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pH titration curves PowerPoint PPT Presentation


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pH titration curves. pH measured as an acid is neutralized by a base follows a characteristic curve that enables the equivalence-point to be determined with precision. Constructing curves. Shapes of curves will depend on the strength of the acid and base

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pH titration curves

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Ph titration curves l.jpg

pH titration curves

  • pH measured as an acid is neutralized by a base follows a characteristic curve that enables the equivalence-point to be determined with precision


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Constructing curves

  • Shapes of curves will depend on the strength of the acid and base

  • Construct curve by calculating [H3O+] as the base (acid) is added.

  • The strong-strong case is easy: all species are always completely ionized. At the equivalence point, pH = 7

  • Weak acid - strong base will use pKa for the acid


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Titration of weak acid with strong base in four zones

  • Zone 1: Initial pH of acid with no base added

  • Zone 2: Addition of OH- up to equivalence point

  • Zone 3: At equivalence point: all the HAc is converted into Ac-

  • Zone 4: After equivalence point


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Constructing a pH curve for the weak acid - strong base case

  • Zone 1: Initial pH of acid with no base added

  • Zone 2: Addition of OH- up to equivalence point

    • Treat like a buffer solution

    • NOTE: Concentrations change as volume increases with addition of base!


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Explicit (contains adult themes) calculation for zone 2

  • Determination of [H+] up to equivalence point

  • [HAc]o is initial concentration of HAc and VH is the initial volume of HAc used

  • [OH-] is initial concentration of the base and VOH is the volume added at that point in the titration


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Weak acid – strong base continued

  • Zone 3: At equivalence point: all the HAc is converted into Ac-

    • Treat like solution of basic salt

    • Where [Ac-] is the initial acetate concentration at the neutralization point prior to dissociation


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Weak acid – strong base continued

  • Zone 4: After equivalence point

    • All OH- results from the excess base


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pH curves assist in identifying suitable indicator

  • The equivalence point with the weak acid is at pH>7

  • Above pH 7, both curves coincide (strong base controls pH)

  • The initial rise in pH is greater with the weak acid (but at much lower [H3O+]


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End-point detection gets harder with weaker acids

  • Initial pH is higher

  • Initial change in pH is greater

  • Change in pH at equivalence point is lower

  • Harder to detect equivalence point in weak acid


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Weak base - strong acid

  • Analogous to the weak acid – strong base case

  • pH at equivalence point < 7

  • Dictates use of different indicator

  • pH after equivalence point controlled by strong acid


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Polyprotic acids – lots to note

  • An amino acid has two dissociations:

    H2A+ + H2O = HA + H3O+

    HA + H2O = A- + H3O+


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Example for alanine

  • Initially, pH determined by pKa1

  • Halfway to 1st equivalence pt pH = pKa1

  • 1st equivalence pt,

  • Halfway to 2nd equivalence pt, pH = pKa2

  • 2nd equivalence pt, pH determined by pKb for the base A- (Kb obtained from Ka2)

  • Beyond 2nd equivalence pt, pH determined from OH- from NaOH


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Solubility products – equilibrium between solute and solid

  • An electrolyte completely dissociated in equilibrium with undissolved solid

  • The solid phase is ignored

  • Writing Ksp expressions for salts


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Values of Ksp at 298 K


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Calculations

  • Determine Ksp from solution concentrations (concentrations of individual ions may not be equal to those you would get by simple dissociation of compound)

  • Determine solubility from Ksp

    MgF2 = Mg2+ + 2F-

    Ksp = x.(2x)2 x =[Mg2+]


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Factors affecting solubility

  • Common ion effect – the addition of an ion from another source

  • Solution pH – where there is a weak acid or base

  • Complex ion formation

  • Amphotericity


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Common ion effect

  • In a solution of a salt AB, addition of a additional B ions from another source will cause [A] to decrease because of Ksp


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Calculation of solubility under these conditions

  • Calculate solubility of MgF2 in a solution of 0.1 M NaF(aq)?

    • At equilibrium, [Mg2+] = x, [F-] = 0.1 + 2x


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Solution pH

  • Basic anions are protonated in acid

    • CaCO3 is insoluble in pure water

    • In acid solution, H+ converts CO32- to HCO3-

    • More Ca2+ is drawn into solution (Le Chatelier)

  • Salts of basic anions increase solubility in acid conditions

  • pH does not affect anions that are not basic


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Complex ion formation affects solubility

  • AgCl is ordinarily highly insoluble (test for chloride ions)

  • Addition of aqueous ammonia redissolves the precipitate by formation of the complex ion Ag(NH3)2+


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  • Two stage process of formation of highly favoured complex

  • Equilibrium lies to right


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Amphoteric substances

  • Dissolve in both acid and basic solutions

  • Examples of complex ion formation

  • Oxides and hydroxide of Al are examples


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Prediction of precipitation and ion products

  • Predict the formation of a precipitate when solutions are mixed

  • Ion product is not an equilibrium quantity

    IP = [Ca2+][F-]2

  • If IP is found to be larger than Ksp then precipitation occurs


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Selective precipitation

  • Mixtures of ions can be separated by combining with an anion that has wide range of solubiity.

    • For example, sulphides of Zn, Pb, Cu and Hg will precipitate leaving the sulphides of Mn, Fe, Co and Ni in solution

    • Ksp for former group are much larger than for latter group


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Qualitative analysis

  • Apply a sequence of precipitation steps to divide a group of many metal ions into smaller groups. These smaller groups will be further analyzed to identify the members therein


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