Learning. All rights reserved 110 Section 21.1 The Transition Metals: A Survey The Complex Ion Co(NH3)63+ Return to TOC Copyright © Cengage Learning. All rights reserved 111 Section 21.1 The Transition Metals: A Survey Ionic Compounds with Transition Metals • Most compounds are colored because the transition metal ion in the complex ion can absorb visible light of specific wavelengths. • Many compounds are paramagnetic. Return to TOC Copyright © Cengage Learning. All rights reserved 112 Section 21.1 The Transition Metals: A Survey Electron Configurations • First-row transition metal ions do not have 4s electrons. Energy of the 3d orbitals is less than that of the 4s orbital. Ti: [Ar]4s23d2 Ti3+: [Ar]3d1 Return to TOC Copyright © Cengage Learning. All rights reserved 113 Section 21.3 The Mole Coordination Compounds Coordination Number • Number of bonds formed between the metal ion and the ligands in the complex ion. 6 and 4 (most common) 2 and 8 (least common) Return to TOC Copyright © Cengage Learning. All rights reserved 114 Section 21.3 The Mole Coordination Compounds Ligands • Neutral molecule or ion having a lone electron pair that can be used to form a bond to a metal ion. Monodentate ligand – one bond to a metal ion Bidentate ligand (chelate) – two bonds to a metal ion Polydentate ligand – more than two bonds to a metal ion Return to TOC Copyright © Cengage Learning. All rights reserved 115 Complex Ion Formation • These are usually formed from a transition metal surrounded by ligands (polar molecules or negative ions). • As a "rule of thumb" you place twice the number of ligands around an ion as the charge on the ion... example: the dark blue Cu(NH3)42+ (ammonia is used as a test for Cu2+ ions), and Ag(NH3)2+. • Memorize the common ligands. Common Ligands Ligands Names used in the ion H2O NH3 aqua ammine OHClBrCNSCN- hydroxy chloro bromo cyano thiocyanato (bonded through sulphur) isothiocyanato (bonded through nitrogen) Names • Names: ligand first, then cation Examples: – tetraamminecopper(II) ion: Cu(NH3)42+ – diamminesilver(I) ion: Ag(NH3)2+. – tetrahydroxyzinc(II) ion: Zn(OH)4 2- • The charge is the sum of the parts (2+) + 4(-1)= -2. When Complexes Form • The odd complex ion, FeSCN2+, shows up once in a while • Acid-base reactions may change NH3 into NH4+ (or vice versa) which will alter its ability to act as a ligand. • Visually, a precipitate may go back into solution as a complex ion is formed. For example, Cu2+ + a little NH4OH will form the light blue precipitate, Cu(OH)2. With excess ammonia, the complex, Cu(NH3)42+, forms. • Keywords such as "excess" and "concentrated" of any solution may indicate complex ions in REACTIONS! AgNO3 + HCl forms the white precipitate, AgCl. With excess, concentrated HCl, the complex ion, AgCl2-, forms and the solution clears. Some Coordination Complexes molecular formula Lewis base/ligand Lewis acid donor atom coordination number Ag(NH3)2+ NH3 Ag+ N 2 [Zn(CN)4]2- CN- Zn2+ C 4 [Ni(CN)4]2- CN- Ni2+ C 4 [PtCl6] 2- Cl- Pt4+ Cl 6 Ni2+ N 6 [Ni(NH3)6]2+ NH3