Chem 163 chapter 19
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CHEM 163 Chapter 19. Spring 2009. Buffers. Solution that resists pH changes Ex. Blood (pH ~ 7.4) Acid must neutralize small amounts of base Base must neutralize small amounts of acid Acid and base must not neutralize each other. Added in as salt (NaCH 3 COO).

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CHEM 163 Chapter 19

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Chem 163 chapter 19

CHEM 163Chapter 19

Spring 2009


Buffers

Buffers

Solution that resists pH changes

  • Ex. Blood (pH ~ 7.4)

  • Acid must neutralize small amounts of base

  • Base must neutralize small amounts of acid

  • Acid and base must not neutralize each other

  • Added in as salt (NaCH3COO)

    Use conjugate acid-base pairs!

    CH3COOH (aq) + H2O (l)

    CH3COO- (aq) + H3O+ (aq)

    Common-ion effect

    Ex: acetate


    Chem 163 chapter 19

    • High concentrations of weak acid/conjugate base

    • Add H3O+ or OH-

      • Added amounts are relatively small

      • Cause only small shifts

      • React with weak acid or conjugate base

    HA (aq) + H2O (l)

    A- (aq) + H3O+ (aq)

    HA (aq) + OH- (aq)

    A- (aq) + H2O (l)

    pH depends on [HA]/[A-] ratio


    Making a buffer

    Making a buffer

    • Choose the conjugate acid-base pair (pKa ≈ pH)

    • Calculate the ratio of acid-base concentrations

    • Determine the buffer concentration

    • Mix solution; adjust pH

    Henderson-Hasselbalch equation:


    Buffer properties

    Buffer Properties

    • Buffer Capacity:

      • Ability to resist pH change

      • Unrelated to pH of buffer

      • Dependent on concentration of weak acid/conj base

      • Highest when [A-] = [HA]

    • Buffer Range:

      • pH range over which buffer is effective

      • Usually within ±1 pH unit of the pKa of weak acid


    Sample problem

    Sample Problem

    Make 200. mL of a pH 3.5 citric acid/sodium citrate buffer with an acid concentration of 0.50 M.

    We are given solid sodium citrate (294 g/mol) and 5.0 M citric acid. The pKa of citric acid is 3.15.


    Measuring ph

    Measuring pH

    • Acid-Base Titration Curves: pH v. volume titrant

      Measuring pH:

      • pH meter

      • Acid-base indicators

        Indicator:

      • Weak organic acid

      • HIn different color than In-

      • Intensely colored (small amount needed)

      • Changes color over ~ 2 pH units


    Titration curves strong acid strong base

    Titration Curves: Strong acid – Strong base

    • Low pH (strong acid)

    • Sudden pH rise (6-8 units)

    • Slow pH increase

     [OH-]added ≈ [H3O+]init

    Equivalence point:

    [OH-]added = [H3O+]init

    pH = 7

    End point:

    when indicator changes color


    Calculating ph during titration

    Calculating pH during titration

    • Original solution of strong HA

    • Before the equivalence point

      • Moles of acid remaining?

      • Calculate [H3O+]

    • At the equivalence point: pH = 7

    • After the equivalence point

      • Excess moles of OH- added?

      • Calculate [OH-]

    moles base

    added

    moles acid

    total

    moles acid initial

    moles acid rxted


    Titration curves weak acid strong base

    Titration Curves:Weak acid – Strong base

    • Higher initial pH (weak acid, lower Ka)

    • Buffer region

      • gradual pH rise

      • Midpoint:

        ½ initial HA reacted

    • Equivalence point:

      pH > 7.00

    • Slow pH increase

     [HA] = [A-]

     pH = pKa


    Calculating ph during titration1

    Calculating pH during titration

    • Original solution of weak HA

      • ICE table

    • Buffer Region

    • At the equivalence point:

    • After the equivalence point

    •  x = [H3O+]

    • Excess moles of OH- added


    Titration curves strong acid weak base

    Titration Curves:Strong acid – Weak base

    • Initial pH > 7.00 (weak base)

    • Buffer region

      • gradual pH decrease

    • Equivalence point:

      pH < 7.00

    • Slow pH decrease

    Less common than strong base-weak acid

    (fewer appropriate indicators)


