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Acids, Bases and Solution Equilibria

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Acids, Bases and Solution Equilibria Acids, bases and solution equilibria A four lecture course for the 1st year Jose M. Goicoechea http://course.chem.ox.ac.uk/acids-bases-and- solutions-equilibria-year-1-2014.aspx http://goicoechea.chem.ox.ac.uk/teaching.html Acid-base reactions + + NH3 + H3O NH4 + H2O + - H2O + HI H3O + I - 2- + HSO4 + H2O SO4 + H3O NH3 + BF3 NH3:BF3 C5H5N + I2 C5H5N:I2 Species highlighted act as acids Redox reactions A redox reaction is a reaction in which there is a change in oxidation state Fe3+(aq) + Cr2+(aq) Fe2+(aq) + Cr3+(aq) + 2+ Zn(s) + 2H3O Zn + H2(g) 2PCl3 + O2 2OPCl3 Ca(s) + H2 CaH2(s) + - O2 + Pt + 3F2 [O2] [PtF6] One or two electrons are transferred entirely Species highlighted act as oxidants Definitions of Acid/Base Arrhenius/Ostwald Brønsted/Lowry Lux/Flood ‘Solvent system’ Lewis Usanovich Arrhenius/Ostwald Acids and bases dissociate in H2O, + + releasing H (H3O ) and OH-. Arrhenius Ostwald + - H2O H (aq) + OH (aq) + H (aq) is an acid - OH (aq) is a base Brønsted/Lowry Proton theory retained but the definition is now independent of solvent Brønsted Lowry An acid is a proton donor and a base is a proton acceptor. + - Other solvents are also NH4 + NH2 2NH3 capable of self-ionisation + - Proton HCl + NH3 [NH ] Cl transferred from 4 acid to base Donor acid Acceptor base Brønsted/Lowry Every acid has a conjugate base and every base has a conjugate acid HA + B- A- + HB The conjugate base of a weak acid is a strong base, and the conjugate base of a strong acid is a weak base. TRUE FACT! Lux/Flood A definition for anhydrous/dry systems: used in solid state chemistry The concept focuses on the oxide ion (O2-) A base is a oxide donor and an acid is an oxide acceptor. CaO + SiO2 CaSiO3 The Lux-Flood base is a basic anhydride CaO + H O Ca2+ + 2OH- 2 The Lux-Flood acid is an acid anhydride. SiO2 + H2O H2SiO3 Solvent system + - This definition 2H2O H3O + OH is based on 2NH NH + + NH - solvent 3 4 2 autoionisation + - 2H2SO4 H3SO4 + HSO4 cationsolv anionsolv An acid is a species that increases the concentration of cationsolv A base is a species that increases the concentration of anionsolv Lewis An acid is an electron pair acceptor A base is a an electron pair donor NH3 + BF3 NH3BF3 H H H N + H N B H B H H H H H H H Has ‘lone’ pair Has six valence electrons More examples of Lewis acid/base interactions OH2 3+ 3+ H2O OH2 Al + 6H2O Al H2O OH2 OH2 All metal cations in donor solvents are Lewis acids interacting with solvent molecules which act as Lewis bases + + Ag + 2NH3 H3N Ag NH3 Usanovich A broad definition which encompasses all other acid/base concepts Usanovich An acid is a species that Lewis reacts with bases. It gives up cations or Ionotropic accepts anions or electrons A base is a species that Brønsted Solvent Lowry Arrhenius system reacts with acids. It gives up anions, combines with cations, or donates electrons. Lux Flood Proton equilibria in