Ammonium salt

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					        EXPERIMENT 8: AMMONIUM ION CONTENT IN AN AMMONIUM SALT
                                           INTRODUCTION
        If one attempts to titrate an extremely weak acid with sodium hydroxide, the endpoint is
difficult to determine (see “Silberberg,” section 19.2, for a general discussion of acid-base
titrations). As you can see from Figure 19.6, page 817 of “Silberberg,” the titration of a strong
acid by a strong base, such as sodium hydroxide, produces a very sharp endpoint. This means
that a very small volume change gives a large change in pH at the endpoint, i.e. the slope of the
titration curve is very steep. However, if the acid being analyzed is weaker, this slope becomes
smaller (figure 19.7, p. 819). If the slope becomes so small that one drop of titrant (0.05 mL)
produces less than about 2 pH units of change, then the titration is not practical, because most
indicators require a change of about 2 pH units to cause their color to change. If the complete
color change requires more than one drop, the experimental error in your seeing the endpoint is
unacceptably large.
     Ammonium ion is an example of an acid that is too weak to be titrated directly. However,
such a system can be analyzed by a method known as a "back titration." In this method, an
excess of one reagent is added to react with the analyte. Then something is done which will
prevent this reaction from being reversed at the next step. Finally, the excess reagent from step
one is titrated by a second reagent. In the case of ammonium ion analysis, an excess of standard
sodium hydroxide is added to react with the ammonium ion.
      NH4+ (aq) + OH-(aq)  NH3(aq) + H2O(l)                                                  (1)
     This solution is then brought to a boil to drive off the ammonia so that reaction 1 cannot be
reversed.
     NH3(aq)  NH3(g)                                                                 (2)
    Now the excess hydroxide ion can be titrated with standard hydrochloric acid, using methyl
red as an indicator. This step is called the “back titration.”
   OH-(aq) + H+(aq)  H2O(l)                                                             (3)
    Since all of the hydroxide has reacted in either step one or step three, we can use this system
to determine the amount of ammonium ion present.
    total mol OH- = mol OH- (rxn 1) + mol OH- (rxn 3)                                         (4)
   The moles of OH- used in reaction 1 is equal to the moles of ammonium ion present; the
moles of OH- used in reaction 3 is equal to the moles of acid added in the back titration.
Therefore, equation 4 can be rewritten as
                                           +
   total mol OH- = (mol NH  ) + (mol H )
                              4                                                            (5)
   Since the hydroxide and the acid were added as standard solutions and since the volumes
were added in carefully measured amounts, the moles of each are given by the volume x the
molarity. These can be substituted into equation 5 to give equation 6.
   (VOH- x M OH-) = (mol NH  ) + (VH+ x MH+)
                                4                                                       (6)
Equation 6 can then be rearranged to give the mol of ammonium ion.
   mol NH  = (VOH- x MOH-) - (VH+ x MH+)
            4                                                                        (7)
    This can then be used to calculate the percent of ammonium ion in the unknown salt that you
will be given, as shown in equation 8.
    % NH  = mol NH  x FW (NH  ) x 100/mass of sample
           4           4              4                                                  (8)




                              Ammonium Ion Content Page 1 of 3
                                       PROCEDURE
(Use your time wisely! Start by weighing out three portions of your sample, adding the 25.00 mL
  of NaOH(aq) to them, and getting them heated, BEFORE you start standardizing the HCl(aq))

              Part 1: Preparation and standardization of 0.1M hydrochloric acid
 Add 4-5 mL of concentrated hydrochloric acid to 500 mL of de-ionized water. Transfer to a
clean brown plastic bottle and mix well by shaking.
       You will use the standardized sodium hydroxide solution from the KHP experiment to
  standardize this hydrochloric acid solution: you will need 100-150 mL of it. Be sure to copy
                          down the NaOH molarity into your lab notebook.
 Rinse a 25 mL volumetric pipet with a small amount of NaOH solution and discard the rinse
solution. If drops of solution hang up in the pipet, the pipet is not clean enough for this
application. Consult with your instructor about how to clean it.
 Transfer a 25.00 mL portion of standard NaOH solution into each of three 250 mL
Erlenmeyer flasks.
 Clean, rinse and fill a buret with the hydrochloric acid solution.
 Add 6 drops of methyl red indicator solution to each flask that contains the 25.00 mL of
sodium hydroxide solution, and titrate each to the endpoint, using the hydrochloric acid solution
in the buret. The endpoint is evident when the indicator changes from yellow to red.
 Calculate the molarity of the hydrochloric acid solution—the standard deviation should be
better than 0.001M.

                     Part 2: Preparation and Analysis of unknown solution
 Obtain a solid sample for analysis and record its number (or letter).
 Clean three 250 mL Erlenmeyer flasks. Into each, accurately weigh 0.11 - 0.12 g of
unknown. Add 50 mL of deionized water to each and swirl to dissolve the unknown.
 Pipet 25.00 mL of standard sodium hydroxide into each.
 Bring these solution to a gentle boil, in the hood. Boil ten minutes, and check the steam
vapor from each flask with wet, red litmus paper. Continue the boiling until the litmus no longer
turns from red to blue in the steam, or the smell of ammonia is no longer evident. If necessary,
add some de-ionized water to the flask to prevent it from going dry.
 Cool each flask to room temperature, add 6 drops of methyl red indicator solution and
titrate with your standard hydrochloric acid as before.
 Calculate the percent of ammonium ion for each sample. Agreement between determinations
should have a standard deviation of less than 0.1%. Compare your value to the theoretical
%NH4+ in the following reagents to determine the identity of you unknown: NH4Br, NH4Cl,
NH4I, or NH4NO3. Be sure to calculate absolute % error based on which unknown you believe
you have.




                             Ammonium Ion Content Page 2 of 3
                                           DISCUSSION
                            Your discussion should have the following:
   your sample number, your average value for the mass % NH4+ in your sample and the error
    analysis values for your determinations;
   state which unknown you believe your sample is and how your data leads you to that
    conclusion;
   discussion of any errors you committed during all the parts of the experiment, and how they
    affected your results;
   discussion of errors associated with the experiment itself and how they would affect your
    results;
   your name, the date of submission, and your signature at the end.




                             Ammonium Ion Content Page 3 of 3

				
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