50
#6. Determination of Hardness of Water
EXPERIMENT 6.
DETERMINATION OF HARDNESS OF WATER WITH EDTA
BACKGROUND
In this experiment the amount of calcium and magnesium present in a water sample
is determined. The determination of the hardness of water is important in water
treatment plants. In this experiment a sample of hard water is titrated with EDTA,
with calmagite as indicator.
In the titration of calcium and magnesium with EDTA the indicator calmagite
works in the following way. When a small amount is added to a solution
containing magnesium at pH 10, red MgIn– forms. As EDTA titrant is added, free
magnesium ions in solution react with it. After all the free magnesium is
complexed, the next portion of EDTA removes magnesium from the indicator
complex, converting the indicator to the blue HIn2– form. The pH must be near 10
because at pH values above 12 or below 8 the free indicator is red like the
magnesium complex, and therefore no color change can be observed.
The calcium-calmagite complex is too weak to function as an indicator (log K
(CaIn) = 3.7 at pH 10). Calcium can, however, be titrated and an end point
obtained when a trace of magnesium is present. In this experiment a small amount
of magnesium is added to the titrant since the primary standard, pure calcium
carbonate, contains none. The optimum pH range for titrations of calcium and
magnesium with EDTA is about 9 to 11. In solutions more acidic than pH 9,
formation of the metal-EDTA complex is incomplete, whereas in solutions above
pH 12, magnesium hydroxide begins to precipitate.
A solution of approximate strength of the disodium salt of EDTA, Na2H2Y.2H2O,
is prepared and standardized against primary standard calcium carbonate. In
addition to having all the properties of a suitable primary standard, calcium
carbonate has the added advantage that the element titrated, calcium, is a major
component of water hardness and of limestone.
#6. Determination of Hardness of Water
51
PROCEDURE
Reagent List:
Unknown sample - do not dry
EDTA , di-sodium salt of - prepare 0.015 M stock soln.
magnesium chloride 1%soln. - supplied
ammonium hydroxide conc. 15M
calcium carbonate - dried
HCl conc.12M
ascorbic acid
Calmagite 0.1 % soln. - indicator
Preparation of EDTA Solution (Must be prepared in advance)
Prepare an approximately 0.015 M EDTA solution by adding 2.5 g of the disodium
salt,1 10 mL of 1% MgCl2 solution, and 1.5 mL of 6 M NH3 to about 450 mL of
water. In this experiment use only distilled water purified by passage through a
column containing either a cation-exchange resin or a mixture of cation- and anionexchange resins. (This water will be provided.)
Preparation of Calcium Carbonate Standard2
Weigh to the nearest 0.1 mg approximately 0.5 g of dry primary-standard CaCO3
into a clean, dry 50 mL beaker. Transfer most of the CaCO3 without loss to a dry
funnel inserted in the top of a dry 100 mL volumetric flask. Transfer of the last
particles is unnecessary. Do not use liquid in this transfer. Weigh the beaker and
any remaining calcium carbonate to the nearest 0.1 mg, obtaining the weight of
CaCO3 in the flask by difference.
Tap the funnel to transfer most of the solid, and then rinse the last particles through
the funnel into the flask with a milliliter of 12 M HCl.3 Add 1 mL more of 12 M
HCl to the flask by rapid dropwise addition. Warm if necessary to complete the
dissolution.4 When dissolution is complete, remove the funnel, rinsing inside and
1
The disodium salt of EDTA is the most commonly used form, being easier to prepare commercially and more
soluble in water than the acid, H4Y. Complete dissolution is somewhat slow, however, requiring 10 to 15 min.
Occasional shaking speeds the process.
2
It may be convenient to prepare the limestone sample solution (see next part of experiment for details) at the same
time as the CaCO3 standard, since dissolution of the limestone is often slow.
3
The 12 M (concentrated) HCl is stored in the fume hood to minimize exposure to its aggressive vapor. Use of
more HCl than the minimum necessary to dissolve the CaCO3 leads to difficulty later with pH adjustment.
