18 Ksp of Ca_IO3_2 by ashrafp


									Name: ________________                                           Date: ________________

Partner: _______________                                         Lab AP 18

                                 DETERMINATION OF A KSP
When a salt of low solubility dissolves in water, equilibrium is established between the
solid solute and the dissolved ions. There is a special term for the equilibrium constant,
the solubility product, Ksp for this equilibrium.
For a general salt, AmBn, the equation would be:
                              AmBn (s)  m A (aq) + n B (aq)
And the equilibrium constant (remember, pure substances, solid AmBn are not included
in equilibrium expressions) would be:
                                       Ksp = [A] m [B]n
The value of the constant is expressed without units. The magnitude of the constant
expresses the degree to which the salt will dissolve (the degree to which the equilibrium
lies to the right). As Ksp values are used for slightly soluble salts, they are quite small.
For example, the Ksp for AgCl is 4 x 10–11 at 10 °C.
In this laboratory you will determine the Ksp for the very slightly soluble salt calcium
iodate, Ca(IO3)2. The equation for the dissolving of the salt is,
                          Ca(IO3)2 (s)  Ca+2 (aq) + 2 IO3– (aq)
and the solubility product expression is,
                                     Ksp = [Ca+2] [IO3–]2
To determine the Ksp you could measure the equilibrium concentration of either Ca+2 or
IO3–. (Since the concentrations of the ions are stoichiometrically related, you only need
to find one to know both.)
In this experiment the concentration of the iodate ion will be determined through a redox
titration with sodium thiosulfate, using starch as an indicator. We will add an excess of
iodide ion (as KI(aq)), which is oxidized by the iodate ion, IO3– , in the presence of acid
to elemental iodine, I2. The ionic equations for the reactions are:
                           IO3– + 5 I– + 6 H+ ——> 3 I2 + 3 H2O
                             I2 + 2 S2O3–2 ——> S4O6–2 + 2 I–
In the presence of starch the I2 forms a dark blue complex. To determine the
concentration of iodate in the unknown solution you will add the starch indicator and
then slowly titrate with Na2S2O3 until the dark blue color disappears.

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      Wear safety goggles at all times.
      Treat all unknown chemicals as hazardous. If case of skin contact, wash the
       affected area immediately

Materials:                            Equipment:
      NaIO3                                 Magnetic stirrer and stir bar
      1.0M Ca(NO3)2 solution                250 mL flask (2)
      0.10M KNO3                            125 mL flask (several)
      0.05M Na2S2O3                         50 or 100 mL graduated cylinder
      KI                                    Filter funnel and filter paper
      Starch solution                       Buret
                                            10 mL pipet

Part I — Preparation of a Saturated Solution of Ca(IO3)2
Note: While you have idle time during Part I, start setting up for your titration, Part II.
   1. Mass 3.96 g of NaIO3 in a weighing boat and transfer the salt to a 250 mL
      Erlenmeyer flask. Use a graduated cylinder to add 100 mL of distilled water.
      Add a magnetic stir bar stir until the NaIO3 is completely dissolved.
   2. Using a graduated cylinder, slowly add 20.0 mL 0.50M Ca(NO3)2 solution while
      stirring. Stir the mixture completely for 5 minutes. A white precipitate of Ca(IO 3)2
      should form.
   3. Let the mixture stand a few minutes, then decant off and discard the supernatant
      liquid. (It’s okay to lose a little solid when decanting. At this stage we are just
      trying to have some clean solid.) Obtain about 200 mL of 0.10M KNO 3 in a clean
      flask. Wash the solid by adding 30 mL of 0.10M KNO3 and stirring thoroughly for
      5 minutes. Repeat the wash process for a total of three washes.
   4. Add 100 mL of the 0.1M KNO3 and stir for another 5 minutes. There must be
      solid left in the flask at this time. (At this stage we are trying to establish the
      equilibrium. We now want the solution. The solid is not important, only that
      some exists.)
   5. Set up a filter funnel with filter paper and decant/filter the suspension into a clean
      dry flask 125 mL Erlenmeyer flask. Most of the solid should remain in the original
      flask. Do not wash the solid. Save the filtrate (liquid) for the next step.

CE80A040-D948-477A-84AD-DDEF604DD13C.DOC                                              2
Part II — Determination of [IO3–] in the Saturated Solution of Ca(IO3)2
   1. Set up a buret and fill it with the standardized solution of approximately 0.05M
      Na2S2O3. Be sure to record the actual concentration in your laboratory notebook.
   2. Pipet 10.0 mL of the filtrate from Part I, above, into a clean Erlenmeyer flask. Add
      about 20 mL of distilled water to the flask.
   3. Weigh about 2 g of solid KI (a large excess) in a weigh boat and transfer into the
      solution. Add 20 drops of 2 M HCl and mix thoroughly. The solution should take
      on a brown color.
   4. Titrate the mixture with the Na2S2O3 until the brown color changes to yellow.
   5. When the solution is a light yellow, add several drops of starch to the solution,
      enough to turn the solution black.
   6. Continue titrating with the Na2S2O3 until the blue/black color just disappears.
   7. Repeat the titration a second time. If time permits, complete a third trial.

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   1. Using the known concentration of sodium thiosulfate, calculate the iodate
      concentration, [IO3–], in the filtrate. (Hint: There’s some stoichiometry involved
      here. Use the volume and concentration of sodium thiosulfate to calculate moles
      of S2O3-2. Then use the stoichiometry in the equations to determine the moles of
      I2 consumed. Then use the stoichiometry of the reaction to determine the moles
      of IO3- used to produce these moles of I2. This can all be done with one
      dimensional analysis string or as separate steps. Either way, clearly show your
      work. Finally, using the aliquot volume of your filtrate, calculate the
      concentration of the iodate ions.)
   2. For every IO3- ion, how many Ca+2 ions are in solution? From the [IO3–],
      calculate the concentration of [Ca+2]. Write out the Ksp expression and then
      determine the Ksp of Ca(IO3)2.


   1. Calculate the moles of Ca+2 from the Ca(NO3)2 used in the experiment.
      Calculate the moles of IO3- from the NaIO3 used. Based on these values,
      calculate he mass of Ca(IO3)2 produced in the lab.

   2. Calculate the theoretical concentration if iodate, [IO3–], from the literature value of

   3. Why do you think we need to go through several washings prior to taking the
      aliquot for the [IO3-] determinations? (Hint: how hard is it to hit a pH of exactly
      7.00 when doing a titration?)

Error analysis

   1. Look up the accepted value of the Ksp and compare your results with the
      expected value.
   2. Calculate the experimental error using both the value for Ksp and the value you
      calculated for the theoretical concentration of IO3- in #2 above. Which do you
      think is a better representation of the error in the lab? Explain why!
   3. Comment on sources of error and possible ways to improve the lab.

      Don’t forget you purpose and conclusions demonstrating the principals involved
      in the lab.

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Notes to self:
The washing of the solid and preparation of the saturated solution took longer than
expected. This could be done prior to the lab to allow more time for multiple titrations.
Perhaps do it as a demo during the prior day’s lecture and prepare enough for the entire

The sodium thiosulfate seems to be of high purity. Standardization of the solution
yielded exactly the concentration calculated from the mass and volume. Don’t bother
standardizing in the future.

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