分析化學實驗
實驗講義
105 學年度下學期
東海大學 化學系
1
Contents
Exp.1 Gravimetric Determination of Nickel in Steel
3
Exp.2 Neutralization Titrations in Aqueous Medium
7
Exp.3 Potentiometric Titration Using a pH Electrode
17
Exp.4 Complex Titrations with EDTA
27
Exp.5 Precipitation Titrations with Silver Nitrate
37
Exp.6 Analysis of Copper in Ore by Iodometric Method
49
Exp.7 Determination of Ascorbic Acid in Vitamin C tablets by Titrations with
Potassium Bromate
59
Exp.8 Determination of Iron in an Ore by Potentiometric Titration
67
Exp.9 Spectrophotometric Determination of Manganese in Steel
79
Exp.10 Simultaneous Determination of Binary Mixture by Spectrophotometric
Method
85
2
Experiment 1
Gravimetric Determination of Nickle in Steel
Introduction
The nickel in a steel sample can be precipitated from a slightly alkaline
medium with an alcoholic solution of dimethylglyoxime (DMG). Tartaric acid is
introduced to complex iron(III) and prevent its interference. After drying at 110 oC,
the organic nickel compound, Ni(C4H7O2N2)2, serves as a convenient weighing form.
Owing to the bulky character of the precipitate, only a small quantity of nickel can
be conveniently handled. The sample weight taken is governed by this consideration.
The excess of precipitating agent must be controlled not only because its solubility
in water is low but also because the nickel compound becomes appreciably more
soluble as the alcohol content of the precipitating medium is increased.
Reagents
1. 6 M hydrochloric acid
2. Concentrated nitric acid
3. Concentrated ammonia
4. 15% (w/v) tartaric acid (15 mL/sample)
5. 1% (w/v) dimethylglyoxime in ethanol (10 mL/sample)
6. 3M ammonium acetate
3
Procedure
1. Dry a piece of glass fiber filter paper to constant weight for at least one hour at
110 oC.
2. Receive one of the unknown Ni sample, and record the unknown number into
the data sheet. Weigh 0.3~0.4 g of the Ni sample into a 250-mL Erlenmeyer
flask, and dissolve it by warming it with about 15 mL of 6 M HCl.
3. Introduce carefully about 1.0 mL of 16 M HNO3, and boil gently for 2 minutes
to expel the oxides of nitrogen. (Attach a gas collecting apparatus to the top of
the Erlenmeyer flask to collect the emitted NOx/HCl vapor.) Transfer the
solution to a 250-mL beaker. Wash the residue in the Erlenmeyer flask with
distilled H2O into the beaker and dilute the resulting solution to about 100 mL.
4. Introduce about 15 mL of 15% tartaric acid (Note 1), and add sufficient
ammonia slowly with good stirring (Note 2) until a slight excess is present (pH
≒9). If the solution is not clear at this stage, proceed as directed in Note 3.
5. Warm up the solution to 60 ~ 80 oC, and then add 10 mL of 1% (w/v) alcoholic
solution of dimethylglyoxime slowly with good stirring. Introduce 10 mL of
3M ammonium acetate. Digest for 10 minutes at about 60 oC, and then standby
to cool it down to room temperature. Weigh the glass fiber paper.
6. Decant the supernatant to the funnel with glass fiber filter paper under vacuum,
then, pour all residues into the funnel. Wash with water until free of chloride
ions (Note 4). Finally, bring the filter paper with the precipitate
(Ni(C4H7O2N2)2) to constant weight by drying at 110 oC.
4
7. Report the % (w/w) of Ni in the sample.
Notes
1. If the tartaric acid solution is not clear, it should be filtered prior to use.
2. Use a separate stirring rod for each sample and leave it in the solution
throughout the determination
3. If a precipitate is formed upon the addition of base, the solution should be
acidified, treated with additional tartaric acid, and again made alkaline.
Alternatively, the precipitate can be removed by filtration. Thorough washing
of the entire filter paper with a hot, dilute NH3/NH4Cl solution is required; the
washings should be combined with the remainder of the sample.
4. The filtrate can be tested for chloride ions by collecting a ~1.0 mL in a test
tube, acidifying with 1 drop of nitric acid, and adding a drop of 0.1 M AgNO3.
Washing is complete when little or no turbidity is observed in the test solution.
References
1. Skoog, West, Holler and Crouch, Fundamentals of Analytical Chemistry, 8th
ed., Chap. 12
2. D.C. Harris, Quantitative Chemical Analysis, 7th ed., pp.628-637
5
Data
Unknown No.
Weight of sample
g
Weight of glass fiber filter paper
g
Weight of glass fiber filter paper + Ni(DMG)2
g
Weight of Ni(DMG)2
g
the % (w/w) of Ni in the sample
%
Accurate value of %Ni
%
Relative error
%
Calculations
Atomic mass of Ni
Molar mass of Ni(C4H7O2N2)2
100%
Weight of sample
Weight of Ni(C4H7O2N2)2 
% (w/w) of Ni =
Discussion
6
Experiment 2
Neutralization Titrations in Aqueous Medium
Introduction
Neutralization titration is performed with standard solutions of strong acids or
strong bases. In principle, a single solution (acid or base) should be sufficient for a
titration; in practice, however, it is convenient to have both standard acid and
standard base available to locate end points more exactly by the needed back
titration. The concentration of one solution is standardized by a primary standard;
the molarity of the other is then found by titration with the previously standardized
solution.
Atmospheric CO2 is in equilibrium with aqueous carbonic acid (H2CO3); the
concentration of CO2 in water is about 1.5x10-5 M at room temperature. The
quantities of acids and bases used in the procedure described in this section are
sufficient that such small concentration of carbonic acid leads to no significant error.
