EXPERIMENT III
DETERMINATION OF CHLORIDE:
GRAVIMETRIC AND VOLUMETRIC METHODS
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INTRODUCTION
For any given analyte, numerous analytical methods are available in the literature, each
having unique advantages and disadvantages. The two most common methods for the
determination of chloride ion are based on its quantitative reaction with silver ion to form
silver chloride. One method involves the isolation of the silver chloride precipitate by
filtration, determining its mass, and using stoichiometry to calculate the percent chloride
in the sample. This is termed Gravimetric Analysis. Alternately, the chloride ion may be
titrated with a standard solution of silver ion and the volume of titrant used to calculate
the percent chloride in the sample. This Volumetric Analysis is often described as a
precipitate-forming titration.
In this experiment you will obtain an unknown sample, containing chloride ion, and you
will carry out both gravimetric and volumetric procedures using the same sample. In a
perfect world, bothe sets of data should result in identical values for the percent of
chloride in the sample as both methods are “tried and true”, and both have similar
precision and accuracy (2-3 ppt). In reality, your results for the two determinations will
likely differ. Your challenge is to determine if the difference is significant, and, most
importantly, you must decide what value to report as the best estimate of the true
percentage of chloride in your sample. You should use a blend of common sense, good
lab technique, statistical tests, and, perhaps, further replication to enhance confidence or
help to identify outliers.
The theory and practice of the Gravimetric Determination of Chloride is described in
Part I of this experiment and the Volumetric Determination of Chloride is presented in
Part II.
PART I: THE GRAVIMETRIC DETERMINATION OF CHLORIDE
Gravimetric methods are among the most accurate and precise methods of quantitative
analysis. These advantages are counterbalanced by the fact that they are often time
consuming, require considerable attention to detail and are limited to sample size and
concentrations which yield a weighable quantity on a conventional analytical balance.
While many gravimetric methods have been at least partially replaced by newer, faster
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techniques, there remain several situations where the only suitable analytical techniques
are gravimetric. For this reason you should become experienced and competent with
gravimetric techniques, such as the gravimetric determination of chloride.
THEORY
Addition of a solution containing silver ion [usually an aqueous silver nitrate (AgNO3)
solution] to a dissolved sample containing chloride ion will quantitatively precipitate the
chloride ion as solid silver chloride (AgCl), due to the small solubility product constant
of silver chloride (ksp = 1.82 x 10-10). The precipitate may be isolated from solution by
filtration, the material dried, weighed and the percent chloride (% Cl-) calculated from the
stoichiometry of the reaction:
A g (aq) + C l (aq)
A g C l(s)
In addition to possible interferences common to all gravimetric methods, the presence of
any other halide will cause high results since they produce a silver halide which is even
more insoluble than silver chloride. Also, silver chloride is light sensitive and excessive
photodecomposition will produce erroneous results, according to the reaction:
AgC l(s)
Ag(s) +
1
2
Cl2 (g)
The precipitate becomes violet-purple, due to the presence of finely divided silver metal,
and results will be low. If silver ion (Ag+) is present, in addition to the above, the
following reaction will also occur as a result of the photodecomposition reaction:
hv
3Cl 2 (g) + 3H 2 O(l) + 5Ag  (aq) 
5AgCl(s)
+ ClO 3- (aq) + 6H  (aq)
If this reaction predominates, the results will be high. To minimize such errors, it is
recommended that unnecessary exposure to light be avoided.
PROCEDURE
Preparation of the Sample for Analysis
Obtain approximately 5 g of an unknown sample from your instructor in a clean, dry,
and properly labeled weighing bottle. Dry the unknown material in the microwave oven
using medium power for two minutes. Weigh accurately, after cooling in a desiccator for
approximately 30 minutes, triplicate samples of approximately 0.5 g into numbered 400mL beakers. Dissolve in deionized water and dilute each to approximately 150 mL
volume. Add approximately 0.5 mL of reagent-grade concentrated nitric acid (HNO3).
