Ammonium Ion by malj



        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 which 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 on 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
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
   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 equations 8 and 9.
   % NH  = mol NH  x FW (NH  ) x 100/mass of sample
        4          4          4                                                           (8)

% NH  = {(VOH- x MOH-) – (VH+ x MH+)} x 18.038 x 100 / sample mass, grams
     4                                                                                   (9)

Safety Precautions

Substance                            Precaution                     Disposal

HYDROCHLORIC ACID                    Strong acid. Corrosive.        Down drain with a lot of
                                     Causes severe skin burns.      water.
                                     Inhalation of vapors           Absorb spills with sand
                                     or mist causes irritation.     sand and soda.
METHYL RED                           Toxic if ingested.             Down Drain.
SODIUM HYDROXIDE                     Corrosive to body tissue.   Down drain with a lot of
                                     Causes burning and itching.  water.

AMMONIUM SALTS                       Corrosive. Release             Down drain.
{Your unknowns}                      ammonia vapors in
                                     strong base solutions.

(Use your time wisely! Start by weighing out three portions of your sample, adding the 25.00 mL
  of NaOH(aq) to them, and geting them heated, BEFORE you start standardizing the HCl(aq))
                     Preparation and standardization of 0.1M hydrochloric acid.
                           {This part will be done with your partner.}
 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--agreement between determinations
should be better than 1.0% relative.

        Preparation of unknown solution. {Each student will do his/her own unknown.}
 Obtain a solid sample for analysis and record its number.
 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: see equation 9. Agreement
between determinations should be within 1.0% relative.

    Your report should include the following, in the order shown:
 A brief, descriptive title, that assumes that the reader knows what your assignment was;
 a brief explanation of the procedure, including evidence that you know the proper handling of
the reagents and equipment that was used;
 your sample number, your three (or more) separate values for the percent of NH4+ in your
sample, the average of these values, and the relative range for your determinations;
 discussion of any errors you committed during the experiment, and how they affected your
results. This section is NOT for speculation, cant, opinion, or anything other than things you
actually did, that might have made a difference;
 your name, the date of submission, and your signature at the end.


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