Acids and Bases Chapter 16

Homework pages 532 – 535 1-7, 9 – 14, 16, 17, 20 - 34, 37, 40, 42, 44, 45, 47, 52 – 56, 58, 59, 61 - 63, 66, 68, 71, 72 74, 82.

Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Bases Have a bitter taste. Feel slippery. Many soaps contain bases. 4.3

Arrhenius acid is a substance that produces H+ (H3O+) in water Arrhenius base is a substance that produces OH- in water 4.3

A Brønsted acid is a proton donor A Brønsted base is a proton acceptor conjugate acid conjugate base base acid 15.1

The conjugate base of an acid is the species formed when one proton is removed. The conjugate acid of a base is the species formed when one proton is added. A conjugate acid-base pair differ by a proton. acid + base conjugate base + conjugate acid What is conjugate base of: HF, H2SO4, NH3? What is conjugate acid of: O2-, SO42-, NH3?

The Conjugate Pairs in Some Acid-Base Reactions Acid + Base Base + Acid Conjugate Pair Reaction 1 HF + H2O F– + H3O+ Reaction 2 HCOOH + CN– HCOO– + HCN Reaction 3 NH4+ + CO32– NH3 + HCO3– Reaction 4 H2PO4– + OH– HPO42– + H2O Reaction 5 H2SO4 + N2H5+ HSO4– + N2H62+ Reaction 6 HPO42– + SO32– PO43– + HSO3–

Identifying Conjugate Acid-Base Pairs Problem: The following chemical reactions are important for industrial processes. Identify the conjugate acid-base pairs. HSO4-(aq) + CN-(aq) SO42-(aq) + HCN(aq) (b) ClO-(aq) + H2O(l) HClO(aq) + OH-(aq) (c) S2-(aq) + H2O(aq) HS-(aq) + OH-(aq) (d) HS-(aq) + H2O(aq) H2S(aq) + OH-(aq)

[ ] - Acid-Base Properties of Water autoionization of water H2O (l) H+ (aq) + OH- (aq) autoionization of water O H + - [ ] conjugate acid base H2O + H2O H3O+ + OH- conjugate base acid 15.2

The Ion Product of Water Kc = [H+][OH-] [H2O] H2O (l) H+ (aq) + OH- (aq) [H2O] = constant Kc[H2O] = Kw = [H+][OH-] The ion-product constant (Kw) is the product of the molar concentrations of H+ and OH- ions at a particular temperature. Solution Is [H+] = [OH-] neutral At 250C Kw = [H+][OH-] = 1.0 x 10-14 [H+] > [OH-] acidic [H+] < [OH-] basic 15.2

What is the [H+] in 0.035M NaOH? What is the concentration of OH- ions in a HCl solution whose hydrogen ion concentration is 1.3 M? Kw = [H+][OH-] = 1.0 x 10-14 [H+] = 1.3 M [OH-] = Kw [H+] 1 x 10-14 1.3 = = 7.7 x 10-15 M What is the [H+] in 0.035M NaOH? 15.2

pH – A Measure of Acidity pH = -log [H+] Solution Is At 250C neutral [H+] = [OH-] [H+] = 1 x 10-7 pH = 7 acidic [H+] > [OH-] [H+] > 1 x 10-7 pH < 7 basic [H+] < [OH-] [H+] < 1 x 10-7 pH > 7 pH [H+] 15.3

pOH = -log [OH-] [H+][OH-] = Kw = 1.0 x 10-14 -log [H+] – log [OH-] = 14.00 pH + pOH = 14.00 15.3

The pH of rainwater collected in a certain region of the northeastern United States on a particular day was 4.82. What is the H+ ion concentration of the rainwater? pH = -log [H+] [H+] = 10-pH = 10-4.82 = 1.5 x 10-5 M The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood? pH + pOH = 14.00 pOH = -log [OH-] = -log (2.5 x 10-7) = 6.60 pH = 14.00 – pOH = 14.00 – 6.60 = 7.40 15.3

Calculating [H3O+], pH, [OH-], and pOH Problem: A chemist dilutes concentrated hydrochloric acid to make two solutions: (a) 3.0 M and (b) 0.0024 M. Calculate the [H3O+], pH, [OH-], and pOH of the two solutions at 25°C. What is the [H3O+], [OH-], and pOH of a solution with pH = 3.67? With pH = 8.05?

