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ANALYSIS OF AN UNKNOWN CHLORIDE*
One of the important applications of precipitation reactions lies in the area of quantitative
analysis. Many substances that can be precipitated from solution are so slightly soluble that the
precipitation reaction by which they are formed can be considered to proceed to completion. Silver
chloride is an example of such a substance. If a solution containing Ag+ ion is slowly added to one
containing Cl- ion, the ions will react to form AgCl:
Ag+(aq) + Cl-(aq) → AgCl(s)
Silver chloride is so insoluble that essentially all of the Ag+ added will precipitate as AgCl
until all of the Cl- is used up. When the amount of Ag+ added to the solution is equal to the amount
of Cl- initially present, the precipitation of Cl- ions will be, for all practical purposes, complete.
A convenient method for chloride analysis using AgCl has been devised. A solution of
AgNO3 is added to a chloride solution just to the point where the number of moles of Ag+ added
is equal to the number of moles of Cl- initially present. You can analyze for Cl- by simply
measuring how many moles of AgNO3 are required. This measurement is rather easily made by an
experimental procedure called a titration.
In this titration, a solution of AgNO3 of known concentration (in moles AgNO3 per liter of
solution) is added from a calibrated buret to a solution with a measured amount of unknown. The
titration is stopped when a color change occurs in the solution, indicating that stoichiometrically
equivalent amounts of Ag+ and Cl- are present. The color change is caused by a chemical reagent,
called an indicator, which is added to the solution at the beginning of the titration.
The volume of AgNO3 solution that has been added up to the time of the color change can
be measured accurately with the buret, and the number of moles of Ag+ added can be calculated
from the known concentration of the solution.
In the Mohr method for the volumetric analysis of chloride, which you will employ in this
experiment, the indicator used is K2CrO4. The chromate ion present in solutions of this substance
will react with silver ion to form a red precipitate of Ag2CrO4. Under the conditions of the titration,
the Ag+ added to the solution reacts preferentially with Cl- until that ion is essentially quantitatively
removed from the system, at which point Ag2CrO4 begins to precipitate and the solution color
changes from yellow to a reddish hue. The end point of the titration is that point at which the color
change is first observed.
In this experiment, weighed samples containing an unknown percentage of chloride will
be titrated with a standardized solution of AgNO3, and the volumes of AgNO3 solution required to
reach the end point of each titration will be measured. From the known molarity of the solution,
MAgNO3, and the measured volume used, VAgNO3, the number of moles of Ag+ added can be
MAgNO3 x VAgNO3 = no. of moles AgNO3 = no. of moles Ag+
*Adapted from Slowinski, E. J., Wolsey, W. C. Chemical Principles in the Laboratory 9th ed.
At the end point of the titration,
no. of moles Ag+ added = no. of moles Cl- present in unknown
no. of grams Cl- present = no. of moles Cl- present x molar mass Cl
% Cl =
no. of grams Clno. of grams unknown
Obtain an unknown chloride from the instructor. Weigh out three samples of the chloride
into weigh trays, each sample weighing about 0.2 grams. Record each mass and make sure not to
mix up which sample is which. Place the first sample into a 250-mL Erlenmeyer flask. Rinse the
weigh tray with distilled water to ensure complete transfer of the sample, and fill the flask up to
50 mL to fully dissolve the sample. Add 6 drops of 0.5 M K2CrO4 indicator solution.
Using the graduated cylinder, measure out about 100 mL of the standardized AgNO 3
solution into a clean dry 125-mL Erlenmeyer flask. This will be your total supply for the entire
experiment so do not waste it. Make sure that the stopcock of the buret is closed, then pour 1-2 mL
of the AgNO3 solution into the buret and tip it back and forth to rinse the inside walls. Allow the
AgNO3 solution to drain out the buret tip completely into a small beaker, then repeat this process
two more times. After the third rinse, fill the buret with the AgNO3 solution. The fill level should
be close to the top of the graduation marks on the buret; do not attempt to fill it to exactly the “0.0
mL” reading. Open the buret stopcock momentarily to flush any air bubbles out of the tip of the
buret. At this time, make sure your stopcock fits snugly and that the buret does not leak.
