CH4. Acids and Bases

Transcription

1 CH4. Acids and Bases 1 Bronsted-Lowry Bronsted-Lowry definitions: Acid = proton donor; Base = proton acceptor HF (aq) + H 2 O BL acid BL base H 3 O + (aq) + F - (aq) Fluoride ion is the conjugate base of HF Hydronium ion is the conjugate acid of H 2 O 2 1

2 Amphiprotic species Amphiprotic species that can act as BL acid or base NH 3 (aq) + H 2 O NH 4 + (aqu) + OH (aqu) BL base BL acid hydroxide K b = base dissociation constant = [NH 4+ ] [OH ] / [NH 3 ] H 2 O is amphiprotic - it s a base with HF, but an acid with NH 3 3 BL acid/base strength K a, the acidity constant, measures acid strength as: K a = [H 3 O + ] [A - ] / [HA] pk a = - log K a When ph = pk a, then [HA] = [A - ] For strong acids pk a < 0 pk a (HCl)

3 BL acid/base strengths 5 K w K w = water autodissociation (autoionization) constant 2 H 2 O H 3 O + (aqu) + OH - (aqu) K w = [H 3 O + ] [OH - ] = 1 x (at 25 C) Using the above, you should prove that for any conjugate acid-base pair: pk a + pk b = pk w =

4 Polyprotic acids Since pk a values are generally wellseparated, only 1 or 2 species will be present at significant concentration at any ph H 3 PO 4 + H 2 O H 2 PO H 3 O + pk a1 = 2.1 H 2 PO H 2 O HPO H 3 O + pk a1 = 7.4 HPO H 2 O PO H 3 O + pk a1 = Solvent leveling The strongest acid possible in aqueous solution is H 3 O + Ex: HCl + H 2 O H 3 O + (aq) + Cl - (aq) there is no appreciable equilibrium, this reaction goes quantitatively; the acid form of HCl does not exist in aqueous solution Ex: KNH 2 + H 2 O K + (aq) + OH - (aq) + NH 3 (aq) this is solvent leveling, the stable acid and base species are the BL acid-base pair of the solvent NH 2 - = imide anion NR 2-, some substituted imide ions are less basic and can exist in aq soln 8 4

5 Solvent leveling Only species with 0 < pk a < 14 can exist in aqueous solutions. The acid/base range for water stability pk w, i.e. 14 orders of mag in [H + ]. Other solvents have different windows and different leveling effects. 9 Solvent leveling 2EtOH EtOH 2+ (solv) + EtO (solv) K ~ chemistry in the range of -3 < pk a < 17 NH 3 NH 4+ (solv) + NH 2 (solv) ammonium imide chemistry in the range of 10 < pk a < 38 O 2 NH 3 (l) OH Na (m) Na + (solv) + NH 2 (solv) + ½ H 2 (g) slow very strong base Na + (solv) + e (solv) 10 5

6 Acid/base chemistry of complexes Aqueous chemistry: H 2 O Fe(NO 3 ) 3 [Fe(OH 2 ) 6 ] 3+ (aq) + 3 NO 3 (aq) 2 [Fe(OH 2 ) 6 ] 3+ (aq) [Fe 2 (OH 2 ) 10 OH] 5+ (aq) + H 3 O + (aq) Hexaaquairon(III), pk a ~ 3 dimer 11 Aqua, hydroxo, oxoacids aqua acid M(OH 2 ) x n+ ex: [Cu(OH 2 ) 6 ] 2+ hexaaquacopper(ii) cation hydroxoacid M(OH) x ex: B(OH) 3, Si(OH) 4 pk a ~ 10 oxoacid MO p (OH) q p and q designate oxo and hydroxo ligands ex: H 2 CO 3 (aq) + H 2 O HCO 3 (aq) + H 3 O + (aq) CO 2 (g) + H 2 O carbonic acid pk a ~ 3.6 bicarbonate 12 6

7 For aqueous ions: 1. Higher charge is more acidic Trends in acidity pk a of [Fe(OH 2 )] 3+ ~ 3 pk a of [Fe(OH 2 )] 6 2+ ~ 9 2. Smaller radius is more acidic Mn 2+ Cu 2+ early TM late TM lower Z* higher Z* => larger radius => smaller radius less acidic more acidic pk a vs z 2 / (r + + d) Na + (aqu) = [Na(OH 2 ) 6 ] + has pk a > 14 so it s a spectator ion in aqu soln 13 Anhydrides Ex: H 2 O + SO 3 H 2 SO 4 Acidic SO 3 / H 2 SO 4 anhydride acid form P 2 O 5 / H 3 PO 4 CO 2 /H 2 CO 3 Basic Na 2 O / NaOH Amphoteric Al 2 O 3 / Al(OH)

8 Trends in acidity 15 Common acids HNO 3 NO 3 (D 3h ) Nitric acid Nitrate HNO 2 NO 2 (C 2v ) Nitrous acid Nitrite You should know these! H 3 PO 4 PO 4 3 (Td) Phosphoric acid Phosphate H 3 PO 3 HPO 3 2 (C 3v ) Phosphorous acid Phosphite 16 8

9 Common acids H 2 SO 4 (Td) Sulfuric acid SO 4 2 Sulfate You should know these! H 2 SO 3 (C 3v ) Sulfurous acid SO 3 2 Sulfite 17 Common acids HClO 4 Perchloric acid ClO 4 (Td) Perchlorate HClO 3 ClO 3 (C 3v ) Chloric acid Chlorate You should know these! HClO 2 ClO 2 (C 2v ) Chlorous acid Chlorite HOCl Hypochlorous acid OCl Hypochlorite 18 9