4
Heating of volumetric flasks to moderate temperatures on a hot plate does not affect calibration, as the original
volume is regained upon cooling room temperature.
52
#6. Determination of Hardness of Water
out with demineralized water from a wash bottle as it is being removed. Fill the
flask to the mark with distilled water, stopper, and mix well.
Sample Preparation
Collect an unknown sample from your instructor. Record the number of your
unknown and the total mass of the sample in your laboratory notebook. Transfer
the unknown sample through a dry funnel into a dry 100 mL volumetric flask. Tap
the funnel to ensure that most of the sample has been transferred into the flask.
Rinse the sample container with approximately 1 mL of concentrated hydrochloric
acid. Transfer this solution into the volumetric flask ensuring that all remaining
particles of your unknown are washed into the volumetric flask. Add an additional
1 mL of concentrated HCl(aq) to your flask by rapid drop-wise addition. Gently
warm the flask to ensure complete dissolution. Wash the walls of the funnel with
distilled water, fill the flask to mark, stopper and mix well.
Titration of Standards
Using a calibrated 10 mL pipet, measure aliquots of the standard calcium solution
into each of three or four 200 or 250 mL conical flasks.8
Titrate the calcium standards and samples in alternation. Immediately before each
titration, take a flask containing the aliquot and add about 50 mg of ascorbic acid to
the flask, then approximately 10 mL of 6 M NH3,9. Swirl the flask to mix the
contents after each addition. Add 4 to 5 drops of 0.1% calmagite indicator
solution, and titrate with EDTA solution until the indicator changes from red to
blue.10 (As before, use the drop of EDTA that gives the greatest color change as
the end point.)
Titration of Unknown
Follow the same titration procedure for the solution of your unknown as the
procedure described above for the titration of the standard samples.
8
If the sample contains some insoluble silicate material, avoid inclusion of the solids in the pipetted portion insofar
as possible.
9
Ammonia alone is added here because the presence of excess HCl from the dissolution must be neutralized. This
normally provides sufficient NH4Cl to serve as the HB+ component of the buffer.
10
Some limestones contain heavy metals that slowly bind strongly to calmagite. Therefore solutions should be
titrated promptly and not too slowly once indicator has been added, otherwise the color change may be poor.
#6. Determination of Hardness of Water
53
CALCULATIONS FOR WATER HARDNESS
The concentration of EDTA solution is found by
(wt. of CaCO3)(V10/100 mL)
M EDTA = ((mol wt CaCO ))(mL EDTA/1000)
3
where V10 is the calibrated volume of the 10 mL pipet.
Water hardness, expressed in parts per million or milligrams per liter of CaCO3,11
can be calculated by
hardness in ppm =
(mL EDTA/1000)(M EDTA)(mol wt CaCO3)(1000 mg/g)
(vol of sample in liters)
The hardness in ppm value represents the CaCO3 concentration in the solution
prepared from your unknown sample.
Report the result as ppm of CaCO3 in the solid unknown sample that was
supplied to you.
11
The results could be reported in several ways, such as % CaCO3, % MgCO3, or % Mg. In accordance with
common practice the approach employed here assumes that all the material titrated is calcium, and is calculated as
calcium carbonate.
54
#7. Iodometric Determination of Copper
EXPERIMENT 7.
IODOMETRIC DETERMINATION OF COPPER IN BRASS
BACKGROUND
This experiment illustrates the analytical method involving the iodine-iodide
couple:
I2 + 2e- <===========> 2 I-
E° = 0.53 V
(1)
This couple is important because it has a standard electrode potential that permits
the analytical use of iodine as an oxidant for substances of lower electrode
potential and iodide as a reductant for substances of higher potential. Because its
electrode potential is little affected by either pH change or complexing agents, this
couple can be used in conjunction with half-reactions that change potential with pH
or with the addition of auxiliary reagents.
