Aqueous Equilibria: Chemistry of the Water World. Chapter Outline

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1 Aqueous Equilibria: Chemistry of the Water World Chapter Outline 15.1 Acids and Bases: The BrØnsted Lowry Model 15.2 Acid Strength and Molecular Structure 15.3 ph and the Autoionization of Water 15.4 Calculations Involving ph, K a, and K b 15.5 Polyprotic Acids 15.6 ph of Salt Solutions 15.7 The Common-Ion Effect 15.8 ph Buffers 15.9 ph Indicators and Acid Base Titrations Solubility Equilibria 2 1

2 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 Have a bitter taste. Bases Feel slippery. Many soaps contain bases. Nomenclature Review Ch 4, Section 4.2 You are only responsible for nomenclature taught in the lab. These ions are part of many different acids and you need to know them! PO 3-4, HPO 2-4, H 2 PO - 4 SO 2-4, HSO - 4 SO 2-3, HSO - 3 CO 2-3, HCO - 3 NO 3-, NO - 2 S 2-, HS - C 2 H 3 O - 2 (CH 3 COO - ) binary acids, oxoacids H 3 PO 4 H 2 SO 4 H 2 SO 3 H 2 CO 3 HNO 3, HNO 2 H 2 S HC 2 H 3 O 2 HCl, HClO 4 2

3 Strong and Weak Acids A Brønsted acid is a proton donor A Brønsted base is a proton acceptor Strong Acid: Completely ionized HNO 3 (aq) + H 2 O(l) NO 3- (aq) + H 3 O + (aq) (H + donor) (H + acceptor) Weak Acid: Partially ionized HNO 2 (aq) + H 2 O(l) NO 2- (aq) + H 3 O + (aq) (H + donor) (H + acceptor) 3

4 Hydronium Ion Conjugate Acid-Base Pairs 4

5 Weak Acids reordered stronger Kc = CH 3 COO (aq) [H 3 O + (aq)] CH 3 COOH(aq) [H 2 O(l)] 5

6 Strong and Weak Bases Strong and Weak Bases 6

7 Weak Bases 7

8 Relative Strengths of Acids/Bases Leveling Effect: H 3 O + is the strongest H + donor that can exist in water. Strong acids all have the same strength in water; they are completely converted into H 3 O + ions. Relative Strengths of Acids/Bases Leveling Effect Bases: OH - is the strongest H + acceptor that can exist in H 2 O 8

9 Chapter Outline 15.1 Acids and Bases: The BrØnsted Lowry Model 15.2 Acid Strength and Molecular Structure 15.3 ph and the Autoionization of Water 15.4 Calculations Involving ph, K a, and K b 15.5 Polyprotic Acids 15.6 ph of Salt Solutions 15.7 The Common-Ion Effect 15.8 ph Buffers 15.9 ph Indicators and Acid Base Titrations Solubility Equilibria 17 Acid Strength and Molecular Structure H 2 SO 4 is a stronger acid because 1. The -2 charge is delocalized over 4 oxygen atoms compared to three 2. the larger number of oxygens in H 2 SO 4 creates a greater electronegativity effect and consequent weakening of the O-H bond. 9

10 10

11 The Acid-Base Properties of Water Water is amphoteric - which means that it can behave either as an acid or a base H 2 O (l) H + (aq) + OH - (aq) H + H O + H O H H autoionization of water H O H H O H + [ ] + - equivalent expressions base conjugate acid H 2 O + H 2 O H 3 O + + OH - acid conjugate base 11

12 Chapter Outline 15.1 Acids and Bases: The BrØnsted Lowry Model 15.2 Acid Strength and Molecular Structure 15.3 ph and the Autoionization of Water 15.4 Calculations Involving ph, K a, and K b 15.5 Polyprotic Acids 15.6 ph of Salt Solutions 15.7 The Common-Ion Effect 15.8 ph Buffers 15.9 ph Indicators and Acid Base Titrations Solubility Equilibria 23 ph and the Autoionization of Water 2 H 2 O(l) = H 3 O + + OH - K c = [H 3 O + ][OH - ] [H 2 O] 2 What is the concentration of H 3 O + and OH - in pure water? Using the RICE table - 12