10 Pauling s rules for pk a s of oxoacids 1. Write formula as MO p (OH) q 2. pk a 8 5p 3. Each succeeding deprotonation increases the pk a by 5 Ex: rewrite HNO 3 as NO 2 (OH) p = 2; pk a 8 5(2) 2 (exptl value is 1.4) Ex: rewrite H 3 PO 4 as PO(OH) 3 p = 1; pk a1 8 5(1) 3 (exptl value is 2.1) pk a2 8 (exptl value is 7.4) pk a3 13 (exptl value is 12.7) 19 pk a values p Pauling pk a calcn exptl Cl(OH) ClO(OH) ClO 2 (OH) ClO 3 (OH) HlO 4 + 2H 2 O H 5 IO

11 Amphoteric oxides [Al(OH 2 ) 6 ] 3+ Al 2 O 3 / Al(OH) 3 [Al(OH) 4 ] Oh H 3 O + OH Td 2 [Al(OH 2 ) 6 ] 3+ (aq) [Al 2 (OH 2 ) 10 (OH)] 5+ (aq) + H 3 O + (aq) pk a ~ 2 dimer 21 polyoxocations linear trimer is [Al 3 (OH 2 ) 14 (OH) 2 ] 7+ Keggin ion [AlO 4 (Al(OH) 2 ) 12 ] 7+ ph 4 charge/volume ratios Al(OH 2 ) 6 3+ > dimer > trimer --- > Al(OH) 3 3+ / Oh 5+ / 2 Oh 7+ / 3 Oh neutral 22 11

12 Polyoxoanions H 3 O + VO 3 4 (aq) V 2 O 5 (s) orthovanadate (Td) 2 VO 4 3 (aq) + H 2 O V 2 O 7 4 (aq) + 2OH (aq) H 3 O + V 3 O 9 3 V 3 O 10 5 oxo bridge H 3 O + V 4 O Lewis acids and bases A + B: A:B LA LB complex LA = electron pair acceptor; LB = electron pair donor Lewis definition is more general than BL definition, does not require aqueous or protic solvent Ex: W + 6 :CO [W(CO) 6 ] BCl 3 + :OEt 2 BCl 3 :OEt 2 D 3h Fe 3+ (g) + 6 :OH 2 [Fe(OH 2 ) 6 ]

13 LA/LB strengths LA strength is based on reaction K f LA/LB strengths depend on specific acid base combination Ex: BCl 3 + :NR 3 Cl 3 B:NR 3 K f : NH 3 < MeNH 2 < Me 2 NH < Me 3 N inductive effect BMe 3 + :NR 3 Me 3 B:NR 3 K f : NH 3 < MeNH 2 < Me 2 NH > Me 3 N inductive + steric H rxn kj/mol 25 log K and ligand type 26 13

14 Drago-Wayland equation A (g) + :B (g) A:B (g) Gas phase reactions (omits solvation effects) - H rxn = E A E B + C A C B look up E, C values for reactants (Table 4.4) 27 Donor/Acceptor numbers Commonly used to choose appropriate solvents (Table 4.5) Donor Number (DN) is derived from H rxn (SbCl 5 + :B Cl 5 Sb:B) higher DN corresponds to stronger LB Acceptor Number (AN) is derived from stability of Et 3 P=O:A complex higher AN corresponds to stronger LA Ex: THF (tetrahydrofuran) C 4 H 8 O DN AN ε dielectric constant THF H 2 O Some Li + salts and BF 3 have similar solubilities in THF, H 2 O NH 3 is much more soluble in H 2 O Most salts are much more soluble in H 2 O 28 14

15 Descriptive chemistry - Group 13 Expect inductive effect BF 3 > BCl 3 > BBr 3 but the opposite is true ex: BF 3 is stable in H 2 O, R 2 O (ethers) BCl 3 rapidly hydrolyzes due to nucleophilic attack of :OH 2 the lower acidity of BF 3 is due to unusually favorable B X bonding in the planar conformation due to interaction AlCl 3 is a dimer (Al 2 Cl 6 ) General trend larger central atom, tends to have higher CN Al 2 Me 6 is isostructural with Al 2 Cl 6 C 6 H 6 C 6 H 5 C(O)R Friedel-Crafts RC(O)-X: + AlCl 3 RC(O) + AlCl 3 X 29 Descriptive chemistry - Group 14 CX 4 is not a Lewis Acid Acidity SiF 4 > SiCl 4 > SiBr 4 > SiI 4 (inductive effect) ex: 2KF(s) + SiF 4 (g) K 2 SiF 6 (s) LB LA SiF 6 2 Oh SnF 4 and PbF 4 have Oh not Td coordination (heavier congener, higher CN) each M has 2 unique axial F and 4 shared F 30 15

16 Descriptive chemistry - Group 15 MF 5 does not exist for M=N; trigonal bipyramidal for M = P, As SbF 5 : Sb has Oh coordination (oligomerizes to Sb 4 F 20 or Sb 6 F 30 ) LB LA transient K 2 MnF 6 (s) + 2 SbF 5 (l) MnF 4 + 2KSbF 6 (s) F transfer KF, H 2 O 2 aqu HF KMnO 4 Sb 2 O 3 MnF 3 + ½ F 2 (g) Dove (1980 s), chemical synthesis of F 2 gas 31 Descriptive chemistry - Group 16 Inductive effect stabilizes conjugate base (anionic form) sulfuric acid fluorosulfonic HSO 3 F / SbF 5 pk a ~ 2 pk a ~ 5 pk a ~ 26 (superacid) HSO 3 F / SbF 5 C 6 H 6 C 6 H 7+ SbF