Analytically useful applications include those in which solutions of iodine are used
to titrate reduced materials directly and those in which oxidizing agents are
determined through oxidation of iodide to iodine. In the latter the iodine formed is
titrated with a standard solution of sodium thiosulfate, Na2SO3.
Iodometric Determination of Copper
In this experiment the copper in a brass sample is determined by a method
involving the liberation of iodine. The sample is dissolved in nitric acid, and the
solution boiled to remove most of the nitrogen oxides formed during the metal
oxidation. The residual nitrogen oxides are eliminated by the addition of urea.
Complete removal of nitrogen oxides is necessary to prevent iodide oxidation.
Iron present in most brasses also causes iodide oxidation. This interference is
eliminated by the addition of fluoride, which forms a stable complex with iron(III).
Next the pH is adjusted to 3.5 to 4.5, an excess of potassium iodide is added, and
the iodine formed is titrated with sodium thiosulfate.
Although the titration of iodine is the only titration of significance for which
thiosulfate is used as a standard solution, it is an important one. Other applications
of thiosulfate are few. One reason is that oxidizing agents stronger than iodine
oxidize it to a mixture of higher oxidation states of sulfur; another is that ions of
#7. Iodometric Determination of Copper
55
transition metals such as copper decrease the stability of sodium thiosulfate
solutions by catalytic air oxidation. Acids also make thiosulfate solutions unstable
by promoting disproportionation to sulfite and elemental sulfur:
–
HS2O3
–
-----------------> HSO3
+ S
(2)
Addition of a small amount of sodium carbonate to make the thiosulfate solution
alkaline prevents this decomposition. Water used for the preparation of standard
thiosulfate solutions may be boiled to destroy sulfur bacteria, which find these
solutions an attractive medium for growth. Alternatively, a bacterial agent such as
a mercury(II) salt or chloroform may be added.
Any of several primary standards, including iodine, potassium chromate, potassium
iodate, and electrolytic copper metal may be used to standardize thiosulfate
solutions. For this experiment copper is chosen because it is the material being
determined and is readily available in primary standard quality as electrical wire.
The principal reactions involved in iodometric copper analysis are
(a) dissolution of sample in dilute nitric acid,
3 Cu + 8 HNO3 -----------> 3 Cu2+ + 2NO + 4 H2O + 6 NO3-
(3)
Cu + 4 HNO3 ------------> Cu2+ + 2 NO2 + 2 H2O + 2 NO3-
(4)
or
(b) removal of residual nitrogen oxides by addition of urea,
2 NO + O2 <===========> 2 NO2
(5)
2 NO2 + H2O <============> HNO2 + HNO3
(6)
and
2 HNO2 + (NH2)2CO ----------------> 2 N2 + CO2 + 3 H2O (7)
(c) neutralization of remaining acid with sodium hydroxide, followed by pH
adjustment;
56
#7. Iodometric Determination of Copper
(d) complexation of iron(III) with fluoride,
Fe3+ + F- <==========> [FeF]2+
(e) addition of excess potassium iodide,
2 Cu2+ + 4 I- <===============> Cu2I2 + I2
(8)
(9)
and
(f) titration of iodine with thiosulfate,
I2 + 2 S2O32- <===========> 2 I- + S4O62-
(10)
Precautions must be taken to avoid side reactions. For instance, iodine will slowly
oxidize tetrathionate to sulfate, especially at high pH values. At low pH sulfurous
acid may be formed by thiosulfate (equation (2)). Another source of error in strong
acid solution is air oxidation of iodide:
O2 + 4 H+ + 4 I- <============> 2 I2 + 2 H2O (11)
This is called oxygen error. These side reactions may be avoided by carrying out
the thiosulfate - iodine reaction in the pH range 2 - 5.
The volatility of iodine creates problems. Iodine loss can be minimized by keeping
the temperature low, titrating promptly, and adding excess iodine to stabilize the
iodine in solution as triiodide:
I2 + I- <==========> I3-
Keq = 7.1 x 102
(12)
Starch gives an intense blue color with triiodide that serves as a specific indicator.
