1. Chapter 16 Acid-Base Equilibria

2. Dissociation of water Autoionization or autoprotolysis Ion-product constant Autoprotolysis constant constant

3. K w = [H + ][OH - ] = 1.0x10 -14 When [H + ] = [OH - ] neutral. Doesn’t usually happen. As one increases, the other decreases; the product must equal 1.0x10 -14. When [H + ] > [OH - ]acidic [OH - ] > [H + ]basic

4. H + is a proton with no electrons. In water:        H HOH Hydronium ion

5. Bronstead-Lowry Acid-Base Acid - Can donate a proton Base - Can accept a proton *Doesn’t have to be in H 2 O. Can be in other solvents.

6. Conjugate Acid-Base Pairs conj base conj acid

7. The stronger an acid, the weaker its conjugate base. The weaker an acid, the stronger its conjugate base.

8. pH scale pH = -log [H + ] Remember K w = (1x10 -7 )(1x10 -7 ) = 1.0x10 -14 pH = -log [H + ] = -log (1x10 -7 ) pH = 7 (neutral) [H + ]pH acidic> 1.0x10 -7 < 7.00 basic 7.00

9. You can also speak in terms of [OH - ] pOH = -log [OH - ] = 14 - pH Because pH + pOH = -log K w = 14

10. Measure pH by pH meter Acid-base indicators Litmus red = pH < 5 blue = pH > 8 Figure 16.7 shows several acid-base indicators and their ranges

11. Strong Acids and Bases Strong electrolytes Completely ionize HA + H 2 O  A - + H 3 O + Bases form hydroxides in solvent

12. In H 2 O, Alkali metal hydroxides Alkaline earth metal Hydroxides (except Be) Many are insoluble Also, substances that will abstract a H + from H 2 O. O 2- + H 2 O  2OH - Na 2 O or CaO would do this. O 2-, H -, N 3- bases that would do this.

13. Weak acids Only partially ionize Acid dissociation constant

14. Larger K a means stronger acid. ex. N O C - O - H O = 0.020M solution pH = 3.26 ? K a pH = -log [H + ] = 3.26 [H + ] = 5.50x10 -4

15. N O C - OH O = N O C - O O = + H +  HA A-A- H+H+ 1:1

16. Can calculate pH in same manner if you have K a and concentration of solution. Let’s use niacin again. N O C - OH O = N O C - O O = + H +  HA A-A- H+H+

17. ** Simplifying Assumption ** x is very very small compared to 0.010M sooooooooo, ignore x in denominator

18. pH = -log [H + ] x = [H + ] = 3.9x10 -4 pH = 3.41 What percent of niacin molecules ionized?

19. Polyprotic Acids ex. H 2 SO 4 H 3 PO 4 H 2 SeO 4 H 2 SO 4  H + + HSO 4 - K a1 = 1.7x10 -2 HSO 4 -  H + + SO 4 2- K a2 = 6.4x10 -8 K a1 always larger than K a2 If K a1 / K a2  10 3, can estimate pH by K a1 only.

20. Weak Bases ex. Amines “an organic substituted ammonia” ammonia NH 3 N H HHN H CH 3 H methyl amine

21. N H CH 3 + H 2 O  HN H CH 3 + OH - H H ClO - + H 2 O  HClO + OH - K b = 3.3x10 -7 Can use this in the same manner in which you used K a. Anions of weak acids

22. K a and K b How are they related?

23. 1) 2) 3) When two reactions are added together, the equilibrium constant for the third reaction is given by the product of equilibrium constants of equations 1 and 2.

24. K 1 x K 2 = K 3 rxn 1 rxn 2 rxn 3

25. Special Case K a x K b = K w For conjugate acid-base pairs.

26. Bond polarity and Bond strength effect on Acid-base behavior: In binary acids  polarity(across a row)  acidity  bond strength(in a group)  acidity  stability of conj. base  acidity

27. Metal hydrides are basic or show no acid/base properties in H 2 O. Nonmetal hydrides are acidic or show no acid/base properties in H 2 O (except NH 3 ) Acidity increases moving down a group.

28. Oxyacids HOS O O O H Have unprotonated and protonated oxygens. YOH H 3 PO 4 As electronegativity of Y increases, acidity increases. As number of unprotonated oxygens increases, acidity increases (effect of formal charge and oxidation number) Ex. HClO, HClO 2, HClO 3, HClO 4

29. Carboxylic Acids R C OH OCOOH = Carboxyl group R = H or an organic group. The more electron withdrawing R is, the greater the acidity (this stabilizes anion and weakens O-H bond) ex. CHC H H O O H Acetic acid K a = 1.8x10 -5 CFC F F O O H Trifluoroacetic acid K a = 5.0x10 -1

30. Lewis Acids and Bases This is a completely different definition for acid/base chemistry than what you have seen thus far!!! Lewis acid = electron pair acceptor Lewis base = electron pair ‘donor’ Not giving them away, just has them available to ‘share’.

31. H + Bronstead-Lowry acid also a Lewis acid H + electron pair acceptor OH - Electron pair donor Lewis base also Bronstead-Lowry base

32. B H H H BH 3 not a Bronstead- Lowry acid, but it’s a Lewis acid Incomplete Octet N H H H Lewis Base has an electron pair available to attack an area that is e - deficient

33. Transition metal ions are often Lewis Acids. They have vacant d orbitals. (s and p also) H H O O = C = OCan be a Lewis Acid because e - density around the C is bound in just 2 directions.

34. H H O = = O O C H H O O O C H H O O O C Carbonic acid Hydrolysis of metal ions Metal ions have positive charge so they attract the lone e - pair on H 2 O molecules

35. 6 of these H O H Fe 3+ H O H H O H H O H H O H H O H H O H Fe 3+ Because the metal is (+), e - density of H 2 O moves toward the metal. When this happens, there is less e - density in water’s O-H bonds, so H + can come off easier…  pH will drop.

36. The higher the charge density of the metal ion, the greater the acidity of its aqua complex.