water + - HX(aq) + H2O H3O + X (aq) [H O+][X-] K = 3 a [HX] pKa = –log10Ka + pH = –log10[H3O ] Autoprotolysis - + (aq.) (aq.) 2H2O H3O + OH + - Kw = [H3O ][OH ] –14 For H2O, Kw = 10 at 25°C Easy way to treat pKa [H O+][X-] Since K = 3 a [HX] + - - Ka = [H (aq.)] when [HX] = [X ] i.e. when HX is 50% converted to X pKa is pH at which HX is 50% dissociated – Acid HA A Ka pKa Hydriodic HI I– 1011 –11 - 10 Perchloric HCIO4 CIO4 10 –10 Hydrobromic HBr Br– 109 –9 Hydrochloric HCl Cl– 107 –7 - 2 Sulfuric H2SO4 HSO4 10 –2 – 2 Nitric HNO3 NO3 10 –2 + Hydronium ion H3O H2O 1 0.0 – –1 Chloric HCIO3 CIO3 10 1 – –2 Sulfurous H2SO3 HSO3 1.5 × 10 1.81 - 2- –2 Hydrogensulfate ion HSO4 SO4 1.2 × 10 1.92 - –3 Phosphoric H3PO4 H2PO4 7.5 × 10 2.12 Hydrofluoric HF F– 3.5 × 10–4 3.45 – –4 Formic HCOOH HCO2 1.8 × 10 3.75 – –5 Ethanoic CH3COOH CH3CO2 1.74 × 10 4.76 + –6 Pyridinium ion HC5H5N C5H5N 5.6 × 10 5.25 – –7 Carbonic H2CO3 HCO3 4.3 × 10 6.37 – –8 Hydrogen sulfide H2S HS 9.1 × 10 7.04 - 2- –8 Dihydrogenphosphate ion H2PO4 HPO4 6.2 × 10 7.21 + –10 Ammonium ion NH4 NH3 5.6 × 10 9.25 Hydrocyanic HCN CN– 4.9 × 10–10 9.31 – 2– –11 Hydrogencarbonate ion HCO3 CO3 4.8 × 10 10.32 2- 3- –13 Hydrogenphosphate ion HPO4 PO4 2.2 × 10 12.67 Hydrogensulfide ion HS– S2– 1.1 × 10–19 19 Polyprotic acids A polyprotic acid loses protons in succession, with each deprotonation becoming progressively less favourable. The behaviour of such species in solution be represented by a distribution diagram pKa1 = 2.21 pKa2 = 7.21 pKa3 = 12.68 [H3PO4] a(H3PO4) = - 2- 3- [H3PO4] + [H2PO4 ] + [HPO4 ] + [PO4 ] Hammet acidity function H+ + X- HX For any base B: H+ + B HB+ [BH+] H = pK + – log 0 BH [B] [BH+] can be measured with a dye [B] Very concentrated solutions: Hammet acidity function HX + Ind IndH+ + X- NH2 Ind is a coloured indicator e.g. [Ind] NO2 H = pK + + log 4-Nitroaniline 0 IndH [IndH+] H0 is the Hammett function of HX Pure H2SO4, H0 = –11.9 Superacid (HSO3F), H0 = –15.1 Strength of Brønsted acids and bases Two principal factors control acid/base strength: 1) Inherent properties of the species itself To quantify this value we must measure proton transfer reactions in the gas phase. Chemistry controlled by ΔH + + H (g) + X(g) HX (g) Ap = –ΔH A = proton affinity of X p 2) Environment (e.g. solvation) WORTH REMEMBERING! Proton-Transfer reactions - + HA(g) + B(g) A (g) + HB (g) Can be considered as two “half reactions” ΔHB + + H (g) + B(g) HB (g) ΔHA + - H (g) + A (g) HA(g) Then ΔHrxn = ΔHB – ΔHA ΔHrxn = Ap(A) – Ap(B) The reaction is favourable is Ap(B) > Ap(A) + - H(g) + e + X(g) EA –ΔHion + - H(g) + X(g) H(g) + X(g) –Ap –BHX HX(g) Ap = BHX + ΔHion – EA Ap = BHX + ΔHion - EA ROW Decreasing EA GROUP Increasing BHX Decreasing Ap; i.e., increasing acidity Proton affinity (Ap; KJ/mol) - - - - - H CH3 NH2 OH F 1675 1745 1689 1635 1554 - - - - SiH3 PH2 SH Cl 1554 1552 1469 1395 Ap for - - - - H2O: 723 GeH3 AsH2 SeH Br 1509 1515 1466 1354 I- 1315 - - PH2 and F have almost identical Ap values but their pKa values differ (27 to 3.45, respectively) making F- a much