Chapter 17 Principles of Reactivity: Chemistry of Acids and Bases Chem 106, Chapter 17 Chapter 17 Principles of Reactivity: Chemistry of Acids and Bases Copyright 2007, David R. Anderson

Acids and Bases Arrhenius definition Brønsted definition Relative strengths of conjugate acid-base pairs Effect of structure on acid-base strength Acid-base equilibria Definitions pH, pOH Ka, Kb, pKa, pKb Kw: autodissociation of water pH Calculations Strong acids and bases Weak acids and bases Polyprotic acids Salts Lewis Acids and Bases

I. Acids and Bases A. Arrhenius definition B. Brønsted definition produces hydronium ion (H3O+) in aqueous solution produces hydroxide ion (OH–) in aqueous solution B. Brønsted definition acid: base: donates a proton (hydrogen ion, H+) accepts a proton HF + H2O H3O+ + F– A conjugate acid is formed by adding a proton to something. A conjugate base is formed by removing a proton from something. acid base conjugate acid conjugate base NH3 + H2O NH4+ + OH– base acid conjugate acid conjugate base

I. Acids and Bases B. Brønsted definition F– + H2O HF + OH– NH4+ + H2O water is amphoteric or amphiprotic (behaves as either an acid or a base) H2O + H2O H3O+ + OH– base acid autoionization of water: Brønsted definition fits even when not in aqueous solution: NH3 + HF NH4+ + F– base acid

I. Acids and Bases C. Relative strengths of conjugate acid-base pairs HA H3O+ + A–  If HA is a stronger acid then A– is a weaker base. If HA is a weaker acid then A– is a stronger base. B BH+ + OH–  If B is a stronger base then BH+ is a weaker acid. If B is a weaker base then BH+ is a stronger acid.

I. Acids and Bases D. Effect of structure on acid-base strength General: HA H3O+ + A– all acids produce this the place to look for a difference is here The strength of HA depends on the stability of A–. If A– is more stable (weaker base)  HA will be a stronger acid If A– is less stable (stronger base)  HA will be a weaker acid A– is more stable when: • A is more electronegative • the negative charge is more delocalized

I. Acids and Bases D. Effect of structure on acid-base strength 1. binary acids: hydrogen + one other element acidity increases (a) (b) HA H3O+ + A– (a) As A becomes more electronegative, A– is more stable (weaker base), and HA is a stronger acid. Thus: HF > H2O > NH3 > CH4 (dissociates: 1% 10–7% 10–16% 10–23%) Similarly: HCl > H2S > PH3 > SiH4

I. Acids and Bases D. Effect of structure on acid-base strength 1. binary acids: hydrogen + one other element acidity increases (a) (b) HA H3O+ + A– (b) As A– becomes larger, the negative charge is more delocalized, A– is more stable (weaker base), and HA is a stronger acid. (overcomes opposite trend in electronegativity) Thus: HI > HBr > HCl > HF strong acids (100% dissociated) weak acid (1% dissociated) Similarly: H2Te > H2Se > H2S > H2O

I. Acids and Bases D. Effect of structure on acid-base strength 2. oxoacids: HnXOm (HNO3, H2SO4, H3PO4, etc.) a. electronegativity of central atom (with same number of oxygens) e.g., HClO4 is a stronger acid than HBrO4 central atom withdraws electron density from the oxygen, stabilizes the anion Cl is more electronegative than Br  more stable anion, stronger acid Thus: HClO4 > HBrO4 > HIO4 HClO3 > HBrO3 > HIO3 HClO2 > HBrO2 > HIO2 HClO > HBrO > HIO And: HClO4 > H2SO4 > H3PO4 HBrO4 > H2SeO4 > H3AsO4 etc.