The triiodide ion appears to be just the right size to enter the helical structure of
starch and thereby form a colored complex. The color of free iodine, red-brown in
high and yellow in low concentrations, also can be used as an indication of its
presence, but it is not so easily seen as that of the complex with starch. In acidic
solutions starch tends to undergo decomposition that is accelerated by high
concentrations of iodine. Therefore, addition of starch is best delayed until near
the equivalence point of the titration.
#7. Iodometric Determination of Copper
57
Most strong oxidizing agents can be determined iodometrically. An excess of
iodide is added, and iodine is produced in an amount equivalent to that of the
oxidant present in the sample. The liberated iodine is titrated with thiosulfate.
PROCEDURE
(Median Time 3.9 hours)
Reagent List:
Unknown Sample
Sodium thiosulfate•5H2O - prepare 0.1M soln.
Na2CO3
HNO3 conc. 16M
copper turnings >99% pure
Urea - 4% soln. needed
sodium hydroxide pellets - 2.5M soln. needed
ammonium acid fluoride(NH4HF2)
potassium iodide
Starch - .2% soln indicator
Standardization of 0.1 M Na2S2O3 Solution
Dissolve 25 g of Na2S2O3.5H2O and 0.1g of Na2CO3 in a liter of distilled water.
N.B. This solution may be saved for Experiment 8.
Weigh accurately 0.2 g samples of clean copper turnings into 250 mL conical
flasks.1 Add 10 mL of 6 M HNO3 to each in a fume hood. When dissolution is
complete, boil the solution for a short time to remove most of the nitrogen oxides.
Add 10 mL of water and 5 mL of 4 % urea solution, and boil again for about one
minute. Allow the solution to cool. When ready to titrate, add 30 mL of water and
then add 2.5 M NaOH until a slight permanent precipitate of Cu(OH)2 is obtained.
This will require about 15 to 20 mL of NaOH, depending on the amount of HNO3
present. Add 1 to 2 g of ammonium acid fluoride, NH4HF2, and swirl until
dissolved.
Cool, add approx. 3 g of KI, and titrate immediately to near the end point with
Na2S2O3. When the solution has become pale yellow or buff, add about 5 mL of
58
#7. Iodometric Determination of Copper
fresh starch solution and titrate to the first complete disappearance of blue.2
PROCEDURE FOR THE SAMPLE
Do not dry the unknown. Determine the total weight of the unknown. Transfer the
unknown to a 250 mL volumetric flask. Dissolve the unknown with 10 mL of 6 M
HNO3. Dilute to 250 mL. Pipet 25 mL into a 250 mL erlenmeyer. Add urea and
boil as described for the standards. When ready to titrate, neutralize with sodium
hydroxide and dilute to 50 mL with water. Add about 1 g of ammonium acid
fluoride and swirl until dissolved. Add approx. 3g of KI and complete the
determination as in the standardization.
Calculate the molarity of the Na2S2O3 solution. One mole of copper requires 1
mole of thiosulfate for titration. Calculate and report the percentage of copper in
the sample.
1.
2.
NOTES
Trace impurities markedly increase the resistance of copper wire. For
electrical use they are removed by electrolytic refining to a level well below a
part per thousand. Copper produced for electrical wiring is therefore an
excellent primary standard. The thin coat of oxide sometimes present on the
wire can be removed by polishing with fine emery cloth, followed by wiping
with clean toweling.
To avoid etching of flasks by HF, empty and rinse immediately after
completing each titration.
#8 . Determination of Ascorbic Acid
59
EXPERIMENT 8.
