Experiment 8
Impurities in Natural Water
A. Purpose
1. To learn how to test for the contaminants of natural water (mostly ions in this
experiment).
2. To learn some techniques for the purification of natural water.
3. To learn the technique of water “softening” using ion-exchange resins.
B. Theory
Water covers nearly three-quarters of the earth and is the most abundant substance
on the planet surface. Less then 3% of the earth’s water is actually “fresh water”.
There is no such thing as “pure” spring water. If it has been in contact with earth,
it holds minerals and, frequently gases in solution. Impurities in natural water
consist of any substances other than the H2O molecule. Today, the production of
ultrapure water that is both sterile and free of dissolved and suspended
impurities, is carried out in research laboratories using different techniques such
as distillation, reverse osmosis (RO), deionization, carbon adsorption and
membrane filtration.
Drinking water must be free of harmful bacteria (sterile), suspended matter, odor,
color and objectionable taste. Dissolved minerals and gases, unless present in
large amounts, usually do not need to be removed. Sometimes they can even be
beneficial. If water is to be used for industrial purposes, appreciable
concentrations of minerals, which are chiefly the salts of calcium or magnesium,
must be removed. The removal of such ions by chemical means is usually referred
to as “softening”.
Testing for the presence of several ions in water involves chemical reactions that
produce visible colored products or precipitates. In some situations, these
reactions can be used for quantitative estimation of the concentration of these ions
in solution. You certainly remember the complexation of M2+ ions with EDTA.
Titration with EDTA provides an example of a quantitative determination of such
ions in water.
C. Materials
Solid CoCl2, 14 M aqueous ammonia (NH4OH), phenolphthalein indicator,
saturated Co(OH)2 solution, 6 M HNO3, 6 M HCl, 0.05 M AgNO3, 0.5 N BaCl2,
0.2 N Al2(SO4)3, 0.0100 M EDTA solution, pH 10 buffer solution, EBT indicator,
ammonium molybdate reagent, SnCl2 in glycerol solution, 5.0 mg/L phosphate
standard, muddy water.
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D. Procedure (with Setup and Theoretical Explanations)

Purification of water by distillation:
Set up the distillation apparatus as shown in Fig. 1.
Fig. 1
1. Add 100 mL of water and a few crystals of CoCl2 to the boiler to give the
water a slight pink color. Put in a few boiling chips, and distill slowly until
about 2 cm of distillate has accumulated in the receiving tube. Note the
color of the distillate as compared with the initial mixture. Is the pinkcolored impurity present in the distillate? Save this distillate (label A1) for
part B, covering the tube with parafilm paper.
2. Empty the still, rinse it well with distilled water then add 5 mL of
concentrated aqueous ammonia, 25 mL of distilled water and distill
slowly, using a fresh receiver. Note the odor of the distillate as compared
with that of the solution being distilled.
Pour a few drops of the distillate into a clean test tube. Add one drop of
phenolphthalein indicator. Note the color. Now add one drop of
phenolphthalein to the original solution. Compare the color with that of
the distillate test.
Save the remainder of the distillate (label A2), for part B.

Minerals in tap water:
Measure about 125 mL of tap water into a beaker and start it boiling gently.
When the volume has been reduced to about half, withdraw a 20 mL sample
(call it B1) and place it into a large test tube. Boiling away some of the water
concentrates ions that may be present in a very small amount of tap water.
Continue boiling of the remaining contents of the beaker to dryness. Remove
the flame as soon as the water is gone (to avoid cracking of the glass), and
save the residue for step 5 below.
3. Calcium and magnesium hardness: Rinse a 50-mL buret with about 10
mL of EDTA solution, mount it vertically in a ring stand, and fill it above
the zero mark with EDTA solution. Open the valve and carefully withdraw
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solution into a waste beaker until the meniscus stands at exactly 0.0 mL.
No bubble should remain at the tip. With a buret measure 100.00 mL of
tap water to be tested and pour it into a 250-mL Erlenmeyer flask. Add 1
mL of the buffer solution and 10 drops of EBT indicator. Titrate with
EDTA solution from the buret with active stirring until the win-red color
of the indicator turns blue (remember to prepare a comparison solution,
see Exp. 2). You should not add more than 1 drop of excess. The volume
of the EDTA solution is read to the nearest 0.01 mL. Calculate the
concentration of CaCO3 (in mg CaCO3/L).
4. Chloride ion test: Divide the 20-mL sample of tap water (B1) into two
parts and to the first one, add a few drops of HNO3 and 1 mL of AgNO3
solution (provided in a dropping bottle). A precipitate or cloudiness is
AgCl, which indicates that Cl ion was present in the sample. Write the
equation. Nitric acid, HNO3, need not be included in the equation because
it merely serves to prevent a false test. Try the test on water directly from
the tap.
Then repeat the test on the distillates A1 and A2. Record your conclusions.
5. Sulfate ion test: To the other half of the sample, add a few drops of HCl
and about 1 mL of BaCl2 solution. What do you observe? The substance
(precipitate or cloudiness) is BaSO4 and indicates that SO 24  ion was
present. Write the equation. The HCl need not be included in the equation.
Try the test directly on tap water.
Then repeat the test on the distillates A1 and A2. Record your conclusions.
6. Phosphate ion test. Transfer 25 mL of a tap water sample to be tested to a
clean Erlenmeyer flask or beaker. Note: This test is particularly interesting
if done on surface water that becomes green with algae in the summer
because of excessive phosphate in the water. Before delivering the water,
the beaker or flask should be cleaned with warm water followed by
thorough rinsing with laboratory (distilled or deionized) water. Cleaning
with soap is not recommended because soaps and detergents often contain
phosphate.
Add 20 drops of 0.025 M ammonium molybdate solution (provided in a
dropping bottle) to the 25-mL water sample, and swirl to mix. Add 2 drops
of stannous chloride (10% SnCl2 in glycerol) solution and mix again by
thorough swirling. If phosphate is present, a blue color will develop to a
maximum intensity in 5 minutes. For the sake of comparison, run the same
phosphate test on a “standard” water sample containing 5.0 mg/L of
phosphate. The intensity of the blue color is directly related to the
concentration of phosphate in the water. How does your water sample
compare in phosphate content to the 5.0 mg/L standard? Record on the
report sheet.
Repeat the same tests on the sea water sample you have collected for
today’s experiment.
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Useful readings: “Phosphate in the Little Miami River”
(http://w3.one.net/~jwclymer/lmr.html); “Phosphate Snap Test Kit”
(http://webpages.charter.net/kwingerden/erhs/aquarium/infophos.htm)
7. Indication of carbonates: Inspect the solid residue from the evaporation
to dryness. It consists for the most part of CaCO3 and MgCO3 produced
by the reactions:
Ca(HCO3)2 (aq)  CaCO3 (s) + H2O (l) + CO2 (g)
Mg(HCO3)2 (aq)  MgCO3 (s) + H2O (l) + CO2 (g)
Put a few drops of HCl solution on the residue and tilt the beaker so as to
let it flow slightly. A slight bubbling at the leading edge indicates the
carbonate ion CO 32  . A typical equation is:
2 HCl (aq) + CaCO3 (s)  H2O (l) + CO2 (g) + CaCl2 (aq)