I. Acids and Bases D. Effect of structure on acid-base strength 2. oxoacids: HnXOm (HNO3, H2SO4, H3PO4, etc.) b. number of lone oxygens (with same central atom) negative charge more delocalized, anion more stable  stronger acid Thus: HClO4 > HClO3 > HClO2 > HClO HBrO4 > HBrO3 > HBrO2 > HBrO etc. And: H2SO4 > H2SO3 H2SeO4 > H2SeO3 HNO3 > HNO2 H3PO4 > H3PO3

I. Acids and Bases D. Effect of structure on acid-base strength Chem 106, Chapter 17 I. Acids and Bases D. Effect of structure on acid-base strength 3. carboxylic acids (10-6 % dissociated) (~1 % dissociated) negative charge delocalized, anion more stable, stronger acid Copyright 2007, David R. Anderson

I. Acids and Bases D. Effect of structure on acid-base strength Which in each pair would be the stronger acid? H2SO3 or HClO3 CH3OH or CH3SH HOBr or HBrO3 HBrO2 or HClO3 F2CHCO2H or FCH2CO2H ClCH2CH2CO2H or CH3CHClCO2H HNO3 or HNO2 PH3 or NH3 HClO4 or HBrO4 H2O or HF H2S or H2Se FCH2CO2H or ClCH2CO2H H2Se or HBr HF or HCl

I. Acids and Bases Summary: acid and base strength “Absolute” strengths (in water): HA H3O+ + A– strong acid nonbasic K ~  (100% dissociated) weak acid weak base K ~ 10–3 - 10–10 neutral strong base K ~ 10–15 nonacid v. strong base K 10–20 ~ <

I. Acids and Bases Summary: acid and base strength

I. Acids and Bases E. Acid-base equilibria HA + B BH+ + A– Chem 106, Chapter 17 I. Acids and Bases E. Acid-base equilibria HA + B BH+ + A– If: stronger stronger weaker weaker then K > 1 acid base acid base (equilibrium lies on right) On which side does the equilibrium lie in the following reactions? OCl– + HCl NH4+ + ClO4– F– + H2O NH2– + H2O Copyright 2007, David R. Anderson

II. Definitions A. pH, pOH Definition: pX = -logX pH = -log[H3O+] pOH = -log[OH–] [H3O+] = 10–pH [OH–] = 10–pOH [H3O+] pH 0.01 M 1 x 10–7 M 3 x 10–10 M pH [H3O+] 3.0 8.4

II. Definitions A. pH, pOH sig figs in pH calculations: e.g., [H3O+] = 2.3 x 10–3 M log [H3O+] = log 2.3 - 3 2 sig figs exact = 0.36 - 3.0000…. log [H3O+] = -2.64 pH = -log [H3O+] = 2.64 from exponent # of sig figs in [H3O+]  # of sig figs in [ ] = # of places past decimal in pH

II. Definitions B. Ka, Kb, pKa, pKb HA + H2O H3O+ + A– [H3O+][A–] Kc = ~ constant (55 M)  Kc·[H2O] = Ka = [H3O+][A–] [HA] acid dissociation constant pKa = -log Ka if pKa = 5 then Ka = 10–5 if pKa = 8 then Ka = 10–8 stronger acid weaker acid

II. Definitions B. Ka, Kb, pKa, pKb B + H2O BH+ +OH– [BH+][OH–] Kc = ~ constant  Kc·[H2O] = Kb = [BH+][OH–] [B] base dissociation constant pKb = -log Kb if pKb = 4 then Kb = 10–4 if pKb = 9 then Kb = 10–9 stronger base weaker base

II. Definitions C. Kw: autodissociation of water H2O + H2O H3O+ + OH– Kw = [H3O+][OH–] = 10–14 (constant at 25ºC)  [H3O+] = 10–14 [OH–] pKw = pH + pOH = 14 pH = 14 - pOH

III. pH Calculations A. Strong acids and bases 100% dissociated  for strong acid: [H3O+]eq = [HA]I base: [OH–]eq = [B]i e.g., 1.0 x 10–3 M HCl [H3O+] = pH = [OH–] = pOH = e.g., 2.5 x 10–2 M NaOH [OH–] = pOH = [H3O+] = pH = [OH–] -1.60 + 14 log pH = 12.40

III. pH Calculations B. Weak acids and bases [H3O+][A–] [HA] HA Solve equilibrium expressions [BH+][OH–] [B] B BH+ + OH– Kb = e.g., What is the pH of 0.10 M HC2H3O2? (Ka = 1.8 x 10–5) HC2H3O2 H3O+ + C2H3O2– 1.8 x 10–5 = x2 (0.10 - x) x = [H3O+] = 1.3 x 10–3 M (assumption valid) pH = 2.87 assume x << 0.10 1.8 x 10–5 = x2 (0.10)  If [HA]i  400·Ka then x << [HA]i (If not, then have to solve quadratic.)

