1. A CIDS AND B ASES II IB C HEMISTRY G R.12 Topic 18 1 Chem2_Dr. Dura

2. T OPICS  18.1 Calculations Involving Acids and Bases  18.2 Buffer Solutions  18.3 Salt Hydrolysis  18.4 Acid-Base Titrations  18.5 Indicators 2 Chem2_Dr. Dura

3. IB S TANDARDS  18.1.1 State the expression for the ionic product constant of water (K w ).  18.1.2 Deduce [H + (aq)] and [OH - (aq)], for water at different temperatures given K w values.  18.1.3 Solve problems involving [H + (aq)], [OH - (aq)], pH and pOH.  18.1.4 State the equation for the reaction for any weak acid or weak base with water, and hence deduce the expressions for K a and K b. Chem 2_Dr. Dura 3

4. IB S TANDARDS 18.1.5 Solve problems involving solutions of weak acids and bases using the expressions: K a x K b = K w pK a + pK b = pK w pH + pOH = pK w 18.1.6 Identify the relative strengths of acids and bases using values of K a, K b, pK a and pK b. 4 Chem2_Dr. Dura

5. 5 H 2 O (l) H + (aq) + OH - (aq) The Ion Product of Water K c = [H + ][OH - ] [H 2 O] [H 2 O] = constant K c [H 2 O] = K w = [H + ][OH - ] The ion-product constant (K w ) is the product of the molar concentrations of H + and OH - ions at a particular temperature. At 25 0 C K w = [H + ][OH - ] = 1.0 x 10 - 14 [H + ] = [OH - ] [H + ] > [OH - ] [H + ] < [OH - ] Solution Is neutral acidic basic

6. K W IS TEMPERATURE DEPENDENT Temp o CKwKw [H+] in pure water pH 01.5 x 10 -15 3.9 x 10 -6 7.47 206.8 x 10 -15 8.2 x10 -6 7.08 251.0 x 10 -14 1.0 x10 -7 7.00 6 Chem2_Dr. Dura

7. 7 pH = - log [H+], [H + ] = 10 -pH pOH = -log [OH-], [OH - ] = 10 -pOH [H + ] = [OH - ] [H + ] > [OH - ] [H + ] < [OH - ] Solution is neutral acidic basic [H + ] = 1 x 10 -7 [H + ] > 1 x 10 -7 [H + ] < 1 x 10 -7 pH = 7 pH < 7 pH > 7 At 25 0 C pH[H + ]

8. Chem2_Dr. Dura 8 [H + ][OH - ] = K w = 1.0 x 10 -14 -log [H + ] – log [OH - ] = 14.00 pH + pOH = 14.00 At 25 0 C

9. Chem2_Dr. Dura 9 Example 1: 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 Example 2: 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

10. Chem2_Dr. Dura 10 HA (aq) + H 2 O (l) H 3 O + (aq) + A - (aq) Weak Acids (HA) and Acid Ionization Constants HA (aq) H + (aq) + A - (aq) K a = [H + ][A - ] [HA] K a is the acid ionization constant KaKa weak acid strength

11. Chem2_Dr. Dura 11 NH 3 (aq) + H 2 O (l) NH 4 + (aq) + OH - (aq) Weak Bases and Base Ionization Constants K b = [NH 4 + ][OH - ] [NH 3 ] K b is the base ionization constant KbKb weak base strength Solve weak base problems like weak acids except solve for [OH-] instead of [H + ].

12. I ONIZATION C ONSTANTS OF C ONJUGATE A CID - B ASE P AIRS  Consider a weak acid (HA) and its conjugate base (A - ) in water: HA (aq)  H + (aq) + A - (aq)  K a = [H + ][A - ] / [HA] A - (aq) + H 2 O(l)  HA(aq) + OH - (aq)  K b = [HA][OH - ] /[A - ]  K a x K b = [H + ][OH - ] = K w Chem 2_Dr. Dura 12

13. IB S TANDARDS 18.2.1 Describe the composition of a buffer solution and explain its action. 18.2.2 Solve problems involving the composition and pH of a specified buffer system. 18.3 Deduce whether salts form acidic, alkaline or neutral aqueous solutions. 13 Chem2_Dr. Dura

14. B UFFER S OLUTIONS 14 Chem2_Dr. Dura A buffer solution is a solution of: 1. A weak acid or a weak base and 2. The salt of the weak acid or weak base Both must be present!