Strong Electrolyte – 100% dissociation NaCl (s) Na+ (aq) + Cl- (aq) H2O Weak Electrolyte – not completely dissociated CH3COOH CH3COO- (aq) + H+ (aq) Strong Acids are strong electrolytes HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq) HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq) HClO4 (aq) + H2O (l) H3O+ (aq) + ClO4- (aq) H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4- (aq) 15.4

Weak Acids are weak electrolytes HF (aq) + H2O (l) H3O+ (aq) + F- (aq) HNO2 (aq) + H2O (l) H3O+ (aq) + NO2- (aq) HSO4- (aq) + H2O (l) H3O+ (aq) + SO42- (aq) H2O (l) + H2O (l) H3O+ (aq) + OH- (aq) Strong Bases are strong electrolytes NaOH (s) Na+ (aq) + OH- (aq) H2O KOH (s) K+ (aq) + OH- (aq) H2O Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq) H2O 15.4

Weak Bases are weak electrolytes (NH3) F- (aq) + H2O (l) OH- (aq) + HF (aq) NO2- (aq) + H2O (l) OH- (aq) + HNO2 (aq) Conjugate acid-base pairs: The conjugate base of a strong acid has no measurable strength. H3O+ is the strongest acid that can exist in aqueous solution. The OH- ion is the strongest base that can exist in aqueous solution. 15.4

15.4

Strong Acid Weak Acid 15.4

What is the pH of a 2 x 10-3 M HNO3 solution? HNO3 is a strong acid – 100% dissociation. Start 0.002 M 0.0 M 0.0 M HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq) End 0.0 M 0.002 M 0.002 M pH = -log [H+] = -log [H3O+] = -log(0.002) = 2.7 What is the pH of a 1.8 x 10-2 M Ba(OH)2 solution? Ba(OH)2 is a strong base – 100% dissociation. Start 0.018 M 0.0 M 0.0 M Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq) End 0.0 M 0.018 M 0.036 M pH = 14.00 – pOH = 14.00 + log(0.036) = 12.56 15.4

Weak Acids (HA) and Acid Ionization Constants HA (aq) + H2O (l) H3O+ (aq) + A- (aq) HA (aq) H+ (aq) + A- (aq) Ka = [H+][A-] [HA] Ka is the acid ionization constant weak acid strength Ka 15.5

15.5

What is the pH of a 0.5 M HF solution (at 250C)? Ka = [H+][F-] [HF] = 7.1 x 10-4 HF (aq) H+ (aq) + F- (aq) HF (aq) H+ (aq) + F- (aq) Initial (M) 0.50 0.00 0.00 Change (M) -x +x +x Equilibrium (M) 0.50 - x x x Ka = x2 0.50 - x = 7.1 x 10-4 Ka << 1 0.50 – x  0.50 Ka  x2 0.50 = 7.1 x 10-4 x2 = 3.55 x 10-4 x = 0.019 M [H+] = [F-] = 0.019 M pH = -log [H+] = 1.72 [HF] = 0.50 – x = 0.48 M 15.5

When can I use the approximation? Ka << 1 0.50 – x  0.50 When x is less than 5% of the value from which it is subtracted. 0.019 M 0.50 M x 100% = 3.8% Less than 5% Approximation ok. x = 0.019 What is the pH of a 0.05 M HF solution (at 250C)? Ka  x2 0.05 = 7.1 x 10-4 x = 0.006 M 0.006 M 0.05 M x 100% = 12% More than 5% Approximation not ok. Must solve for x exactly using quadratic equation or method of successive approximation. 15.5

EH Assignment Due Mar 7 Unit 18 Minimum score Sec 1 Acid-Base Reactions 90 Sec 2 Kw and pH Sec 3 pH meter Sec 4 Strong Acids and Bases Sec 5 Weak Acids and Bases

Finding the Ka of a Weak Acid from the pH of its Solution and Vice Versa Problem: The weak acid hypochlorous acid is formed in bleach solutions. If the pH of a 0.12 M solution of HClO is 4.19, what is the value of the Ka of this weak acid? What is the % ionization? Ka = 3.5 x 10-8

Solving weak acid ionization problems: Identify the major species that can affect the pH. In most cases, you can ignore the autoionization of water. Ignore [OH-] because it is determined by [H+]. Use ICE to express the equilibrium concentrations in terms of single unknown x. Write Ka in terms of equilibrium concentrations. Solve for x by the approximation method. If approximation is not valid, solve for x exactly. Calculate concentrations of all species and/or pH of the solution. 15.5

What is the pH of a 0.122 M monoprotic acid whose Ka is 5.7 x 10-4? HA (aq) H+ (aq) + A- (aq) Initial (M) 0.122 0.00 0.00 Change (M) -x +x +x Equilibrium (M) 0.122 - x x x Ka = x2 0.122 - x = 5.7 x 10-4 Ka << 1 0.122 – x  0.122 Ka  x2 0.122 = 5.7 x 10-4 x2 = 6.95 x 10-5 x = 0.0083 M 0.0083 M 0.122 M x 100% = 6.8% More than 5% Approximation not ok. 15.5