III. pH Calculations B. Weak acids and bases e.g., What is the pH of 0.25 M N2H4? (Kb = 1.7 x 10–6) N2H4

III. pH Calculations B. Weak acids and bases e.g., If a 0.10 M solution of a weak acid has a pH of 4.28, what is the Ka for the acid?

III. pH Calculations B. Weak acids and bases e.g., What concentration of C5H5N (Kb = 1.7 x 10–9) will give a solution with pH = 9.80? C5H5N

III. pH Calculations C. Polyprotic acids H2SO4, H2SO3, H2CO3, etc. Lose their protons in separate steps: H2A H3O+ + HA– Ka1 = [H3O+][HA–] [H2A] (Usually, Ka1 >> Ka2) HA– H3O+ + A2– Ka2 = [H3O+][A2–] [HA–] Assume: 1) [H2A], [H3O+], and [HA–] can be determined from the 1st step. (i.e., HA– dissociates only very little.) 2) [A2–] can be determined from the 2nd step.

III. pH Calculations C. Polyprotic acids e.g., What are the equilibrium concentrations of all species in 0.10 M Vitamin C (ascorbic acid, H2C6H6O6)? H2A H3O+ + HA– Ka1 = 7.9 x 10–5 HA– H3O+ + A2– Ka2 = 1.6 x 10–12 Ka1 = [H3O+][HA–] [H2A] Ka2 = [H3O+][A2–] [HA–] 7.9 x 10–5 = x2 (0.10 - x) (0.10) = ~ 1.6 x 10–12 = (2.8 x 10–3 + y)(y) (2.8 x 10–3 - y) = ~ (2.8 x 10–3)(y) (2.8 x 10–3) = y x = 2.8 x 10–3 M y = 1.6 x 10–12 M  [H2A] = 0.10 M [H3O+] = [HA–] = 2.8 x 10–3 M [A2–] = 1.6 x 10–12 M

III. pH Calculations C. Polyprotic acids e.g., What are the equilibrium concentrations of all species in a 0.10 M solution of carbonic acid, H2CO3? (Ka1 = 4.2 x 10-7, Ka2 = 4.8 x 10-11)

III. pH Calculations D. Salts HA + B  BH+A– (salt) Salt may be acidic or basic depending on the nature of BH+ and A–. conjugate acid: BH+ + H2O B + H3O+ conjugate base: A– + H2O HA + OH–

III. pH Calculations D. Salts Ka and Kb for conjugate acid-base pairs: [H3O+][A–] [HA] HA H3O+ + A– Ka = A– HA + OH– Kb = [HA][OH–] [A–] usually only one or the other given in a table Ka · Kb = [H3O+][A–] [HA] · [HA][OH–] [A–] = [H3O+][OH–] = Kw or Ka = or Kb = Kw Ka Kb Ka · Kb = Kw for a conjugate acid-base pair

III. pH Calculations D. Salts 1. acidic salts e.g., What is the pH of a 0.10 M solution of NH4Cl? NH4+: Cl–: conjugate acid of a weak base  weak acid conjugate base of a strong acid  nonbasic (spectator) NH4+ NH3 + H3O+ Ka(NH4+) = Kw Kb(NH3) = [H3O+][NH3] [NH4+] 10–14 1.8 x 10–5 x2 (0.10 - x) x2 (0.10) = = ~ x = [H3O+] = 7.5 x 10–6 M pH = 5.13

III. pH Calculations D. Salts 2. basic salts e.g., What is the pH of a 0.10 M solution of NaCN? Ka(HCN) = 4.9 x 10–10 Na+: CN–:

IV. Lewis Acids and Bases acid: accepts a pair of electrons base: donates a pair of electrons HBr + H2O H3O+ + Br– Arrhenius, Brønsted, Lewis NH3 + H2O NH4+ + OH– Arrhenius, Brønsted, Lewis

IV. Lewis Acids and Bases HBr + NH3 NH4+ Br– Brønsted, Lewis NaI + CH3Br CH3I + NaBr Lewis only

IV. Lewis Acids and Bases NaOH + CO2 NaHCO3 Lewis only