Water softening using ion-exchange resins:
8. Household type (cation-exchange resin in which cations are replaced
by Na). Pour 50 mL of tap water from a beaker through a column of
single resin (DOWEX 50W-X8) water softener like the model shown in
Fig. 2.
Fig. 2
By ion exchange, the resin captures all metal ions and releases Na+ ions in
their place. Hence the water now contains increased concentrations of
sodium salts such as NaCl. A typical equation for this ion exchange
process is:
Na2R + CaCl2  CaR + 2NaCl
In this equation, the symbol “R” represents the negatively charged site of
the cation exchange resin. Make the hardness test on resin-softened water.
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Also make the chloride, sulfate, and phosphate ion tests and record your
result.
9. Deionizing type (ion exchange using both cation- and anion-exchange
resins). Combinations of ion exchange resins are used in deionizing units.
The cation-exchange resin removes all metal ions and replaces them with
H ions, and the anion-exchange resign removes negative ions and
replaces them with OH ions. The neutralization reaction H + OH 
H2O follows. Since the total positive charge must equal the total negative
charge in solution, all cations and anions are effectively replaced by water
molecules. Preparation of deionized water is less expensive than
distillation of water due to lower energy requirements. Deionized water is
often substituted for distilled water in scientific applications. Obtain a
sample of deionized water and test for hardness, chloride ions, sulfate
ions, and phosphate ions. Record your results.
Useful reading: “Theory of Ion Exchange”
http://www.esb.ucp.pt/~bungah/ionex/theoryd.htm
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E. CHEM 203 Lab Report
NAME : _______________________
EXPERIMENT No. : ______________
TITLE : __________________________________
DATE : ________________
Partner : ________________________

Distillation
1. Distillation of a solution containing a dissolved solid (CoCl2)
Is CoCl2 impurity present in the distillate? _______________________
Explain from your observations.
2. Distillation of a solution containing a dissolved gas (NH3)
Is NH3 impurity present in the distillate? _______________________
Explain from your observations.

Minerals in tap water (or other potable water supply)
3. Calcium and magnesium hardness
Initial buret reading _________
Final buret reading __________
Volume of EDTA used __________
Hardness of tap water _________________ (include correct concentration units)
4. Chloride ion test
Evaporated water ________________
Tap water ________________
Equation for test: _______________________________________________
5. Sulfate ion test
Evaporated water ________________
Tap water ________________
Equation for test: _______________________________________________
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6. Phosphate ion test
Does tap water produce a blue color? _______________
How does it compare to the 1.0 mg/L standard? ___________________________
7. Indication of carbonate
Is the CO32 ion present in the residue? ____________
How do you know?
Write a balanced equation for the reaction of HCl with any MgCO3 that may have
been present:

Water softening using ion-exchange resins
8. Household type
Initial buret reading ___________
Final buret reading ____________
Volume of EDTA used ____________
Hardness of softened water _____________
Is chloride ion present? ____________ Is sulfate ion present? _____________
Is phosphate ion present?
9. Deionizing type
Questions
1. What reaction between soil, water and limestone rock accounts for the relatively
large amounts of Ca(HCO3)2 and Mg(HCO3)2 found in ground water. Where does
the necessary CO2 come from?
2. “Boiler scale” forms in kettles, consisting mainly of calcium carbonate. Write a
chemical equation for the corresponding reaction. Explain why the scale forms
much more heavily at high temperature (like the boiling point of water).
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