13 ph - A Measure of Acidity ph = -log [H + ] ph [H + ] Solution neutral acidic basic [H + ] = [OH - ] [H + ] > [OH - ] [H + ] < [OH - ] [H + ] = 1 x 10-7 [H + ] > 1 x 10-7 [H + ] < 1 x 10-7 ph = 7 ph < 7 ph > 7 The ph Scale 13

14 ph, poh, and K poh is defined the same way as ph - poh = - log[oh - ] p-functions are very common in chemistry, i.e. the negative log of any physical constant is calculated the same way. Since K a and K b values for weak acids and bases tend to be very small, it s convenient to take the negative log of these values as well pk a = - log[k a ] pk b = - log[k b ] The smaller the K a, the weaker the acid The weaker the acid, the larger the pk a The same concepts apply for weak bases 14

15 Useful Equation for Acid-Base Calculations 1. Starting with K w K w = [H 3 O + ][OH - ] = 1.00 x Taking the negative log of both sides - - log K w = - log [H 3 O + ][OH - ] - log(1.00 x ) = - log [H 3 O + ] - log[oh - ] 14 = ph + poh 15

16 Chapter Outline 15.1 Acids and Bases: The BrØnsted Lowry Model 15.2 Acid Strength and Molecular Structure 15.3 ph and the Autoionization of Water 15.4 Calculations Involving ph, K a, and K b 15.5 Polyprotic Acids 15.6 ph of Salt Solutions 15.7 The Common-Ion Effect 15.8 ph Buffers 15.9 ph Indicators and Acid Base Titrations Solubility Equilibria 31 Weak Acids Most acids are weak. How do you know if an acid is weak? Because it s not one of the 6 strong ones you ve memorized! HCl HBr HI hydrochloric hydrobromic hydroiodic HNO 3 HClO 4 H 2 SO 4 nitric perchloric sulfuric The monster in the movie Alien had blood that contained molecular acid and ate through six decks of the spaceship! 16

17 General Weak Acid Equilibrium Equation and K a HA(aq) + H 2 O(l) = H 3 O + (aq) + A - (aq) K a = [H 3 O + ][A - ] [HA] If you measure the ph of a solution containing a weak acid, you can calculate the equilibrium constant Calculating K a for a weak acid when the ph is known, e.g M, ph = 2.20 HA(aq) + H 2 O(l) = H 3 O + (aq) + A - (aq) M 17

18 Percent Ionization Percent Ionization = [H+ ] equil [HA] initial X 100% 18

19 Weak Bases Weak bases frequently contain nitrogen because the lone pair makes a good proton acceptor NH 3 ammonia C 6 H 5 NH 2 aniline NH(CH 3 ) 2 dimethylamine General Weak Base Equilibrium Equation and K b B(aq) + H 2 O(l) = BH + (aq) + OH - (aq) K b = [BH + ][OH - ] [B] Ordinary bleach contains the weak base ClO - 19

20 Relationship Between K a and K b K a = [H 3 O + ][A - ] [HA] K b = [HA][OH - ] [A - ] 20

21 Chapter Outline 15.1 Acids and Bases: The BrØnsted Lowry Model 15.2 Acid Strength and Molecular Structure 15.3 ph and the Autoionization of Water 15.4 Calculations Involving ph, K a, and K b 15.5 Polyprotic Acids 15.6 ph of Salt Solutions 15.7 The Common-Ion Effect 15.8 ph Buffers 15.9 ph Indicators and Acid Base Titrations Solubility Equilibria 41 Polyprotic Acids Two or more ionizable protons K a1 > K a2 > K a3 H H + K a1 = 7.11 x 10-3 H H + K a2 = 6.32 x 10-8 H H + K a3 = 4.5 x

