Titration Technique

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Titration Technique

• http//wps.prenhall.com/wps/media/objects/3312/339
  2202/blb1703.html

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• Sites
• http//chemed.chem.wisc.edu/chempaths/GenChem-Text
  book/Titrations-875.html
• IB http//ibchem.com/IB/ibnotes/full/aab_htm/18.5
  .htm
• Simulation
• http//www.avogadro.co.uk/chemeqm/acidbase/titrati
  on/phcurves.htm
• http//www.ausetute.com.au/titrcurv.html

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• A titration is a volumetric technique in which a
  solution of one reactant (the titrant) is added
  to a solution of a second reactant (the
  "analyte") until the equivalence point is
  reached.
• The equivalence point is the point at which
  titrant has been added in exactly the right
  quantity to react stoichiometrically with the
  analyte.
• An indicator may be added which has an "endpoint"
  (changes color) at the equivalence point, or the
  equivalence point may be determined from a
  titration curve.

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Neutralization of Acids and Bases
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• In an acid-base titration, a solution containing
  a known concentration of base is slowly added to
  an acid (or the acid is added to the base).
• Acid-base indicators can be used to signal the
  equivalence point of a titration (the point at
  which stoichiometrically equivalent quantities of
  acid and base have been brought together).

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The titrant is added to the solution from a buret, and the pH is continually monitored using a pH meter. To understand why titration curves have certain characteristic shapes, we will examine the curves for three kinds of titrations (1) strong acid-strong base (2) weak acid-strong base
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• a pH meter can be used to monitor the progress of
  the reaction producing a pH titration curve, a
  graph of the pH as a function of the volume of
  the added titrant.
• The shape of the titration curve makes it
  possible to determine the equivalence point in
  the titration.
• The titration curve can also be used to select
  suitable indicators and to determine the Ka of
  the weak acid or the Kb of the weak base being
  titrated.

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Terminology

• Titration the progressive transfer of a
  solution from a buret (called the titrant) into a
  measured volume of another solution (called the
  sample).
• Equivalence point the volume of titrant
  required to neutralize the sample ( mol acid
  mol base).
• Endpoint the pH at the equivalence point of a
  titration, where the indicator changes color.
• Indicator a chemical which is added to the
  sample that changes colour at the equivalence
  point of a titration.
• Buffering region a horizontal region of the pH
  curve where pH is not changing significantly.

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• The endpoint of a titration is NOT the same thing
  as the equivalence point
• The equivalence point is a single point defined
  by the reaction stoichiometry as the point at
  which the base (or acid) added exactly
  neutralizes the acid (or base) being titrated.
• The endpoint is defined by the choice of
  indicator as the point at which the colour
  changes. Depending on how quickly the colour
  changes, the endpoint can occur almost
  instantaneously or be quite wide.

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• If an appropriate indicator is chosen such that
  the endpoint of the titration occurs at the
  equivalence point, then a colour change in the
  solution being titrated can be used as a signal
  that the equivalence point has been reached.

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• Intuition may suggest that the endpoint of the
  titration will occur at the equivalence point if
  we choose an indicator whose pKa is equal to the
  pH of the equivalence point. If such an indicator
  was chosen, the colour change would be half
  complete at the equivalence point.
• Thus for titrating a weak acid with a strong base
  where, at equivalence point the solution is
  slightly acidic, an optimum indicator should have
  a pKa gt 7, for example Thymol Blue (pKa 8.9) or
  Phenolphthalein (pKa 9.4)
• For titrating a weak base with a strong acid at
  equivalence point encounters a slightly basic
  solution. Thus an optimum indicator should have a
  pKa lt 7, for example Bromocresol Green (pKa
  4.7) or Methyl Red (pKa 5.1)
• For titration involving strong base and a strong
  acid an optimum indicator should have a pKa 7,
  for example Bromothymol Blue (pKa 7.0) or
  Phenol Red (pKa 7.9

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Strong Acid-Strong Base Titrations

• 1. Consider the titration when 0.100 M NaOH
  solution have been added to 50.00 mL of 0.100 M
  HCl.

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• The initial pHThe pH of the solution before the
  addition of any base is determined by the initial
  concentration of the strong acid.
• For a solution of 0.100 M HCl
• H 0.100 M
• pH -log(0.100) 1.00

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• Between the initial pH and the equivalence point
  ( adding 49 mL of NaOH)
• As NaOH is added, the pH increases slowly at
  first and then rapidly in the vicinity of the
  equivalence point.
• The pH of the solution before the equivalence
  point is determined by the concentration of acid
  that has not yet been neutralized.
• nH 0.1 x 0.05 5 x 10-3 nOH- 0.1 x
  0.049 4.9 x 10-3
• nH left 5 x 10-3 4.9 x 10-3 1 x 10-4
  mol H
• Vsolution 50 49 99 mL 0.099L
• H 1x10-4 / 0.099 0.001M
• gt pH 3

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• The equivalence pointAt the equivalence point
  equal number of moles of the NaOH and HCl have
  reacted, leaving only a solution of their salt,
  NaCl. No calculation is required to deduce that
  the pH is 7.00, because the cation of a strong
  base (in this case Na) and the anion of a strong
  acid (in this case Cl) do not hydrolyze and
  therefore have no appreciable effect on pH.

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• After the equivalence pointThe pH of the
  solution after the equivalence point is
  determined by the concentration of the excess
  NaOH in the solution.(adding 51 mL of NaOH)
• V50 51 101 mL 0.101 L
• nH 5 x 10-3 nOH- 0.1 x 0.051 5.1 x
  10-3
• nOH- left 5.1 x 10-3 5.0 x 10-3 1 x
  10-4 mol OH-
• OH- 1x10-4 / 0.101 9.9 x 10-4 M
• pOH 3.00
• pH 11.00

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• 3. What is the pH when 48.00 ml of .100 M NaOH
  solution have been added to 50.00 ml of .100 M
  HCl solution?

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GO TO INDICATORS!!!!!
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Weak Acid Strong Base Titration

• As an example, consider the titration curve for
  the titration of 50.0 mL of 0.100 M acetic acid,
  HC2H3O2, with 0.100 M NaOH.
• The initial pHThis pH is just the pH of the
  0.100 M HC2H3O2.
• The calculated pH of
• 0.100 M HC2H3O2 is 2.89
• (using Ka)

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• Between the initial pH and the equivalence
  pointTo determine pH in this range, we must
  consider the neutralization of the acid
• Calculate the pH of the solution formed when 45.0
  mL of 0.100 M NaOH is added to 50.0 mL of 0.100 M
  HC2H3O2 (Ka 1.8 105). Consider the
  neutralization of the acid.
• Prior to reaching the equivalence point, part of
  the HC2H3O2 is neutralized to form C2H3O2. Thus,
  the solution contains a mixture of HC2H3O2 and
  C2H3O2.
• pH 5.70

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At the equivalence point

• The equivalence point is reached after adding
  50.0 mL of 0.100 M NaOH to the 50.0 mL of
  0.100 M HC2H3O2. At this point the 5.00  103 mol
  of NaOH completely reacts with the 5.00  103 mol
  of HC2H3O2 to form 5.00  103 mol of their salt,
  NaC2H3O2.
• The Na ion of this salt has no significant
  effect on the pH.
• The C2H3O2 ion, however, is a weak base, and the
  pH at the equivalence point is therefore greater
  than 7.

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• Finding the concentration of the salt
• na 0.005
• nsalt 0.005 n C2H3O2-
• V 50 50 100 mL 0.100 L
• Because the salt is a weak base
• Kb Kw/Ka OH- 5.3 x 10-6 pH 8.72