Quantitative Determination of Chloride Ion in Solution using Two Techniques:
The Mohr Method and Gravimetric Analysis
Purpose:
-
To determine the molarity of a sample of chloride ion in an unknown solution using the Mohr
Method
To isolate the precipitate containing the chloride ion using gravimetric analysis and mass it for
comparison to the Mohr Method mass
Pre-Lab Questions:
1. A chloride unknown weighing 0.4221 g is dissolved in acidic solution and is treated with silver
nitrate. The silver chloride precipitate that forms is filtered, dried, and weighed. The weight of
silver chloride obtained is 0.7632 g. Calculate the percentage of Cl in the unknown chloride
sample.
2. Given the Ksp for each substance to be precipitated in this laboratory, AgCl = 1.8 X 10 -10 and
Ag2CrO4 = 1.2 X 10 -12, which would you predict would precipitate first from solution?
3. Given that the solubility product constant, Ksp, for AgCl is 1.8 X 10 -10, calculate the concentration
of each ion at room temperature in moles per liter and in grams per liter.
4. Given that the solubility product constant, Ksp, for Ag2CrO4 is 1.2 X 10 -12, calculate the
concentration of each ion at room temperature in moles per liter and in grams per liter.
5. Given the answers to 3 and 4, does your answer to #2 change at all? Give a rationale for your
answer here!
Background Information:
The chloride ion (Cl-) is an important anion found in solids and solutions. Chloride is the predominant
anion of sea water with a molarity of approximately 0.54 M. It is an essential component of fluid
systems in the body and is part of the regulatory mechanism for body fluid volume and osmotic
pressure in cells. An average human excretes approximately 16 mg of chloride per day, thus the
detection of the chloride ion becomes germane to water treatment facilities. Bleaching agents can also
contribute to chloride ion concentrations in drinking water. The chemical detection of the chloride ion
is an important process in water purification and treatment. In this experiment, we will determine the
amount of chloride ion in an unknown sample of water using the Mohr method, which relies on the
solubility differences of two anions and the titration endpoint of a precipitate. The net ionic primary
reaction during the titration is as follows:
Ag+ + Cl-  AgCl(s)
The Ksp for AgCl is 1.8 x 10-10 and that for Ag2CrO4 is 1.2 x 10-12. Thus, as silver ions are slowly
added to solution, AgCl precipitates first, removing the chloride ions from solution. As can be
inferred, Ag+ ions react with almost all chloride ions before precipitation of Ag2CrO4 begins. Thus, the
secondary reaction is:
2AgNO3(aq) + K2CrO4(aq)
Ag2CrO4(s) + 2KNO3(aq)
And results in a color change from yellow to reddish-brown at the end point. Knowing the molarity of
the titrant allows for calculation of the chloride concentration.
Gravimetric analysis is a standard classical method for determining the amount of a given component
present in many solutions or solid-unknown samples. The method involves precipitating the
component of interest from the unknown by means of some added reagent. From the mass of the
precipitate, the percentage of the unknown component in the original sample may be calculated.
Generally, the reagent causing precipitation is chosen to be as specific as possible for the component of
interest. The process is intended to remove and weigh only the particular component of interest from
the unknown sample. The precipitating agent must also be chosen carefully so that it completely
precipitates the component under study. The resulting precipitate must have an extremely low
solubility, must be of known composition, and must be chemically stable. Handbooks of chemical data
list suggested precipitating agents for routine gravimetric analysis of many unknowns.
A complete gravimetric analysis includes a series of distinct steps. First, a precise, known amount of
original unknown sample must be taken for the analysis. If the unknown is a solid, an appropriately
sized portion is taken and weighed as precisely as permitted by the balances available (typically to the
nearest 0.001 or 0.0001 g). If the unknown is a solution, an appropriately sized aliquot (a known
fraction) is taken with a pipette. The sample of unknown taken for the analysis must be large enough
that precision may be maintained at a high level during the analysis but not so large that the amount of
precipitate generated cannot be handled easily.
Next, the unknown sample must be brought into solution (if it is not already dissolved). For some solid
samples, such as metal ores, this may involve heating with acid to effect the dissolving. Frequently, the
pH of the solution of unknown must be adjusted before precipitation can take place; this is usually
done with concentrated buffer systems to maintain the pH constant throughout the analysis.
The reagent that causes the precipitation is then added to the sample. The precipitating reagent is added
slowly, and in fairly dilute concentration, to allow large, easily filterable crystals of precipitate to form.
The precipitate formed is allowed to stand for an extended period, perhaps at an elevated temperature,
to allow the crystals of precipitate to grow as large as possible. This waiting period is called digestion
of the precipitate.
The precipitate must then be filtered to remove it from the liquid. Although filtration could be
accomplished with an ordinary gravity funnel and filter paper, this would probably be very slow. The
precipitates produced in gravimetric analysis are often very finely divided and would tend to clog the
pores of the filter paper. Specialized sintered glass filtering funnels have been prepared for routine
gravimetric analyses. Rather than filter paper, such funnels have a frit plate constructed of several
layers of very fine compressed glass fibers that act to hold back the particles of precipitate. Such glass
funnels can use suction to speed up the removal of liquid from a precipitate, can be cleaned easily
before and after use, and are not affected by reagents in the solution.
Once the precipitate has been transferred to the funnel, it must be washed to remove any adhering ions.