Ka = x2 0.122 - x = 5.7 x 10-4 x2 + 0.00057x – 6.95 x 10-5 = 0 -b ± b2 – 4ac  2a x = ax2 + bx + c =0 x = 0.0081 x = - 0.0081 HA (aq) H+ (aq) + A- (aq) Initial (M) Change (M) Equilibrium (M) 0.122 0.00 -x +x 0.122 - x x [H+] = x = 0.0081 M pH = -log[H+] = 2.09 15.5

For a monoprotic acid HA Ionized acid concentration at equilibrium Initial concentration of acid x 100% percent ionization = For a monoprotic acid HA Percent ionization = [H+] [HA]0 x 100% [HA]0 = initial concentration 15.5

The Stepwise Dissociation of Phosphoric Acid Phosphoric acid is a weak acid, and normally only loses one proton in solution, but it will lose all three when reacted with a strong base with heat. The ionization constants are given for comparison. H3PO4 (aq) + H2O(l) H2PO4-(aq) + H3O+(aq) H2PO4-(aq) + H2O(l) HPO42-(aq) + H3O+(aq) HPO42-(aq) + H2O(l) PO43-(aq) + H3O+(aq) H3PO4 (aq) + 3 H2O(l) PO43-(aq) + 3 H3O+(aq)

15.8

Terms to know: diprotic and polyprotic Calculate the concentration of all species in 0.10 M H2S.

Weak Bases and Base Ionization Constants NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq) Kb = [NH4+][OH-] [NH3] Kb is the base ionization constant weak base strength Kb Solve weak base problems like weak acids except solve for [OH-] instead of [H+]. 15.6

Most anions can act as weak bases. The base-dissociation constant Kb refers to the equilibrium that occurs when a weak base is added to water. Ammonia and amines are the most common molecules that act as weak bases. Most anions can act as weak bases.

15.6

Determining pH from Kb and Initial [B] Problem: Ammonia is commonly used cleaning agent in households and is a weak base, with a Kb of 1.8 x 10-5. What is the pH of a 1.5 M NH3 solution?

Ionization Constants of Conjugate Acid-Base Pairs HA (aq) H+ (aq) + A- (aq) Ka A- (aq) + H2O (l) OH- (aq) + HA (aq) Kb H2O (l) H+ (aq) + OH- (aq) Kw KaKb = Kw Weak Acid and Its Conjugate Base Ka = Kw Kb Kb = Kw Ka 15.7

All anions can act as weak bases except those that are the conjugate bases of the strong acids.

Determining the pH of a Salt Solution Problem: Sodium cyanide in water can produce a basic solution. What is the pH of a 0.25 M solution of NaCN? Ka for HCN is 4.9 x 10-10 Plan: We have to find the pH of a solution of the cyanide ion, CN -, which acts as a base in water:

The stronger an acid, the weaker is its conjugate base and vice versa. An acid-base reaction proceeds to the greater extent in the direction to form the weaker acid and base. If the weaker acid and base are on the right Kc > 1 Sample Problem H2PO4-(aq) + NH3(aq) NH4+(aq) + HPO42-(aq) H2O(l) + HS-(aq) OH-(aq) + H2S(aq)

The reaction of an ion with water is frequently called hydrolysis. Salt solutions are usually acidic or basic because of hydrolysis of the cation or anion. Cations: all cations act as acids except those in groups 1 and 2. Anions: A few are acids HSO4- Those that are conjugate bases of strong acids are neutral. (Cl-, NO3-, …) All others are bases.

Acid-Base Properties of Salts Acid Solutions: It is easy to see how a salt like NH4Cl can produce an acidic solution. NH4Cl (s) NH4+ (aq) + Cl- (aq) H2O NH4+ (aq) NH3 (aq) + H+ (aq) How can salts Al(NO3)3, CrCl3, or FeBr3 produce acidic solutions? Al(H2O)6 (aq) Al(OH)(H2O)5 (aq) + H+ (aq) 3+ 2+ 15.10

Acid Hydrolysis of Al3+ 15.10

Problem 16.62 Which of these salts will undergo hydrolysis? KF, NaNO3, NH4I, KCN, C6H5COONa, RbI, Na2CO3, CaCl2, HCOOK, AlCl3. Are the solutions acidic, basic or neutral? Write hydrolysis equation.

Definition of An Acid Arrhenius acid is a substance that produces H+ (H3O+) in water A Brønsted acid is a proton donor A Lewis acid is a substance that can accept a pair of electrons A Lewis base is a substance that can donate a pair of electrons + OH- • H O H • H+ acid base N H • H N H H + H+ + acid base 15.12

No protons donated or accepted! Lewis Acids and Bases F B F N H • H F B F N H H + acid base No protons donated or accepted! 15.12

Problem 16.74 Describe this equation according to the Lewis theory of acids and bases. AlCl3 + Cl-  AlCl4-