22 Chapter Outline 15.1 Acids and Bases: The BrØnsted Lowry Model 15.2 Acid Strength and Molecular Structure 15.3 ph and the Autoionization of Water 15.4 Calculations Involving ph, K a, and K b 15.5 Polyprotic Acids 15.6 ph of Salt Solutions 15.7 The Common-Ion Effect 15.8 ph Buffers 15.9 ph Indicators and Acid Base Titrations Solubility Equilibria 44 22

23 ph of Salt Solutions 1. Neutral Salts (ph = 7) are from strong electrolytes (100% ionization) (a)ionic Compounds: NaCl(aq) Na + (aq) + Cl - (aq) base conj. acid HCl(aq) + H 2 O(l) H 3 O + (aq) + Cl - (aq) acid conj. base Infinitely strong Infinitely weak ph of Salt Solutions 2. Basic Salts (ph > 7) are conjugate bases of weak acids HClO(aq) + H 2 O(l) = H 3 O + (aq) + ClO - (aq) weak conj. acid base ClO - (aq) + H 2 O(l) = OH - (aq) + HClO(aq) conj. base ph > 7 K a (HClO) = 2.9 x 10-8 K b = K b (ClO - ) = 3.4 x 10-7 K w K a K b = 1.0 x x

24 ph of Salt Solutions 3. Acidic Salts (ph < 7) are conjugate acids of weak bases NH 3 (aq) + H 2 O(l) = OH - (aq) + NH 4+ (aq) weak conj. base acid NH 4+ (aq) + H 2 O(l) = H 3 O + (aq) + NH 3 (aq) conj. acid ph < 7 K b (NH 3 ) = 1.8 x 10-5 K a = K a (NH 4+ ) = 5.6 x K w K b K a = 1.0 x x

25 Calculating the ph of Solutions of Weak Acids and Bases: Use the RICE Table as Before 1. Calculating the ph of a Solution of a Basic Salt ClO - (aq) + H 2 O(l) = OH - (aq) + HClO(aq) M - x K b = 3.4 x 10-7 = x M - x 3.4 x 10-7 = x M x = (3.4 x 10-7 )(0.100) x Since K a is < 10-5, assume that x << M x K a = 2.9 x 10-8 K b = K w K a K b = 3.4 x 10-7 x = [OH - ] = 1.9 x 10-4 M Assumption OK poh = - log (1.9 x 10-4 ) = 3.7 ph = = 10.3 Calculating the ph of Solutions of Weak Acids and Bases: Use the RICE Table as Before 2. Calculating the ph of a Solution of an Acidic Salt (Ex. 15.8) What is the ph of a 0.25 M solution of NH 4 Cl? 25

26 Chapter Outline 15.1 Acids and Bases: The BrØnsted Lowry Model 15.2 Acid Strength and Molecular Structure 15.3 ph and the Autoionization of Water 15.4 Calculations Involving ph, K a, and K b 15.5 Polyprotic Acids 15.6 ph of Salt Solutions 15.7 The Common-Ion Effect 15.8 ph Buffers 15.9 ph Indicators and Acid Base Titrations Solubility Equilibria 51 The Common Ion Effect A shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. CH 3 COONa (s) CH 3 COOH (aq) Na + (aq) + CH 3 COO - (aq) H + (aq) + CH 3 COO - (aq) common ion The common ion effect can be used to produce a BUFFER SOLUTION = a solution of a weak acid or base and it's conjugate, e.g. CH 3 COOH and CH 3 COONa By controlling the ratio of weak acid/base to it's conjugate, we can shift the equilibrium to whatever [H + ] and therefore ph we want. 26

27 The Henderson-Hasselbach Equation HA(aq) + H 2 O(aq) = H 3 O + (aq) + A - (aq) K a = [H 3 O + ][A - ] [HA] 1. Solution using the RICE table 27

