This document describes the precipitation method for determining the chloride ion concentration of a solution by titration with silver nitrate. Silver nitrate is added until all chloride ions are precipitated as silver chloride. Additional silver ions then react with potassium chromate indicator to form a red-brown silver chromate precipitate, signaling the endpoint. The method can be used to analyze water samples. It involves titrating aliquots of the sample with a standardized silver nitrate solution until concordant results are obtained.

1. Precipitation method
Adit Muktadir

2. Introduction
 This method determines the chloride ion concentration of
a solution by titration with silver nitrate. As the silver
nitrate solution is slowly added, a precipitate of silver
chloride forms.
 Ag+(aq) + Cl–(aq) → AgCl(s)
 The end point of the titration occurs when all the chloride
ions are precipitated. Then additional silver ions react
with the chromate ions of the indicator, potassium
chromate, to form a red-brown precipitate of silver
chromate.
 2 Ag+(aq) + CrO42–(aq) → Ag2CrO4(s)

3. This method can be used to
determine the chloride ion
concentration of water samples from
many sources such as, stream water,
river water and many drugs and
chemicals.

4.  Solutions Needed
 Silver nitrate solution: (0.1 mol L−1) If possible, dry
5 g of AgNO3 for 2 hours at 100°C and allow to
cool. Accurately weigh about 4.25 g of solid AgNO3
and dissolve it in 250 mL of distilled water in a
conical flask. Store the solution in a brown
bottle.Potassium chromate indicator solution:
(approximately 0.25 molL-1) Dissolve 1 g of K2CrO4
in 20 mL distilled water.

5.  Equipment Needed
 burette and stand
 10 and 20 mL pipettes
 100 mL volumetric flask
 250 mL conical flasks
 10 mL and 100 mL measuring cylinders

7. Method
Sample Preparation
 Titration
 1. Dilute sample by pipetting a 20 mL sample into a 100
mL volumetric flask and making it up to the mark with
distilled water.
 2. Pipette a 10 mL aliquot of diluted sample into a conical
flask and add about 50 mL distilled water and 1 mL of
chromate indicator.
 3. Titrate the sample with 0.1 mol L−1 silver nitrate
solution. Although the silver chloride that forms is a white
precipitate, the chromate indicator initially gives the cloudy
solution a faint lemon-yellow colour.

8.  Repeat the titration with further aliquots of
diluted sample until concordant results (titres
agreeing within 0.1 mL) are obtained.

9. Figure 1 Before the addition of any silver nitrate the chromate indicator gives the
clear solution a lemon-yellow colour.

10.  The endpoint of the titration is identified as
the first appearance of a red-brown colour of
silver chromate

11.  Figure 2 Left flask: before the titration endpoint, addition of Ag+
ions leads to formation of silver chloride precipitate, making the
solution cloudy. The chromate indicator gives a faint lemon-yellow
colour. Centre flask: at the endpoint, all the Cl− ions have
precipitated. The slightest excess of Ag+ precipitates with the
chromate indicator giving a slight red-brown colouration. Right
flask: If addition of Ag+ is continued past the endpoint, further
silver chromate precipitate is formed and a stronger red-brown
colour results. NB: The titration should be stopped when the first
trace of red-brown colour is observed. Using an incompletely
titrated reference flask for comparison is a helpful way to identify
the first appearance of red-brown colouration.

12.  Result Calculations
 1. Determine the average volume of silver
nitrate used from your concordant titres.
 2. Calculate the moles of silver nitrate reacting.
 3. Use the following reaction equation to
determine the moles of chloride ions reacting.
Ag+(aq) + Cl–(aq) → AgCl(s)
 4. Calculate the concentration of chloride ions
in the diluted sample.
 5. Calculate the concentration of chloride ions
in the original undiluted sample.
 6. Calculate the concentration of sodium
chloride in the sample in molL−1, gL−1 and
g/100 mL (%)

13. Additional Notes
 Residues containing silver ions are usually
saved for later recovery of silver metal.
 The Mohr titration should be carried out
under conditions of pH 6.5 – 9. At higher pH
silver ions may be removed by precipitation
with hydroxide ions, and at low pH chromate
ions may be removed by an acid-base
reaction to form hydrogen chromate ions or
dichromate ions, affecting the accuracy of the
end point.

14.  It is a good idea to first carry out a “rough” titration in
order to become familiar with the colour change at the
end point.
 The Mohr titration is sensitive to the presence of both
chloride and bromide ions in solution and will not be too
accurate when there is a significant concentration of
bromide present as well as the chloride. However, in
most cases, such as sample, the bromide
concentration will be negligible. For this reason, the
method can also be used to determine either the total
concentration of chloride and bromide in solution, or the
concentration of bromide when the chloride
concentration is known to be negligible.

