Chem 1310: Introduction to physical chemistry Part 5: Buffers and solubility

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Chem 1310 Introduction to physical chemistry
Part 5 Buffers and solubility

• Peter H.M. Budzelaar

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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).

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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)?

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The pH of a buffer solution

• Take a mixture of HOAc and NaOAc (both 0.1 M)
• HOAc H2O ? OAc- H3O

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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)

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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).

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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.

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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 0.10.03 0.13,new NaOAc
  0.1-0.030.07

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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 gt
  7, e.g. phenolphtalein.

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Titration

• Weak base titrated with strong acid
• Solution will originally not be very basic, but
  will go till very acidic. Use indicator for pH lt
  7, e.g. methyl red.
• Do not titrate a weak acid with a weak base!
• No clear equivalence point.

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Solubility in water

• Just another equilibrium (see MSJ p839 for Ksp
  table)
• AgCl(s) ? Ag(aq) Cl-(aq)
• KC Ksp AgCl-
• The standard rules for writing equilibrium
  constants apply
• Mg3(PO4)2 (s) ? 3 Mg2(aq) 2 PO43-(aq)
• Ksp Mg23PO43-2

No AgCl, becausethat is a pure solid.
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Calculating the solubilityof a compound in pure
water

• Add excess AgCl to water it starts to dissolve
• AgCl- x2 Ksp 1.810-10
• x 1.310-5 mol/L

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Calculating the solubilityof a compound in pure
water (2)

• Add excess PbCl2 to water it starts to dissolve
• Pb2Cl-2 4x3 Ksp 1.710-5
• x 0.015 mol/L

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Calculating solubilityin the presence of "common
ions"

• Dissolve AgCl in a solution of 0.1 M NaCl
• AgCl- x(0.1x) Ksp 1.810-10

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Calculating solubilityin the presence of "common
ions" (2)

• Assume x 0.1x(0.1x) 0.1x Ksp
  1.810-10Þ x 1.810-9 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!

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Calculating whethera salt will precipitate

• Calculate Qsp ......(same formula as for
  Ksp)
• Qsp lt Ksp more could dissolve
• Qsp Ksp saturated solution
• Qsp gt Ksp super-saturated salt will precipitate
• (c.f. Q and K for other equilibria)

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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.

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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)

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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!

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Solubility and acid-base reactions

• CaCO3(s) ? Ca2(aq) CO32-(aq)
• CO32- H2O ? HCO3- OH-
• Part of CO32- removed via reaction with waterÞ
  more will dissolved than you would calculate from
  Ksp.
• With added acid
• CO32- H3O HCO3- H2O
• HCO3- H3O H2CO3 H2O
• H2CO3 CO2 H2O

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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, ...).