Acids and Bases - the Three Definitions 1.Measurement of pH - the pH meter 2.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 K a of a conjugate acid and K b of a conjugate base 3.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

Ionization Constants (1)

Ionization Constants (2)

Exercise on Acid/Bases Strength For each conjugate acid/base pair, (1) Write the reactions defining K a and K b. (2)Find the values of pK a, pK b, and K b. (3)Which species is the strongest conjugate acid, which is the strongest conjugate base? nitrous acid:HNO 2 / NO 2 - oxalic acid (2):HC 2 O 4 - / C 2 O 4 2- arsenic acid (2):HAsO 4 2- / AsO 4 3- carbonic acid (1):H 2 CO 3 / HCO 3 -

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. 1.sodium acetate (basic (pH=8.88), acetate (pK b =9.25) hydrolyses to produce OH - ) 2.ammonium chloride (acidic (pH=5.12), ammonium (pK a =9.25) hydrolyses to produce H 3 O + ) 3.calcium chloride (neutral, no hydrolysis) 4.sodium monohydrogen phosphate (basic (pH=10.12), HPO 4 2- (pK b 2 =6.79) hydrolyses to produce OH - ) (you need to consider two conjugate acid/base pairs..)

pH and % Dissociation of a Monoprotic Weak Acid CH 3 COOH = CH 3 COO - + H + K a = 1.75 x x x x (We let x = [H + ]) What is the pH of 0.10 M CH 3 COOH? [CH 3 COO - ] [H + ] [CH 3 COOH] x2x x 1.75 x K a = = = Approximation Method: Since K a <<1, assume x<<0.1 x 2 = 0.1 * 1.75 * = 1.75 x (x << 0.1) and x = [H + ] = M

Calculating % Dissociation and the pH CH 3 COOH = CH 3 COO - + H + % dissociation = 100 * [CH 3 COO - ] = x 0.1 = 1.3% [CH 3 COOH] init [H + ] = 1.3 x M pH = - log 10 [H + ] = 2.88

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

K a and Acid Strength The stronger the acid, the larger the K a and the smaller the pK a : CH 3 COOH (aq) = CH 3 COO - (aq) + H + (aq) K a = 1.76 x HCN (aq) = CN - (aq) + H + (aq) K a = 6.17 x pK a = 4.75 HNO 2 (aq) = NO 2 - (aq) + H + (aq) K a = 4.6 x pK a = 3.34 pK a = 9.21 stronger weaker

Weak Acids = = = = = = = weaker stronger

K b and pK b Arrhenius bases liberate OH - in solution. K b is the equilibrium constant for this reaction. NH 4 OH (aq) = NH 4 + (aq) + OH - (aq) K b = [NH 4 + ] [OH - ] [NH 4 OH] = 1.76 x pK b = - log 10 K b (definition) pK b = - log 10 (1.8 x ) = 4.74

K b and Base Strength The stronger the base, the larger the K b and the smaller the pK b : NH 4 OH (aq) = NH 4 + (aq) + OH - (aq) K b = 1.8 x pK b = 4.74 stronger weaker PO 4 3- (aq) + H 2 O (l) = HPO 4 2- (aq) + OH - (aq) K b = 4.5 x pK b = 1.34 Conclusion: phosphate anion is a stronger base than NH 4 OH.

K b 's of weak bases Strength (Ranked)

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

Some Lewis Acids and Bases Lewis bases are characterized by having an available lone pair. Examples are: O-H -.. : N-H H H : O-H H :.. I :.. : S :.. 2- hydroxide iodide ammonia water sulfide Lewis acids are electron deficient - i.e., electron pair acceptors Examples are: H + Zn 2+ Hg 2+ Ag + BF 3 metal cations electron deficient compounds

Lewis Acids/Bases - the Most General Definition Least general definition Most general definition Arrhenius Bronsted-Lowry Lewis H + (aq) + :OH - (aq) = H 2 O(l) Electron pair acceptor electron pair donor 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:

The Lewis Acid-Base Reaction Lewis definition: Acid: electron pair acceptor Base: electron pair donor HNO 2 + ClO 2 - = HClO 2 + NO 2 - Bronsted-Lowry: acid1 base2 acid2base1 The Lewis acid is H + (the electron deficient species) There are 2 bases (NO 2 - and ClO 2 - ), which compete for the acid The lone pairs donated by these bases are on oxygen atoms: O=N- O.. : O-N= O.. : [ ] - O=Cl- O.. : H O=N- O.. : H O=Cl- O.. : - O-Cl= O.. : [ ] + +

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 H + + : O-H - = H-O Lewis acid Lewis base.. : H F - B + N - H = F - B : N - H F F : F H H F H H Lewis acid Lewis base

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: Hg 2+ (aq) + I - (aq) = HgI + (aq) HgI + (aq) + I - (aq) = HgI 2 (s) (red-brown ppt) HgI 2 (s) + I - (aq) = HgI 3 - (aq) HgI 3 - (aq) + I - (aq) = HgI 4 2- (aq) Identify the Lewis acid and Lewis base in each reaction.

Complex Ions and Solubility Stepwise Lewis acid/base complexes form between Al 3+ (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. Al 3+ + :OH - = AlOH 2+ AlOH 2+ + :OH - = Al(OH) 2 + Al(OH) :OH - = Al(OH) 3 (s) Al(OH) 3 + :OH - = Al(OH) 4 - white precipitate