Preparation of Crucibles
Clean and suitability mark three sintered glass crucibles of medium (M) porosity. Your
instructor will outline for you the best method of cleaning the crucibles, as the method of
cleaning is dependent upon their past history. Then bring the crucibles to constant weight
in the microwave oven using medium power. The first heating should be at least 30
seconds; let the crucible cool, heat for a further 30 seconds, cool, then finally heat for 4
2
minutes. Re-heat for 2-minute periods until constant weight is obtained. The weight
should become constant to within ±0.3 mg.
Preparation of Precipitating Reagent
An aqueous solution of 0.5 M silver nitrate (AgNO3) is available in the laboratory.
Assuming the unknown to be pure sodium chloride, calculate the volume of the silver
nitrate solution required to just precipitate all chloride ion in each sample. Include 10%
excess in the final volume to compensate for the fact that the calculation is based on an
assumption that could be incorrect.
Precipitation
Add the calculated amount of silver nitrate solution slowly to each sample, preferably
from a buret, with adequate stirring, (remember each sample will require a different
volume of silver nitrate solution as each sample is of a different weight). Cover each
solution with a watch glass (remember to leave the stirring rod in the beaker), heat each
solution to nearly boiling (hot plate) to coagulate the silver chloride precipitate. Let each
solution cool and settle, then test for complete precipitation as demonstrated by the
instructor. Add more precipitating reagent, if necessary, and re-check for complete
precipitation. Allow each solution to digest (in the dark) at least 1-2 hours, or, if
necessary overnight.
Filtration
Filter each solution through the corresponding sintered glass crucible, adhering to the
guidelines for proper filtration, presented by the instructor. Be certain to "police" each
beaker to assure quantitative transfer. Wash each precipitate with a solution of 1:500
nitric acid/water (nitric acid is added to prevent peptization) until the washings give a
negative test for Ag+. After you have a negative test for Ag+, place the sintered glass
crucible in the filtration apparatus and pull a vacuum on it for several minutes. It is
important that you get the filtrate as dry as possible before drying in an oven.
NOTE:
Washings may be tested for their silver ion content by using the following
procedure. After you have filtered your precipitate and washed with an initial amount of
wash solution, carefully remove the filter crucible and touch the base to the surface of a
clean, dry watch glass. Several drops of the residual wash will transfer to the watch
glass. If this does not occur, add a small amount (1 or 2 mL) of wash solution to the filter
crucible and wait for gravity to cause the transfer of some of the wash solution. Add 2-3
drops of 12 M HCl to the wash solution on the watch glass. The formation of a cloudy
white solution indicates the presence of Ag+ in the wash solution. If this occurs,
reassemble and wash with ~10 mL of wash solution and retest. Repeat until a negative
test for Ag+ is obtained.
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Drying the Crucibles
Place the crucibles in a beaker covered with a ribbed watch glass (as with the empty
crucibles). Dry in the microwave oven at medium power for 30 seconds; let the crucible
cool, heat for a further 30 seconds and cool, then finally heat for 4 minutes, cool in a
desiccator and weigh. Reheat for 2-minute periods until constant weight is obtained.
Cleaning the Crucibles
After completion of the analysis, remove carefully, without damaging the sinter, the cake
of silver chloride from the crucible(s) and place the cake in the silver waste bottle. Place
the crucibles in a beaker and fill each crucible with concentrated aqueous ammonia
solution (Do this in a hood). Allow the crucibles to sit until all the silver chloride has
been dissolved by the ammonia. Rinse each crucible with several portions of deionized
water and dry in a microwave oven at medium power for 4 minutes or until dry.
CALCULATION OF RESULTS
Calculate your results as the % Cl- in your sample. Be certain to report each individual
value, the mean value, the absolute deviation of each value from the mean, the relative
average deviation in parts per thousand, the standard deviation and the confidence
interval at an appropriate confidence level. A relative average deviation of > 5 ppt (5
parts per thousand) should be considered unsatisfactory.