    Titration curves polyprotic acids

    Titration Curves:Polyprotic Acids


    Salts

    Salts

    H2O

    • soluble

    NaCl (s)

    Na+ (aq) + Cl- (aq)

    • “slightly soluble”

    • Equilibrium between solid and dissolved ions

    H2O

    PbSO4 (s)

    Pb2+ (aq) + SO42- (aq)

    Ion-product expression

    Solubility product

    Solubility-Product Constant

    (at saturation)

    larger Ksp: more dissolution at equil. (saturation)

    Smaller Ksp: less dissolution at equil. (saturation)


    Insoluble metal sulfides

    Insoluble Metal Sulfides

    H2O

    MnS (s)

    Mn2+ (aq) + S2- (aq)

    S2- (aq) + H2O (l)

    HS- (aq) + OH- (aq)

    MnS (s)

    + H2O (l)

    Mn2+ (aq) +

    HS- (aq) + OH- (aq)


    3 minute practice

    3-minute Practice

    Write Ksp expression for each of the following:

    Silver bromide in H2O

    H2O

    AgBr (s)

    Ag+ (aq) + Br - (aq)

    Silver sulfide in H2O

    Ag2S (s)

    + H2O (l)

    2Ag+ (aq) +

    HS- (aq) + OH- (aq)


    Higher k sp greater solubility

    2-minute Practice

    Higher Ksp = greater solubility?

    Yes, for compounds with same total number of ions


    What else affects solubility

    What else affects solubility?

    • Presence of a common ion:

    Decreases solubility

    H2O

    PbSO4 (s)

    Pb2+ (aq) + SO42- (aq)

    Add Na2SO4?

    • pH:

    ↑ [H3O+] ↑ solubility

    if compound contains anion of weak acid

    H2O

    CaCO3 (s)

    Ca2+ (aq) + CO32- (aq)

    CO32- (aq) + H3O+ (aq)

    H2O (l) + HCO3- (aq)


    Homework problems

    Homework problems

    Chap 19: #9, 13, 19, 29, 50, 63, 70, 76, 78, 90

    Due Tuesday, 4/28

    More lecture notes will be added next week!

    Stay tuned.


    Precipitation

    Precipitation

    Will it occur?

    • Qsp = Ksp:

    • Qsp > Ksp:

    • Qsp < Ksp:

  • Selective precipitation

    • Way to separate ions

    • Form slightly soluble compounds with different Ksp

  • Saturated solution

    Precipitation occurs

    Unsaturated solution


    Selective precipitation

    Selective Precipitation

    Mix 0.2 M Zn(NO3)2 and 0.4 M Mn(NO3)2.

    Precipitate?

    Add NaOH…

    Products?

    Zn(OH)2 and Mn(OH)2

    3-minute Practice

    Ksp Zn(OH)2 = 3.0 x 10-16

    KspMn(OH)2 = 1.6 x 10-13

    Which product is more soluble?

    What [OH-] would need to make a saturated solution of the more soluble product? Hint: use Ksp expression!


    Complex ions

    Complex Ions

    • Central metal ion + ligands

      Ionic ligands:

    Lewis base

    Lewis acid

    OH-, CN-, halides

    Molecular ligands:

    H2O, NH3

    M(H2O)42+ (aq) + 4 NH3(aq)

    M(NH3)42+ (aq) + 4 H2O(l)

    Formation constant:


    Effects of ligands

    Effects of ligands

    A slightly soluble compound becomes more soluble when its cation forms a complex ion.

    AgBr(s)

    Ag+ (aq) + Br- (aq)

    Add Na2S2O3:

    Ag+ (aq) + S2O32- (aq)

    2

    Ag(S2O3)23-(aq)

    Amphoteric Hydroxides:

    • Very slightly soluble in water

    • More soluble in acidic or basic solutions

    Al(OH)3 (s)

    + 3H3O+

    Al3+(aq) + 6 H2O (l)

    Al(H2O)6 (s)

    + 4 OH-

    Al(H2O)2(OH)4- (aq)

    + 4 H2O (l)


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