THE DETERMINATION OF ASCORBIC ACID BY
TITRATION WITH POTASSIUM BROMATE
DISCUSSION
Ascorbic acid, C6H8O6, is cleanly oxidized to dehydroascorbic acid by bromine:
H
H
O O
H
H
O O
C C
O O
C C
O C
C C C
O
H H
H +
H
Br2
O
C
H H
O
C C
O
H
O
C H + 2 Br - + 2H +
H H
An unmeasured excess of potassium bromide is added to an acidified solution of
the sample. The solution is titrated with standard potassium bromate to the first
permanent appearance of excess bromine; this excess is then determined
iodometrically with standard sodium thiosulfate. The entire titration must be
performed without delay to prevent air-oxidation of the ascorbic acid.
Reagent List:
Unknown sample
Potassium Bromate - dried 110oC
Sodium thiosulfate•5H2O - 0.1 M soln. from Exp.#7
Na carbonate
potassium iodide
H2SO4 conc. 18M
starch 0.2% soln. indicator
potassium bromide
uncalibrated buret
60
#8 . Determination of Ascorbic Acid
PREPARATION OF A STANDARD 0.015 M
POTASSIUM BROMATE SOLUTION
Transfer about 1.5 g of reagent-grade potassium bromate to a weighing bottle, and
dry at 110°C for at least 1 hour. Cool in a desiccator. Weigh approximately 1.3 g
(to the nearest 0.1 mg) into a 500 mL volumetric flask; use a powder funnel to
ensure quantitative transfer of the solid. Rinse the funnel well, and dissolve the
KBrO3 in about 200 mL of distilled water. Dilute to the mark, and mix
thoroughly.
Solid potassium bromate can cause a fire if it comes into contact with damp
organic material (such as paper toweling in a waste container). Consult with
the instructor concerning the disposal of any excess.
PREPARATION OF 0.1 M SODIUM THIOSULFATE
(NOT REQUIRED IF YOU SAVED SOLUTION FROM EXPERIMENT #7.)
Boil about 1 L of distilled water for 10 to 15 min. Allow the water to cool to room
temperature; then add about 25 g of Na2S2O3.5H2O and 0.1 g of Na2CO3. Stir
until the solid has dissolved. Transfer the solution to a clean glass or plastic bottle,
and store in a dark place.
STANDARDIZATION
POTASSIUM
OF
BROMATE
SODIUM
(NOT
THIOSULFATE
REQUIRED
IF
AGAINST
YOU
SAVED
SOLUTION FROM EXPERIMENT #7)
Iodine is generated by the reaction between a known volume of standard potassium
bromate and an unmeasured excess of potassium iodide:
BrO3 + 6I + 6H+ Æ  Br + 3I2 + 3H2O
#8 . Determination of Ascorbic Acid
61
The iodine produced is titrated with the sodium thiosulfate solution.
Pipet 25.00 mL aliquots of the KBrO3 solution into 250 mL conical flasks and
rinse the interior wall with distilled water. Treat each sample individually beyond
this point. Introduce 2 g of KI and about 5 mL of 3 M H2SO4. Immediately titrate
with 0.10 M Na2S2O3 until the solution is pale yellow. Add 5 mL of starch
indicator, and titrate to the disappearance of the blue color.
Calculate the concentration of the thiosulfate solution.
EXPERIMENTAL
Collect an unknown sample from your instructor. Immediately copy the sample
number into your laboratory note book and record the total mass of the sample.
Ensure that your sample is homogeneous and then weigh individual 0.35 to 0.40 g
samples (to the nearest 0.1 mg) into dry 250 mL conical flasks. Treat each sample
individually beyond this point. Dissolve the sample in 50 mL of 1.5 M H2SO4;
then add about 5 g of KBr. Titrate immediately with standard KBrO3 to the first
faint yellow due to excess Br2. Record the volume of KBrO3 used. Add 2 g of KI
and 5 mL of starch indicator; back-titrate (Note 1) with standard 0.10 M Na2S2O3.
Calculate the average weight (in milligrams) of ascorbic acid (FW = 176.13) in
your unknown sample. Report the wt. % of ascorbic acid in your original sample.
1.
NOTES
The volume of thiosulfate needed for the back-titration seldom exceeds a few
milliliters.