15. Chem2_Dr. Dura 15 The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. The presence of a common ion suppresses the ionization of a weak acid or a weak base. Example: A buffer consisting of CH 3 COONa (salt) and CH 3 COOH (weak acid). CH 3 COONa (s) Na + (aq) + CH 3 COO - (aq) CH 3 COOH (aq) H + (aq) + CH 3 COO - (aq) common ion

16. 16 A buffer solution has the ability to resist changes in pH upon the addition of small amounts of either acid or base. Add strong acid H + (aq) + CH 3 COO - (aq) CH 3 COOH (aq) Add strong base OH - (aq) + CH 3 COOH (aq) CH 3 COO - (aq) + H 2 O (l) How buffers work Consider an equal molar mixture of CH 3 COOH and CH 3 COONa Chem2_Dr. Dura

17. 17 Consider mixture of salt NaA and weak acid HA. HA (aq) H + (aq) + A - (aq) NaA (s) Na + (aq) + A - (aq) K a = [H + ][A - ] [HA] [H + ] = K a [HA] [A - ] -log [H + ] = -log K a - log [HA] [A - ] -log [H + ] = -log K a + log [A - ] [HA] pH = pK a + log [A - ] [HA] pK a = -log K a

18. H ENDERSON -H ASSELBALCH EQUATION 18 pH = pK a + log [conjugate base] [acid] Chem2_Dr. Dura pH of buffers can be determined from Henderson- Hasselbalch Equation.

19. 19 Exercise: Which of the following are buffer systems? (a) KF/HF (b) KBr/HBr, (c) Na 2 CO 3 /NaHCO 3 (a) KF is a weak acid and F - is its conjugate base buffer solution (b) HBr is a strong acid not a buffer solution (c) CO 3 2- is a weak base and HCO 3 - is its conjugate acid buffer solution Chem2_Dr. Dura

20. 20 = 9.20 Problem: Calculate the pH of the 0.30 M NH 3 /0.36 M NH 4 Cl 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? NH 4 + (aq) H + (aq) + NH 3 (aq) pH = pK a + log [NH 3 ] [NH 4 + ] pK a = 9.25 pH = 9.25 + log [0.30] [0.36] = 9.17 NH 4 + (aq) + OH - (aq) H 2 O (l) + NH 3 (aq) start (moles) end (moles) 0.0290.001 0.024 0.0280.00.025 pH = 9.25 + log [0.25] [0.28] [NH 4 + ] = 0.028 0.10 final volume = 80.0 mL + 20.0 mL = 100 mL [NH 3 ] = 0.025 0.10 Chem2_Dr. Dura

21. 21 Acid-Base Properties of Salts Neutral Solutions: Salts containing an alkali metal or alkaline earth metal ion (except Be 2+ ) and the conjugate base of a strong acid (e.g. Cl -, Br -, and NO 3 - ). NaCl (s) Na + (aq) + Cl - (aq) H2OH2O Basic Solutions: Salts derived from a strong base and a weak acid. NaCH 3 COOH (s) Na + (aq) + CH 3 COO - (aq) H2OH2O CH 3 COO - (aq) + H 2 O (l) CH 3 COOH (aq) + OH - (aq)

22. Chem2_Dr. Dura 22 Acid-Base Properties of Salts Acid Solutions: Salts derived from a strong acid and a weak base. NH 4 Cl (s) NH 4 + (aq) + Cl - (aq) H2OH2O NH 4 + (aq) NH 3 (aq) + H + (aq) Salts with small, highly charged metal cations (e.g. Al 3+, Cr 3+, and Be 2+ ) and the conjugate base of a strong acid. Al(H 2 O) 6 (aq) Al(OH)(H 2 O) 5 (aq) + H + (aq) 3+2+

