1. Acid-Base chemistry Acidity of blood (pH range of Heartburn (acid-reflux) – Tums, Rolaids, Milk of Magnesia; The Purple Pill , Nexium Acidity regulation (tropical fish / goldfish tanks) Pepsi (most sodas); just how acidic? –Loosens rusty bolts; cleans windshields Battery acid (H 2 SO 4 ) Acid Rain (SO 2, NO 2, CO 2 ) Drain cleaners (Drano, Liquid Plumber)

2. Nature of acids and bases Acids – sour, tart taste (vinegar, lemon juice) Bases – bitter taste, slippery feel between fingers (drano, detergent) Arrhenius definitions (1884): Acids: produce H + ions (protons) in solution Bases: produce OH - (hydroxide) ions in solution

3. Lowry-Brønsted definitions Danish & British chemists More general definitions than Arrhenius definition: Acids: Proton donors Bases: proton acceptors H 2 O + HCl  H 3 O + + Cl -

4. Lewis Acids / Lewis Bases LA: electron pair acceptor LB: electron pair donor NH 3 (aq) + H + (aq)  NH 4 + (aq) LB LA OH - (aq) + H + (aq)  H 2 O(l)

5. General reaction for HA + H 2 O Position of equilibrium dictated between bases in equation (competition for H + ) H 2 O stronger base: A - stronger base: K eq = [H 3 O + ][A - ]K a = Acid dissociation constant [HA][H 2 O][H 3 O + ] = [H + ] in H 2 O Ka = [H 3 O + ][A - ]K a = Acid dissociation constant [HA] HA(aq) + H 2 O(l)  H 3 O + (aq) + A - (aq) Acid BaseConjugate Conjugate Acid Base

6. Acid dissociation reactions Dissociation of acid (HA) most important here HA (aq)  H + (aq) + A - (aq) H 2 O still important (required for aqueous conditions) K a for this equilibrium process? Can predict dissociation reaction for any acid (no matter how complex-looking) HCl (aq) HC 2 H 3 O 2 (aq) NH 4 + (aq) C 6 H 5 NH 3 + (aq)

7. Acid Strength Defined by equilibrium position of dissociation reaction: HA(aq) + H 2 O(l)  H 3 O + (aq) + A - (aq) Strong acid – lies far to right side (a) Weak acid – lies far to left side (b)

8. Describing acid strength Insert Figure 14.6 Insert table 14.1 strong acid weak acid PropertyStrong Acid Weak Acid K a value Equilibrium Posn. [H + ] e vs [HA] o Strength of conj. base compared to water

9. Common strong acids HCl, HClO 4, HNO 3, HI, HBr (monoprotic acids) H 2 SO 4 (diprotic acid) K a values – very large

10. Relative base strengths Arrange F -, Cl -, NO 2 -, CN - in order of increasing base strength:

11. Water: Amphoteric substance Can exist as both an acid and a base Autoionization of water H 2 O(l) + H 2 O(l)  H 3 O + (aq) + OH - (aq) K w = [H 3 O + ][OH - ]; K w = Dissociation constant for H 2 O K w = [H + ][OH - ] [H + ]=[OH - ] = 1.0 x 10 -7 M (at 25 °C, in pure water) K w = [H + ][OH - ] = 1.0 x 10 -14 In any aqueous solution, the product of [H + ] and [OH - ] must always equal 1.0 x 10 -14

12. KwKw K w = [H + ][OH - ] = 1.0 x 10 -14 3 possible situations: Neutral solutions; [H + ] = [OH - ] Acidic solutions; [H + ] > [OH - ] Basic solutions; [H + ] < [OH - ] In all cases: K w = [H + ][OH - ] = 1.0 x 10 -14 Varies with T (as do all K values)

13. Calculating [H + ],[OH - ] [H + ] with [OH - ] = 1.0 x 10 -5 M [OH - ] with [H + ] = 10.0 M [OH - ] and [H + ] in neutral solution at 60 °C (K w = 1x10 -13 at 60 °C)

14. pH scale (acidity measurement) pH = -log [H + ] [H + ] = 1.0 x 10 -7 M; pH = pOH = -log[OH - ]; pK = -log K a Log scale; pH changes by 1 for every power of 10 change in [H + ] pH decreases as [H + ] increases

15. Calculating pH / pOH 1.0 x 10 -3 M OH - 1.0 M H + pH + pOH = 14.00 pH of sample of human blood = 7.41 @ 25 °C. Calculate pOH, [H + ], [OH - ]

16. pH of strong acid solutions What species are present? 1M HCl What are the major species that can furnish H + ions? pH of 0.10 M HNO 3 pH of 1.0 x 10 -10 M HCl

17. pH of weak acid solutions pH of 1.00 M HF (weak acid); K a = 7.2 x 10 -4 Major species in solution? HF, H 2 O Major species which furnish H + ? Which of these two is the stronger acid? K a = [H + ][F - ] / [HF] = 7.2 x 10 -4 In order to calculate pH, need equilibrium value of [H + ] [HF] o = 1.00 M [H + ] o = 0 M (approximation, as H + from H 2 0 not included here) [F - ] o = 0 M Let x be the change required to reach equilibrium…..

18. pH of weak acid solutions Equilibrium concentration of [HF] = Equilibrium concentration of [H + ] = Equilibrium concentration of [F - ] = Substitute these into K a = [H + ][F - ] / [HF] = 7.2 x 10 -4 eqn, and solve for x….

19. Bases Cleaning solutions (ammonia, bleach) Antacids (Tums, Rolaids, Milk of Magnesia) Arrhenius: produces OH - ions Lowry – Brønsted: H + acceptor Strong Bases - NaOH

20. pH of strongly basic solution 5.0 x 10 -2 M NaOH solution (same procedure as for acidic pH calculations); Expected pH range? Major species:

21. Weak bases Many types of bases don’t contain OH -, but do increase [OH - ] when dissolved in water (through reaction with water). Base Acid Lone pair of electrons on N picks up H + from H 2 O

22. Weak bases General reaction with H 2 O: K b = K b always refers to reaction of a base with H 2 O to produce a conjugate acid and OH -

23. pH of weak bases - calculations Very similar to those of weak acids pH of 15.0 M solution of NH 3 (K b = 1.8 x 10 -5 )

24. Polyprotic acids More than 1 acidic H; H 2 SO 4, H 3 PO 4 ; H 2 CO 3 Consider H 3 PO 4 Only 1 st dissociation step usually important for [H + ] determination K a1 >>> K a2 >>> K a3

25. Polyprotic acids pH of 5.0 M H 3 PO 4 solution; eq. concs. of H 3 PO 4, H 2 PO 4 2-, HPO 4 3- and PO 4 3-.

26. Sulfuric acid (H 2 SO 4 ) Unique acid Strong acid in 1 st dissociation step Weak acid in 2 nd dissociation step pH of 1.0 M H 2 SO 4 solution Does HSO 4 - make significant contribution to [H + ]? No.

27. Acid-base properties of salts Salts producing neutral solutions Salts producing basic solutions Salts producing acidic solutions

28. Acid-base properties of salts

29. Structure effects on acid/base properties Hydrogen halides Oxyacids