D17 _Acid-Base Titration_ by liamei12345


									                               Acid-Base Titration
    In a titration, the concentration of one solution is determined by quantitatively
     observing its reaction with a standard solution (ie. a solution of known concentration).

    The observations can be used to standardize the solution (ie. determine its unknown

    The predicted yield is calculated on the assumption that all the limiting reactant reacts
     to make product on the ratio described by the balanced equation.

    A typical titration set-up is shown here.


                                                                      burette clamp

             titrant                                                   retort stand

                                                          stopcock     valve

    Erlenmeyer flask

    The solution that is placed into the burette is known as the titrant. By measuring the
     initial burette volume (prior to beginning the titration) and the final burette volume (at
     the completion of the titration), the volume of titrant required to complete the reaction
     can be determined.
   The solution that is placed into the Erlenmeyer flask is known as the aliquot. The
    volume of the aliquot is pre-determined before the reaction begins - a volumetric
    pipette is used to measure the aliquot.

   Either the aliquot or the titrant can be the standard solution.

   The stage of the titration at which the reaction is complete is called the equivalence
    point. At this point, stoichiometrically equivalent amounts of each reactant have been

   In acid-base titrations, an acid-base indicator is added to the aliquot to provide visual
    evidence of the end of the reaction. A dramatic colour change of the indicator
    identifies when the reaction is complete.

   The point at which the indicator changes colour is called the endpoint.

                            The Equivalence Point
   In an acid-base titration, an acid titrant is added to a base aliquot, or vice versa.

   For monoprotic acids and bases, the point at which equal moles of reactant acid and
    base combine is called the equivalence point of the titration.

   For example, the titration of sodium hydroxide with hydrochloric acid.

                       HCl(aq) + NaOH(aq)            H2O(l) + NaCl(aq)

   The mole ratio is 1:1. Therefore, the equivalence point occurs when an equal amount
    of HCl(aq) has been added to the NaOH(aq).

   In every reaction between a strong monoprotic acid and a strong base, the equivalence
    point has a pH of 7 because all hydronium ions from the acid have been neutralized
    by an equal amount of hydroxide ions from the base.

   Acid-base titrations are performed with repeated trials until at least 3 concordant
    results are obtained. A concordant result means the titrant volumes required to reach
    the equivalence point are within a range of 0.2 mL.

   Most neutralization reactions involve colourless solutions with no obvious visible
    evidence that a reaction is taking place.

   An acid-base indicator is a substance that changes colour over a given pH range.
   Usually, indicators are weak monoprotic acids. The molecular and ionized forms of
    the indicator have different colours.

                        HIn(aq) + H2O(l)         H3O+(aq) + In-(aq)

                      colour 1                              colour 2

   For example, bromothymol blue is a commonly used indicator for titrations. It is
    yellow between pH 0 and pH 6. It turns blue between pH 6 and pH 7.6.

   This indicator is commonly used for titrations between a strong monoprotic acid and
    a strong base.

                    The titration of 10.0 mL of hydrochloric acid with
                               0.16 mol/L sodium hydroxide
Trial #                           1        2         3          4         5            6
Initial Burette Volume
Final Burette Volume
Equivalence Point

Use this data to calculate [HCl(aq)].
Chem 20                                         Titrations

1. Find the molar concentration of the sodium carbonate solution, given the following:

         Titration of 25.0 mL of sodium carbonate with 0.352 mol/L hydrochloric acid
     Trial #                               1            2           3             4
     Initial buret reading (mL)           0.0         15.9         31.2          1.4
     Final buret reading (mL)            15.9         31.2         46.4         16.6
     Equivalence point (mL)

2. Find the molar concentration of the potassium hydroxide solution, given the

          Titration of 10.0 mL of potassium hydroxide with 0.150 mol/L sulfuric acid
     Trial #                                1          2            3              4
     Initial buret reading (mL)            0.2        12.8         25.3           0.6
     Final buret reading (mL)             12.8        25.3         37.8          13.3
     Equivalence point (mL)

3. Find the molar concentration of the potassium permanganate solution, given the

                Titration of 10.0 mL of acidified 0.100 mol/L iron(II) sulfate with
                                    potassium permanganate
     Trial #                                  1           2              3           4
     Initial buret reading (mL)              0.1         11.3          21.9         32.5
     Final buret reading (mL)               11.3         21.9          32.5         42.9
     Equivalence point (mL)

   The balanced equation that you will use for this reaction is:

 10 FeSO4(aq) + 2 KMnO4(aq) + 8 H2SO4(aq)  5 Fe2(SO4)3(aq) + K2SO4(aq) + 2 MnSO4(aq) + 8 H2O(l)

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