15. FAJANS METHOD
 Fajans method of chloride determination employs an
adsorption indicator. The indicator reaction takes
place at the surface of the precipitate. The indicator
is a weakly acidic dye and exists in solution in the
ionized form, In-. The titrant is a silver solution, and
during the titration a precipitate of AgCl is formed.
Initially this precipitate is colloidal, consisting of very
small non-settling particles with a diameter of less
than 1 μm. While this would be undesirable for a
gravimetric determination (colloids cannot be
filtered), it is advantageous for an adsorption
indicator method. What happens is the following:

16. Theory
 Precipitates have a tendency to adsorb “their own” ions to
the surface to form what is known as the primary adsorption
layer, i.e., AgCl preferentially adsorbs Ag+ or Cl-, whichever
happens to be in excess. A colloidal precipitate has a very
large surface area and, therefore, presents an abundance of
room for adsorption. Before the equivalence point of the
titration of Cl- with Ag+, the Cl- ion is in excess and forms the
primary adsorption layer on the surface of the AgCl
precipitate. The particles have a negative surface charge and
repel each other; the colloid is stabilized by this. The
indicator ion, In , is also r ‑ epelled and stays well away from
the surface. Because the particles are negatively charged,
they attract cations that are in solution more strongly than
anions. Thus there is weakly bound secondary adsorption
layer consisting of the cation that forms the most insoluble
chloride to AgCl (probably Na+); these ions form the
secondary adsorption layer.

17. IN-IN-IN-IN-IN-IN-IN-IN-
Ag+
Ag+
Ag+
Ag+
Ag+
Ag+
Ag+
Ag+
Ag+ Ag+
Ag+
IN-IN-IN-

18. Beyond the equivalence point, Ag+ is in
excess and the surface of the precipitate
becomes positively charged, with the
primary layer being Ag+. These positively
charged colloidal particles will now attract
the indicator anion and adsorb it into
secondary adsorption layer.

19. The indicator forms a colored complex
with silver ion, imparting a color to the
precipitate. Only at the surface is the
silver ion concentration high enough for
the solubility product of the complex to
be exceeded; this does not happen
anywhere else in the solution, and the
color is therefore confined to the
precipitate surface.

20.  The pH must be controlled for reliable results.
If it is too low, the indicator (a weak acid) is
dissociated too little to produce enough In-.
In the case of Fajans method,
dichlorofluorescein is the preferred indicator
and it gives good results at pH values around
7.

21.  Since the end point does not exactly coincide with the
equivalence point, the titrant should the standardized by the
same titration as used for the sample (this eliminates the
inherent error). Photodecomposition of AgCl, creating a purple-black
hue on top of the white AgCl, is another source of error. It
should be minimized by carrying out the titration expeditiously
and in relative low light. As explained above, a colloidal
precipitate is preferred in this titration. At the equivalence point,
neither titrant nor titrate ions are in excess, and the precipitate is
momentarily without a surface charge. This causes the colloidal
particles to coagulate, thereby reducing the precipitate surface
area. It can be prevented by the addition of a small amount of
dextrin (hydrolyzed starch) to the solution.

22.  The titration reaction is:
AgNO3 (aq) + NaCl (aq) ® AgCl (s) +
NaNO3 (aq)
 In fact, because neither the Na+ nor NO3-
ions are involved in the reaction, this reaction
may be written more accurately as:
Ag+ + Cl-  AgCl(s)

23. The nitrate ions are only weakly adsorbed
to the precipitate after the equivalence
point is reached, and they are easily
displaced by indicator ion. The end point is
signaled by the appearance of the pink
color of silver dichlorofluoresceinate.

24. Procedure
Dry about 2 g of NaCl at 120°C for 1 hr.
Cool in a desiccator. Prepare an ~0.10 M
NaCl standard solution by dissolving the
appropriate amount of salt in a 250-mL
volumetric flask. Pipet 25.0 mL of solution
into a 250-mL Erlenmeyer flask, add 5
drops of indicator, 0.05 g of dextrin, and
titrate the solution in diffuse light with
constant swirling. The silver chloride
flocculates shortly before the equivalence
point is reached.

25. Carry out a rough titration quickly first of
all to find approximately how much AgNO3
is needed to reach the end point. Don’t
worry if you overshoot the end point.
Then do three more replicate titrations
accurately. For each the three samples,
samples add all but one milliliter of the
titrant quickly while swirling the flask.

26. Continue the titration dropwise,
again with vigorous agitation,
until the precipitate suddenly
turns pink. Use the average of
three results to calculate the
silver nitrate concentration.

27.  Dry about 1.5 g of unknown as above.
Accurately (i.e., to the nearest 0.1 mg) weigh a
~0.25-g sample by difference into a 250-mL
Erlenmeyer flasks, dissolve in about 25 mL of
water and titrate as above with the standardized
AgNO3 solution. Accurately weigh three more
samples into Erlenmeyer flasks and dissolve
each sample in ~25 mL of water. (The flasks do
not have to be dry, just clean.) Based on your
first titration, calculate how much each sample
should require to reach the end point. Titrate
each sample using the same procedure as
above, adding all but 1 mL of the titrant rapidly
and then approaching the end-point dropwise.

28.  For each titration, calculate the percent
chloride in the unknown (to two decimal
places). Report the mean of the three
determinations as your final answer. Also
report the relative standard deviation (as a
percentage) of the three determinations.