23. Chem2_Dr. Dura 23 Acid-Base Properties of Salts Solutions in which both the cation and the anion hydrolyze : K b for the anion > K a for the cation, solution will be basic K b for the anion < K a for the cation, solution will be acidic K b for the anion  K a for the cation, solution will be neutral

24. IB S TANDARDS  18.4.1Sketch the general shapes of graphs of pH and volume for titrations of strong and weak acids and bases and explain their important features.  18.5.1 Describe qualitatively the action of acid- base indicator.  18.5.2 State and explain how the pH range of an acid-base indicator relates to its pKa value.  18.5.3 Identify an appropriate indicator for a titration given the equivalence point of titration and the pH range of the indicator. 24 Chem2_Dr. Dura

25. 25 Titrations In a titration a solution of accurately known concentration is added gradually to another solution of unknown concentration 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) Chem2_Dr. Dura

26. 26 Strong Acid-Strong Base Titrations NaOH (aq) + HCl (aq) H 2 O (l) + NaCl (aq) OH - (aq) + H + (aq) H 2 O (l) Chem2_Dr. Dura

27. 27 Weak Acid-Strong Base Titrations CH 3 COOH (aq) + NaOH (aq) CH 3 COONa (aq) + H 2 O (l) CH 3 COOH (aq) + OH - (aq) CH 3 COO - (aq) + H 2 O (l) CH 3 COO - (aq) + H 2 O (l) OH - (aq) + CH 3 COOH (aq) At equivalence point (pH > 7): Chem2_Dr. Dura

28. 28 Strong Acid-Weak Base Titrations HCl (aq) + NH 3 (aq) NH 4 Cl (aq) NH 4 + (aq) + H 2 O (l) NH 3 (aq) + H + (aq) At equivalence point (pH < 7): H + (aq) + NH 3 (aq) NH 4 Cl (aq) Chem2_Dr. Dura

29. 29 Acid-Base Indicators HIn (aq) H + (aq) + In - (aq)  10 [HIn] [In - ] Color of acid (HIn) predominates  10 [HIn] [In - ] Color of conjugate base (In - ) predominates Chem2_Dr. Dura

30. 30 The titration curve of a strong acid with a strong base. Chem2_Dr. Dura

31. 31 Which indicator(s) would you use for a titration of HNO 2 with KOH ? Weak acid titrated with strong base. At equivalence point, will have conjugate base of weak acid. At equivalence point, pH > 7 Use cresol red or phenolphthalein Chem2_Dr. Dura

32. 32 Acid-Base Properties of Salts Acid Solutions: Salts derived from a strong acid and a weak base. NH 4 Cl (s) NH 4 + (aq) + Cl - (aq) H2OH2O NH 4 + (aq) NH 3 (aq) + H + (aq) Salts with small, highly charged metal cations (e.g. Al 3+, Cr 3+, and Be 2+ ) and the conjugate base of a strong acid. Al(H 2 O) 6 (aq) Al(OH)(H 2 O) 5 (aq) + H + (aq) 3+2+ Chem2_Dr. Dura

33. 33 Acid-Base Properties of Salts Solutions in which both the cation and the anion hydrolyze: K b for the anion > K a for the cation, solution will be basic K b for the anion < K a for the cation, solution will be acidic K b for the anion  K a for the cation, solution will be neutral Chem2_Dr. Dura

34. 34 Chemistry In Action: Antacids and the Stomach pH Balance NaHCO 3 (aq) + HCl (aq) NaCl (aq) + H 2 O (l) + CO 2 (g) Mg(OH) 2 (s) + 2HCl (aq) MgCl 2 (aq) + 2H 2 O (l) Chem2_Dr. Dura