28 2. Solution using the Henderson-Hasselbach equation Chapter Outline 15.1 Acids and Bases: The BrØnsted Lowry Model 15.2 Acid Strength and Molecular Structure 15.3 ph and the Autoionization of Water 15.4 Calculations Involving ph, K a, and K b 15.5 Polyprotic Acids 15.6 ph of Salt Solutions 15.7 The Common-Ion Effect 15.8 ph Buffers 15.9 ph Indicators and Acid Base Titrations Solubility Equilibria 56 28

29 ph Buffers Calculate the response of buffers to an influx of acid or base as compared to the same amount in pure water, p.686. [H 2 CO 3 ] = 1.3 x 10-5 M and [HCO 3- ] = 1.0 x 10-4 M 1.0 L samples of river water and pure water, and then add 10.0 ml of 1.0 x 10-3 HNO 3 (a) Calculate the ph change in pure water ph Buffers (b) Calculate the ph change in the buffer 29

30 A Buffer in Action Weak Acid and its Salt e.g. CH 3 COO - /CH 3 COOH Weak Base and its Salt e.g. NH 4+ /NH 3 Add H + Add OH - 30

31 Chapter Outline 15.1 Acids and Bases: The BrØnsted Lowry Model 15.2 Acid Strength and Molecular Structure 15.3 ph and the Autoionization of Water 15.4 Calculations Involving ph, K a, and K b 15.5 Polyprotic Acids 15.6 ph of Salt Solutions 15.7 The Common-Ion Effect 15.8 ph Buffers 15.9 ph Indicators and Acid Base Titrations Solubility Equilibria 61 ph Indicators 31

32 Acid-Base Titrations In a titration a solution of accurately known concentration (titrant) is added gradually added to another solution of unknown concentration (analyte) until the chemical reaction between the two solutions is complete. Equivalence point the point at which the reaction is complete Indicator substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL The indicator changes color (pink) Acid-Base Titrations (p. 692): Strong Acid NaOH + HCl NaCl + H 2 O titrant analyte 20.0 ml M each ph = 7 (a) Initial ph = - log (0.100 M) = 1.00 (b) ph eq. pt. = 7 32

33 Acid-Base Titrations (p. 692): Weak Acid NaOH + CH 3 COOH NaCH 3 COO + H 2 O titrant analyte 20.0 ml M each (a) Initial ph = use RICE table (b) ph eq. pt. < 7, RICE table again (c) ph midpoint = Henderson- Hasselbach eqn. Titration Calculations: Concentration of the Unknown aa + bb products acid M A, V A unknown base M B, V B known Remember: M x V = moles 33

34 34

35 Chapter Outline 15.1 Acids and Bases: The BrØnsted Lowry Model 15.2 Acid Strength and Molecular Structure 15.3 ph and the Autoionization of Water 15.4 Calculations Involving ph, K a, and K b 15.5 Polyprotic Acids 15.6 ph of Salt Solutions 15.7 The Common-Ion Effect 15.8 ph Buffers 15.9 ph Indicators and Acid Base Titrations Solubility Equilibria 69 Solubility Equilibria Calculating the concentration of sparingly soluble salts in solution (compounds that violated the solubility rules do ionize to a very limited extent) 35

36 Solubility Equilibria Solubility Equilibria 36

37 Solubility Equilibria and the Solubility Product K sp Another equilibrium constant that allows calculation of the amount of a compound that will dissolve in water. AgCl(s) Ca 3 (PO 4 ) 2 (s) Calculating the Molar Solubility from K sp Mg(OH) 2 (s) = Mg 2+ (aq) + 2 OH - (aq) K sp = 5.6 x

38 General Equation for Calculating the Molar Solubility A m B n (s) = m A n+ (aq) + n B m- (aq) Common Ion Effect BaSO 4 (s) = Ba 2+ (aq) + SO 4 2- (aq) K sp = 9.1 x

39 Need to include the hydrolysis of the conjugate acid or base K sp and Q Q > K sp forms too much product so it precipitates Q = K sp at equilibrium Q < K sp not enough products so no precipitate forms 39

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