Acids and Bases - the Three Definitions

1. Measurement of pH - the pH meter Bronsted-Lowry definition of acids and bases - an acid is a proton donor - a base is a proton acceptor - conjugate acid/conjugate base pairs - relationship of Ka of a conjugate acid and Kb of a conjugate base Lewis definition of acids and bases - a base is an electron pair donor - an acid is an electron pair acceptor - some examples of Lewis acids and Lewis bases Acids and Bases - the Three Definitions

2. Ionization Constants (1)

3. Ionization Constants (2)

4. Exercise on Acid/Bases Strength For each conjugate acid/base pair, (1) Write the reactions defining Ka and Kb. Find the values of pKa, pKb, and Kb. Which species is the strongest conjugate acid, which is the strongest conjugate base? nitrous acid: HNO2 / NO2- oxalic acid (2): HC2O4- / C2O42- arsenic acid (2): HAsO42- / AsO43- carbonic acid (1): H2CO3 / HCO3-

5. Hydrolysis Reactions Which salts undergo hydrolysis? Is the resulting solution acidic, basic, or neutral? Write the hydrolysis reaction (if any). Calculate the pH of a 0.10 M solution. sodium acetate (basic (pH=8.88), acetate (pKb=9.25) hydrolyses to produce OH-) ammonium chloride (acidic (pH=5.12), ammonium (pKa=9.25) hydrolyses to produce H3O+) calcium chloride (neutral, no hydrolysis) sodium monohydrogen phosphate (basic (pH=10.12), HPO42- (pKb2=6.79) hydrolyses to produce OH-) (you need to consider two conjugate acid/base pairs..)

6. x2 0.1 - x pH and % Dissociation of a Monoprotic Weak Acid What is the pH of 0.10 M CH3COOH? CH3COOH = CH3COO- + H+ Ka = 1.75 x 10-5 0.1 - x x x (We let x = [H+]) [CH3COO-] [H+] Ka = = = 1.75 x 10-5 [CH3COOH] (x << 0.1) Approximation Method: Since Ka <<1, assume x<<0.1 x2 = 0.1 * 1.75 * 10-5 = 1.75 x 10-6 and x = [H+] = 0.0013 M

7. Calculating % Dissociation and the pH CH3COOH = CH3COO- + H+ [H+] = 1.3 x 10-3 M pH = - log10[H+] = 2.88 [CH3COO-] x 0.1 % dissociation = 100 * = 1.3% = [CH3COOH]init

8. Measurement of pH: the pH Meter pH varies linearly with output voltage and can be measured over the range pH 0 to pH 14

9. HNO2 (aq) = NO2- (aq) + H+ (aq) Ka = 4.6 x 10-4 pKa = 3.34 Ka and Acid Strength The stronger the acid, the larger the Ka and the smaller the pKa: stronger CH3COOH (aq) = CH3COO- (aq) + H+ (aq) Ka = 1.76 x 10-5 pKa = 4.75 HCN (aq) = CN- (aq) + H+ (aq) Ka = 6.17 x 10-10 pKa = 9.21 weaker

10. = = = = = = = Weak Acids stronger weaker

11. [NH4+] [OH-] Kb = [NH4OH] Kb and pKb Arrhenius bases liberate OH- in solution. Kb is the equilibrium constant for this reaction. NH4OH (aq) = NH4+ (aq) + OH- (aq) = 1.76 x 10-5 pKb = - log10 Kb (definition) pKb = - log10 (1.8 x 10-5) = 4.74

12. NH4OH (aq) = NH4+ (aq) + OH- (aq) Kb = 1.8 x 10-5 PO43- (aq) + H2O (l) = HPO42- (aq) + OH- (aq) Kb = 4.5 x 10-2 pKb = 4.74 pKb = 1.34 Kb and Base Strength The stronger the base, the larger the Kb and the smaller the pKb: stronger Conclusion: phosphate anion is a stronger base than NH4OH. weaker

13. Kb's of weak bases Strength (Ranked) 3 7 6 1 5 2 4

14. Acids and Bases - Three Definitions Arrhenius Definition: Acids: increases [H+] in aqueous solution Bases: increases [OH-] in aqueous solution Bronsted-Lowry Definition: (based on proton transfer reactions) Acids: proton (H+) donor Bases: proton (H+) acceptor Lewis Definition: Acids: electron pair acceptor Bases: electron pair donor

15. H .. .. .. .. O-H- : N-H .. : : : O-H 2- I : .. :S : .. H H H+ Zn2+ Hg2+ Ag+ BF3 metal cations electron deficient compounds Some Lewis Acids and Bases Lewis bases are characterized by having an available lone pair. Examples are: hydroxide iodide ammonia water sulfide Lewis acids are electron deficient - i.e., electron pair acceptors Examples are:

16. Least general definitionMost general definition ArrheniusBronsted-Lowry Lewis Lewis Acids/Bases - the Most General Definition The Lewis definition generalizes the acid/base concept: Every Arrhenius acid/base is also a Lewis acid/base. Every Bronsted acid/base is also a Lewis acid/base. Example: A strong acid reacts with a strong base: H+(aq) + :OH-(aq) = H2O(l) Electron pair acceptorelectron pair donor

17. Lewis definition: Acid: electron pair acceptorBase: electron pair donor HNO2+ ClO2- = HClO2 + NO2- Bronsted-Lowry: acid1 base2 acid2 base1 The Lewis acid is H+ (the electron deficient species) There are 2 bases (NO2- and ClO2-), which compete for the acid The lone pairs donated by these bases are on oxygen atoms: .. .. .. : O=N- O H .. .. .. .. .. .. .. .. .. .. : .. O=Cl- O H .. .. .. : O=N- O .. .. .. .. .. .. .. : .. O-N= O - .. - .. : : ] [ O-Cl= O O=Cl- O ] [ .. .. .. .. .. .. The Lewis Acid-Base Reaction + +

18. .. .. H+ + :O-H- = H-O Lewis acid Lewis base : .. H H F H F : F - B + N - H = F - B :N - H F H F H Lewis acid Lewis base Lewis Acids and Bases The acid/base concept is further generalized by the Lewis acid/base definition. The driving force is the donation of an electron pair to electron-deficient atom. Lewis acid - an electron pair acceptor Lewis base - an electron pair donor

19. Complex Ions in Solution One example of Lewis acid-base neutralization involves the stepwise complexes formed between Hg(II) and I- ion. There are four stepwise reactions, each of which is an acid/base neutralization by the Lewis definition: Hg2+(aq) + I-(aq) = HgI+(aq) HgI+(aq) + I-(aq) = HgI2(s) (red-brown ppt) HgI2(s) + I-(aq) = HgI3-(aq) HgI3-(aq) + I-(aq) = HgI42-(aq) Identify the Lewis acid and Lewis base in each reaction.

20. Al3+ + :OH- =AlOH2+ AlOH2+ + :OH-=Al(OH)2+ Al(OH)2+ + :OH-=Al(OH)3 (s) white precipitate Al(OH)3 + :OH-=Al(OH)4- Complex Ions and Solubility Stepwise Lewis acid/base complexes form between Al3+(aq) and OH- ion. All charged species are soluble in aqueous solution. Only the uncharged Al(OH)3(s) forms a white precipitate. There are four stepwise reactions, each of which is an acid/base neutralization by the Lewis definition: Identify the Lewis acid and Lewis base in each reaction.