Published on January 2, 2008
Chem 1310: Introduction to physical chemistry Part 5: Buffers and solubility: Chem 1310: Introduction to physical chemistry Part 5: Buffers and solubility Peter H.M. Budzelaar Buffers: Buffers Water has a very small [H3O+] (10-7). Adding just a little bit of acid or base can change the pH drastically. Add 0.001 M HCl: pH goes from 7 to 3! For many applications this sensitivity is undesirable. One of the best ways to prevent pH swings is buffering: the use of a mixture of a weak acid and its conjugate base (which will be a weak base). Buffers: Buffers Two important aspects: What will be the resulting pH? What will be the buffer capacity (how much acid/base can be absorbed before the pH starts to change drastically)? The pH of a buffer solution: The pH of a buffer solution Take a mixture of HOAc and NaOAc (both 0.1 M): HOAc + H2O ⇋ OAc- + H3O+ The pH of a buffer solution (2): The pH of a buffer solution (2) We usually assume x « buffer concentration, so (Always check afterwards! If not valid, solve the full quadratic equation) The pH of a buffer solution (3): The pH of a buffer solution (3) General formula (using the same x assumption): where we replace the actual [HA], [A-] by the amounts weighed in (Henderson-Hasselbalch). Buffer capacity: Buffer capacity How much added acid or base can a buffer absorb? At most until the buffer acid or its conjugate base is consumed. If you have 1L of a buffer containing 0.2 M HOAc and 0.35 M NaOAc, this can absorb up to 0.35 moles of acid (all NaOAc consumed) or 0.2 moles of base (all HOAc consumed). As long as you do not exceed the buffer capacity, you can calculate the new pH using Henderson-Hasselbalch. Buffer capacity (2): Buffer capacity (2) 0.03 moles of HCl is added to 1L of a buffer of 0.1 M each of HOAc and NaOAc. What is the resulting pH? New [HOAc] = 0.1+0.03 = 0.13, new [NaOAc] = 0.1-0.03=0.07: Titration: Titration Slowly add acid of known concentration from a burette to a solution of base (or vv). Use an indicator to detect moment of fast pH change (happens at equivalence point). Strong acid, base largest pH change, almost any indicator will work. Weak acid titrated with strong base: Solution will originally not be very acidic, but will go till very basic. Use indicator for pH > 7, e.g. phenolphtalein. Titration: Titration Weak base titrated with strong acid: Solution will originally not be very basic, but will go till very acidic. Use indicator for pH < 7, e.g. methyl red. Do not titrate a weak acid with a weak base! No clear equivalence point. Solubility in water: Solubility in water Just another equilibrium (see MSJ p839 for Ksp table): AgCl(s) ⇋ Ag+(aq) + Cl-(aq) KC = Ksp = [Ag+][Cl-] The standard rules for writing equilibrium constants apply: Mg3(PO4)2 (s) ⇋ 3 Mg2+(aq) + 2 PO43-(aq) Ksp = [Mg2+]3[PO43-]2 No AgCl, because that is a pure solid. Calculating the solubilityof a compound in pure water: Calculating the solubility of a compound in pure water Add excess AgCl to water; it starts to dissolve: [Ag+][Cl-] = x2 = Ksp = 1.8·10-10 x » 1.3·10-5 mol/L Calculating the solubilityof a compound in pure water (2): Calculating the solubility of a compound in pure water (2) Add excess PbCl2 to water; it starts to dissolve: [Pb2+][Cl-]2 = 4x3 = Ksp = 1.7·10-5 x » 0.015 mol/L Calculating solubilityin the presence of "common ions": Calculating solubility in the presence of "common ions" Dissolve AgCl in a solution of 0.1 M NaCl: [Ag+][Cl-] = x(0.1+x) = Ksp = 1.8·10-10 Calculating solubilityin the presence of "common ions" (2): Calculating solubility in the presence of "common ions" (2) Assume x « 0.1: x(0.1+x) » 0.1x = Ksp = 1.8·10-10 Þ x = 1.8·10-9 mol/L (verify: x « 0.1!) A lot less soluble than in pure water! Without assumption: solve the quadratic equation. This is often not a good idea! Calculating whethera salt will precipitate: Calculating whether a salt will precipitate Calculate Qsp = [...][...] (same formula as for Ksp) Qsp < Ksp: more could dissolve Qsp = Ksp: saturated solution Qsp > Ksp: super-saturated: salt will precipitate (c.f. Q and K for other equilibria) Solubility calculations are not always straightforward...: Solubility calculations are not always straightforward... The real solution equilibrium might be more complicated: PbCl2(s) ⇋ PbCl+ + Cl- ⇋ Pb2+ + 2 Cl- The original Ksp expression is still valid, but we cannot assume all Pb in solution is present as Pb2+. There will also be some PbCl+, so the amount of Pb that goes into solution will be higher than expected. Solubility calculations are not always straightforward...: Solubility calculations are not always straightforward... Added reagents may complex with the solutes and reduce their concentrations, setting up new equilibria: AgCl(s) ⇋ Ag+(aq) + Cl-(aq) Ag+(aq) + 2 CN-(aq) ⇋ Ag(CN)2- (aq) AgCl(s) + 2 CN-(aq) ⇋ Ag(CN)2- (aq) + Cl-(aq) (Hess's law) Solubility and complexation: Solubility and complexation We have 1L of a solution 0.1 M in NaCN. Will it dissolve 0.01 moles of AgCl? Assuming complete conversion to Ag(CN)2-: [Cl-] = 0.01, [Ag(CN)2-] = 0.01, [CN-] = 0.008: a) b) Either way: it will easily dissolve! Solubility and acid-base reactions: Solubility and acid-base reactions CaCO3(s) ⇋ Ca2+(aq) + CO32-(aq) CO32- + H2O ⇋ HCO3- + OH- Part of CO32- removed via reaction with water Þ more will dissolved than you would calculate from Ksp. With added acid: CO32- + H3O+ ® HCO3- + H2O HCO3- + H3O+ ® H2CO3 + H2O H2CO3 ® CO2 + H2O Solubility and acid-base reactions: Solubility and acid-base reactions Eventually, all CaCO3 dissolves in acid! This happens with many poorly soluble salts of weak acids (S2-, CO32-, F-), except when Ksp is really very small (PbS, HgS, ...).