Read the initial buret level to 0.01 mL. You may find it useful when making readings to
put a white card marked with a thick black stripe behind the meniscus. If the black fine is held just
below the level to be read, its reflection in the surface of the meniscus will help you obtain an
accurate reading. Begin to add the AgNO3 solution to the chloride solution in the Erlenmeyer flask.
A white precipitate of AgCl will form immediately, and the amount will increase during the course
of the titration. At the beginning of the titration, you can add the AgNO 3 fairly rapidly, a few
milliliters at a time, swirling the flask as best you can to mix the solution. You will find that at the
point where the AgNO3 hits the solution, there will be a red spot of Ag2CrO4, which disappears
when you stop adding the silver nitrate and swirl the flask. As you proceed with the titration, the
red spot will persist more and more, since the amount of excess chloride ion, which reacts with the
Ag2CrO4 to form AgCl, will slowly decrease. Gradually decrease the rate at which you add AgNO3
as the red color becomes stronger. At some stage you may find it convenient to set your buret
stopcock to deliver AgNO3 slowly, drop by drop, while you swirl the flask. When you are near the
end point, add the AgNO3 drop by drop, swirling between drops. The end point of the titration is
that point where the mixture first takes on a permanent reddish-yellow color that does not revert
to pure yellow on swirling. If you are careful, you can hit the end point within 1 drop of AgNO 3.
When you have reached the end point, stop the titration and record the buret level. If you are
uncertain if you have reached the end point, note the buret level before adding one additional drop;
this ensures that if the additional drop takes you past your end point, you still have the data from
before it was added.
Pour the solution you have just titrated into another 250-mL Erlenmeyer flask. Use the
color of this mixture as a reference against which you compare your samples in the remaining
Rinse out the 250-mL Erlenmeyer flask in which you carried out the titration. Take your
second sample and carefully pour it from the weigh tray into the Erlenmeyer flask. Wash the tray
a few times with distilled water to complete the transfer of the sample into the flask. Add water to
the flask to a volume of about 50 mL and swirl to dissolve the solid. Refill your buret if necessary,
take a volume reading, add the indicator, and proceed to titrate to an end point as before. This
titration should be more accurate than the first, since the volume of AgNO3 used is proportional to
sample size and therefore can be estimated rather well on the basis of the relative masses of the
two samples. In addition, you have a reference for color comparison that should make it easier to
recognize when a color change has occurred.
Titrate the third sample as you did the second. With care it should be possible to obtain
volume-mass ratios that agree to within less than 1% in the last two titrations.
CHEM 403 – Exp 7
Name:______________________________ Section: _________
Data & Calculations
Molarity of standard AgNO3 solution: ____________ M
Mass of sample
Initial buret reading
Final buret reading
Volume of AgNO3 used
to titrate sample
No. of moles of AgNO3
used to titrate sample
No. of moles of Clpresent in sample
Mass of Clpresent in sample
Percentage of Clin sample
Mean value of percentage of Cl- in unknown ___________ %
Unknown number ______
CHEM 403 – Exp 7
Name:______________________________ Section: _________
1. A sample containing 0.221 g Cl- is dissolved in 50.0 mL water.
a. How many moles of Cl- ion are in the solution?
__________ moles Clb. What is the molarity of the Cl- ion in the solution? (M = moles per liter)
2. A solid chloride sample weighing 0.3147 g required 43.75 mL of 0.05273 M AgNO3 to reach
the Ag2CrO4 end point.
a. How many moles Cl- ion were present in the sample? (Use Equations 2 and 3.)
__________ moles Clb. How many grams Cl- ion were present? (Use Equation 4.)
__________ g Clc. What was the mass percent Cl- ion in the sample? (Use Equation 5.)
__________ % Cl3. How would the following errors affect the mass percent Cl- obtained in Question 2c? Give your
reasoning in each case.
a. The student did not rinse their buret with AgNO3 solution before filling it with that solution.
b. The student was past the end point of the titration when they took the final buret reading.
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