stronger acid in water. Predictions ignoring solvation - + HI + H2O I + H3O 723 1315 - + NH3 + H2O OH + NH4 872 1635 This can’t be right, other factors must be at play Taking into account solvation effects B(aq) + H+(aq) BH+(aq) Large –ΔHhyd. or alternatively… - B (aq) + H+(aq) BH(aq) The Born formula tells us 2 Large –ΔH –z hyd that AHsolv is proportional to: r Effective proton affinity (Ap′; KJ/mol) and pKa - - - - - H CH3 NH2 OH F 1380 1351 1188 1150 49 39 15.7 3.2 SiH - PH - SH- Cl- 3 2 ′ Ap for 1283 1090 H2O: 1130 35 27 7.05 –6.1 - - - - GeH3 AsH2 SeH Br 1079 25 23 3.8 –8 I- 1068 –9 Predictions taking into account solvation - + HI + H2O I + H3O 1130 1068 - + NH3 + H2O OH + NH4 1195 1188 Range of discrimination + - HY + HY H2Y + Y + - Kauto = [H2Y ][Y ] pKauto = –logKauto Acid-base discrimination windows NH3 H2O CH3COOH HF 1400 1200 1000 800 Ap′/KJ/mol Ap vs. Ap′ 1900 - CH3 - NH2 1700 OH- F- - PH2 CN- 1500 Cl- Br- I- 1300 1100 900 py NH3 700 H2O 500 Solvent levelling + + H2F (aq) + H2O HF + H3O (aq) - - NH2 (aq) + H2O OH (aq) + NH3 + Strongest acid that can exist is H2Y Strongest base is Y- Metal aqua ion acidity (n-1)+ OH2 n+ OH2 H2O OH2 H2O OH2 M M H2O OH2 H2O OH OH2 OH2 n+ (n–1)+ + Maq. MOH + Haq. Arrange in order of increasing acidity: + 3+ 2+ 2+ [Na(OH2)6] , [Sc(OH2)6] , [Mn(OH2)6] , [Ni(OH2)6] + 2+ 2+ 3+ [Na(OH2)6] < [Mn(OH2)6] < [Ni(OH2)6] < [Sc(OH2)6] Hydrolysis constants (pKhyd) for cations at 25°C Li+ Be2+ 13.9 6.2 Na+ Mg2+ Al3+ 14.7 11.4 5.0 K+ Ca2+ Sc3+ Ti3+ V2+ Cr2+ Mn2+ Fe2+ Co2+ Ni2+ Cu2+ Zn2+ Ga3+ - 12.6 4.7 2.3 6.5 8.7 10.6 10.1 9.6 10.0 7.6 9.5 2.6 V3+ Cr3+ Fe3+ Co3+ 2.6 3.9 2.0 3.2 Rb+ Sr2+ Y3+ Rh3+ Pd2+ Ag+ Cd2+ In3+ - 13.1 8.0 3.2 1.6 11.8 7.9 3.2 Cs+ Ba2+ La3+ Ir3+ Pt2+ Hg2+ Tl+ - 13.3 9.5 4.8 >2.5 2.5 13.3 Burgess, J. Metal Ions in Solution; Ellis Horwood: Chichester, 1978, pp. 264-265 Pauling’s rules for oxoacids OpE(OH)q Rule 1: pKa1 = 8 – 5p WORTH REMEMBERING! The higher the number of oxo- groups bound to the central element, the greater the acidity of the acid. HClO4; O3Cl(OH); pKa = –10 O O Cl Cl The higher acidity O O O O is due to the O O greater number of resonance structures of the conjugate base O O Cl Cl O O O O O O Pauling’s rules to the test pKa1 = 8 – 5p p = 0 1 2 3 O Cl Cl HO Cl Cl HO O HO O HO O Hypochlorous Chlorous O Chloric Perchloric O pKa= 7.2 2.0 -1.0 -10 Pauling’s rules for oxoacids The successive pKa values of polyprotic acids (those with q > 1), increase by five for each successive proton transfer Rule 2: pKa2 = pKa1 + 5 pKa3 = pKa2 + 5 pKa4 = pKa3 + 5 WORTH REMEMBERING! O pKa1 = 2.1 P OH pKa2 = 7.4 HO OH pKa3 = 12.7 O pK = 1.8 P a1 OH HO pKa2 = 6.6 H O pKa1 = 2.3 As pKa2 = 6.9 OH HO pKa3 = 11.5 OH O Compare the acidities of the following acids: P OH pK = 8 – 5p = 3 HH3OPO4 a OH Real value: 2.12 O H SO 2 4 pK = 8 – 5p = –2 S a OH O Real value: –2 OH O HClO4 Cl pKa = 8 – 5p = –7 HO O O Real value: –10 So what’s up with boric acid? OH According to Pauling’s first rule for oxo- HO B acids it should have a pK of 8.
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