PART II: THE VOLUMETRIC DETERMINATION OF CHLORIDE
Precipitate-forming titrations are not common. However, they are widely used for
determination of the halides, chloride (Cl-), bromide (Br-) and iodide (I-), using silver
nitrate (AgNO3) as the titrant. In this part of the experiment, chloride will be determined
by titration with silver nitrate using dichlorofluorescein as an indicator. This is
commonly known as Fajan's Method.
THEORY
Chloride present in a sample is quantitatively insoluble in a solution containing excess of
silver ion:
Ag+(aq) + Cl- (aq) 
 AgCl(s)
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Ksp = 1.82 x 10-10
This reaction is the basis of a gravimetric procedure described in Part I of this
experiment. The same determination may be accomplished volumetrically if a standard
solution of Ag+ is available. The end point of the reaction may be determined through
the use of an adsorption indicator, dichlorofluorescein. Its function may be described as
follows:
If the reaction is run in neutral or basic solution some of the indicator will dissolve to
form the dichlorofluorescinate anion, which is represented as In-. Before the equivalence
point, with Ag+ as titrant, excess Cl- is present in solution. The excess Cl- is adsorbed
onto the precipitate particles formed and the indicator anion is repelled by the negativelycharged precipitate.
Ag+(aq) + 2Cl-(aq) 
 AgCl:Cl (s)
At the equivalence point, there is little or no excess Cl-, and just beyond the equivalence
point Ag+ is in excess and becomes the primary adsorbed ion. The charge on the
precipitate changes from negative to positive and the indicator anion is adsorbed.
+
AgCl:Cl-(s) + Ag+(aq) 
 AgCl:Ag (s) + Cl (aq)
+ AgCl:Ag+(s) + In-(aq) 
 AgCl:Ag In (s)
yellow
rose-pink
The color change is:
yellow  rose-pink
It is believed that the indicator anion (yellow) forms a complex ion with Ag+, adsorbed
on the silver chloride precipitate, which alters its light-absorbing properties, and hence its
color. The indicator function is critically dependent on the availability of a large
precipitate surface area to allow adsorption. The greatest surface area results from a
precipitate comprised of very small particles (colloidal). Stabilization of these colloidal
particles (recall that a colloid has a very high surface-to-volume ratio) is accomplished by
adding a protective colloid, such as dextrin.
It is important that the titration be conducted quickly in diffuse light as
photodecomposition of the silver chloride renders the solution purple, making it difficult
to discern the pale pink end point signal from the purple background.
PROCEDURE
Preparation of a Standard Silver Nitrate Solution
Obtain from the instructor approximately 8.5 g of AgNO3 in a clean, dry, weighing
bottle. Grind the contents to a fine powder with an agate mortar and pestle. Return the
AgNO3 to the weighing bottle, dry in a microwave oven at medium power for 2 minutes,
and accurately weigh, by difference, the contents into a 500-mL volumetric flask.
5
Dissolve with deionized water and dilute to the mark. Store in a dark area and calculate
the molarity of the solution.
Preparation of the Sample for Analysis
Using the unknown chloride sample from part 1, accurately weigh three samples of the
unknown into three 250-mL Erlenmeyer flasks (the sample size should be approximately
0.3 g). Add approximately 50 mL of deionized water to each flask and swirl to dissolve.
Add 10 mL of a 2% dextrin suspension, 5 drops of dichlorofluorescein indicator, and
titrate each sample to the rose-pink end point.
CALCULATION OF RESULTS
Calculate your results as the % Cl- in your sample. Be certain to report each individual
value, the mean value, the absolute deviation from the mean of each value, the relative
average deviation in parts per thousand, the standard deviation and the confidence
interval at an appropriate confidence level.
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