As the precipitate forms, certain ions of the other species present in the mixture may have been
adsorbed on the surface of the crystals. The liquid chosen for washing the crystals of precipitate is
usually chosen so as to remove such chemically adhering ions. The washing of the precipitate must be
performed carefully to prevent peptization of the precipitate. The wash liquid must not redissolve the
precipitate or break up the crystals to the point where they might be lost through the pores of the filter.
After the precipitate has been filtered and washed it must be dried. This is usually accomplished in an
oven whose temperature is rigorously controlled at 110o C. The oven must be hot enough to boil off
water adhering to the crystals but cannot be so hot that it might decompose the crystals. Some
precipitating reagents are organic in nature and cannot stand very strong heating.
Finally, the dried precipitate is weighed. From the mass and composition of the precipitate, the mass of
the component of interest in the original sample is determined. Generally, gravimetric analyses are
done in triplicate (or even quadruplicate) as a check on the determination. The precision expected in a
good gravimetric analysis is very high, and if there is any major deviation in the analyses, it is likely
due to some source of error either in the procedure or by the operator.
One of the most common gravimetric analyses is that of the chloride ion, Cl-1. Chloride samples
determined by this method typically include simple inorganic salts, sample of brine, or even body
fluids.
In one type of analysis, a weighed, dried, powdered sample containing chloride ion is first dissolved in
dilute nitric acid. Then a dilute solution of silver nitrate is added, which causes the formation of a
precipitate of silver chloride. The silver chloride is then digested, filtered, dried, and weighed. The net
ionic reaction for this process is simply:
Ag+1 + Cl-1  AgCl(s).
The weight of chloride ion present in the original sample is given by (mass of AgCl precipitated) x
(molar mass of Cl/molar mass of AgCl) and the percent of chloride ion in the original sample is given
by [(mass of Cl)/ (mass of sample)] x 100.
The original sample is dissolved in dilute nitric acid (rather than distilled water) to prevent
precipitation of any other negative ions that might be present in the sample. For example, if carbonate
ion (or bicarbonate ion) were present in a sample, they too would be precipitated by silver ion. Nitric
acid can destroy or mask such interferences.
As initially formed, the particles of AgCl are too small to be filtered easily. Digestion of the AgCl
allows smaller particles of AgCl to combine to form larger crystals that can be filtered more
quantitatively. Digestion of the precipitate may be effected by either of two methods. If it is desired to
complete the analysis during a single laboratory period, then the precipitate and supernatant solution
can be heated almost to boiling for approximately an hour. If the analysis is to be finished at a later
time, the precipitate and supernatant liquid can be allowed to stand in a dark place for a few days.
Silver ion can be reduced by light to metallic silver, with loss of chloride ion to the atmosphere. Loss
of chloride ion by this mechanism would affect the weight of precipitate obtained.
Materials:
One volumetric pipette, one 250 mL Erlenmeyer flask, 50 mL buret, ring stand and buret clamp, 100
mL graduated cylinder, stirrer, 0.100 M AgNO3, 1.0 M K2CrO4, unknown chloride solution
Procedure:
1) Using the volumetric pipette, transfer 20.0 mL of your unknown chloride solution into a 250 mL
Erlenmeyer flask. Add about 50 mL of distilled water to your unknown.
2) Clean a 50 mL buret and rinse with two 5 mL portions of AgNO3 solution. Do not discard the
rinse. Fill the buret with the 0.100 M AgNO3 titrant solution.
3) Add 3 drops of 1.0 M potassium chromate (K2CrO4) solution to the first flask which will give the
solution a yellow color.
4) Record the initial reading of AgNO3 in the buret, and then slowly add AgNO3 to the flask while
swirling or stirring. Reddish specks form upon contact; slow down and continue until the reddishbrown color persists. Record the volume of titrant used.
5) Allow the supernatant and precipitate to sit in a cool, dark, place to allow the crystals to digest.
6) After two days, filter the precipitate using a Gooch-type crucible and a vacuum apparatus. Make
sure that you mass your crucible first. Carefully rinse and wash your precipitate in your crucible
with small amounts of distilled water from a wash bottle. Following your rinse technique, heat
your filtering funnel in a drying oven at 1100C for at least an hour to dry your precipitate.
7) Dispose of the contents in the flask, as well as the remainder of the buret’s silver nitrate, in the
waste bottle labeled silver waste. Clean your buret twice with distilled water.
Data:
Create a data table that arranges all data, collected or derived, neatly and coherently.
Calculations and Conclusions:
1) Calculate the weight of chloride ion present in the original sample. Then, calculate the percent of
chloride ion in the original sample. Show all relevant calculations.
2) Calculate your % error using the measurement from question #2, and from an accepted value.
3) Compare the concentration of chlorine ion in solution obtained during experimentation to that
obtained in pre-lab question #3. Are they the same? Should they be the same? Explain!
4) Discuss the significance of the Ksp of AgCl on the accuracy of the Mohr test.
5) Suggest a method for detecting the presence of another anion of your choice using the technique
employed in this lab, i.e., titration to an end point.
6) Suggest a titration method for detecting the amount of the transition metal ion Cu2+ present in a
solution of CuSO4 using a titration method.
7) The Ksp of AgCl is larger than the Ksp of Ag2CrO4, yet Ag2CrO4 is more soluble than AgCl.
Explain the fallaciousness of the above statement.