On the other hand, the distilled water is sometimes supersaturated with the CO2 and
thus contains sufficient carbonic acid to cause detectable errors. To test the water to
be used in neutralization titration, take about 500 ml of the distilled water from the
source, add 5 drops of phenolphthalein, and titrate it with 0.1 M NaOH. Less than
0.2 to 0.3 mL of base should be needed to form the first faint pink color. If a larger
volume is required, standard solutions should be prepared from water that has been
boiled briefly to remove CO2 and cooled to room temperature. Furthermore, the
water used to dissolve and dilute the samples should be also boiled and cooled;
alternatively, corrections can be made by blank titration with identical composition
7
but the sample. The instructions that follow assume that dissolved carbon dioxide
can be neglected.
Analysis of a mixture of carbonate and bicarbonate involves two titrations. First,
total alkalinity (= [HCO3 -] + 2[CO32-]) is measured by titrating the mixture with
standard HCl to a bromocresol green end point:
HCO3- + H+ → H2CO3
CO32- + 2H+ → H2CO3
A separate aliquot of unknown is treated with excess standard NaOH to convert
HCO3 - to CO32-:
HCO3- + OH- → CO32- + H2O
Then all the carbonate is precipitated with BaCl2:
Ba2+ + CO32- → BaCO3(s)
The excess NaOH is immediately titrated with standard HCl to determine how
much HCO3- was present. A blank titration is required for the Ba 2+ addition because
Ba2+ precipitates some of the NaOH:
Ba2+ + 2OH- → Ba(OH)2(s)
From the total alkalinity and bicarbonate concentration, you can calculate the
original carbonate concentration.
Reagents
1. Primary standard grade sodium carbonate
2. Bromocresol green indicator solution
8
3. NaCl
4. Primary standard grade potassium hydrogen phthalate (KHP)
5. Phenolphthalein indicator solution
6. 0.1 M NaOH and 0.1 M HCl
7. 10% (w/w) aqueous BaCl2
8. Unknown solutions
Procedure
A. Preparation of 0.1 M HCl solution
Add 1.3~1.5 mL of concentrated HC1 to approximately 150 mL of distilled
water. Mix thoroughly, and store in a glass stoppered bottle.
Note: For very diluted HCl (<0.05 M) solutions, it is advisable to eliminate CO2
from the water by a preliminary boiling.
B. Standardization of HCl against Na2CO3
1. Dry a quantity of primary standard sodium carbonate for 2 hours at 110 oC, and
cool in a desiccator. Weigh two 0.10~ 0.12 g of sodium carbonate (to l mg) into
250-mL Erlenmeyer flasks, and dissolve each in about 50 mL of distilled water.
Introduce 3 drops of bromocresol green, and titrate the solutions with HCl
solution until the color of the solution changes from blue to green. Boil the
solution for 2 to 3 minutes. Cool down to room temperature (Note 1), and
complete the titration with the HC1 solution.
2. Determine an indicator correction by titrating 50 ml of 0.05 M NaCl (Note 2)
9
and 3 drops of the indicator. Subtract the volume of the blank from the titration
data.
Notes
1. The color of the solution should change from green to blue as a result of loss of
CO2 in the heating step. If it does not, it means the amount of acid added is
excess at beginning. This excess can be back-titrated with a base, provided it has
been standardized; otherwise, the sample must be discarded.
2. Weigh 0.14~0.15g NaCl and dissolved in 50 mL of distilled water.
C. Preparation of carbonate-free 0.1 M NaOH solution
Weigh 0.4~0.5 g of NaOH pellets into a 250-mL beaker. Dissolve the NaOH
pellets with 100 mL distilled water (if the water contains significant amounts of
carbonic acid, it should be boiled and cooled) and mix well. Protect the solution
from unnecessary contact with the atmosphere.
D. Standardization of NaOH against potassium hydrogen phthalate (KHP)
1. Dry a quantity of primary standard potassium hydrogen phthalate (KHP,
KHC8H4O4) for 2 hours at 110 oC and let it cool in a desiccator.
2. Weigh two 0.2~0.3 g of KHP (to l mg) into two 250-mL Erlenmeyer flasks and
dissolve in 50 mL of distilled water. Add 2 drops of phenolphthalein, and titrate
the solutions with NaOH until the pink color persists for 30 s.
3. Determine an indicator correction by titrating 50 mL of distilled water with 2
10
drops of phenolphthalein. Subtract the volume of the blank from the titration
data.
E. Analysis of a mixture of carbonate and bicarbonate
1. Total alkalinity: Pipet a 10.0 mL aliquot of unknown solution into a 250-mL
Erlenmeyer flask and titrate with standard 0.1 M HCl, using bromocresol green
indicator as above for standardizing HCl. Boil the solution to yield a sharper
endpoint.
2. Bicarbonate content: Pipet 10.0mL of unknown and 20.00 mL of standard 0.1
M NaOH into a 250-mL Erlenmeyer flask. Swirl and add 5 mL of 10% (w/w)
BaCl2, using a graduated cylinder. Swirl again to precipitate BaCO3, add 2 drops
of phenolphthalein indicator, and immediately titrate with standard 0.1 M HCl
until the red color just disappeared.
3. Blank titration of Ba2+: Pipet 10.0 mL of distilled water, and add 20.00 mL of
standard 0.1 M NaOH 5 mL of 10% (w/w) BaCl2, using a graduated cylinder.
Swirl well to precipitate BaCO3, add 2 drops of phenolphthalein indicator, and
immediately titrate with standard 0.1 M HCl until the red color just disappeared.
The difference in moles of HCl needed in steps 3 and 2 equals the moles of
bicarbonate in the mixture.
4. From the result of step 1, calculate the total alkalinity. From the results of steps
2 and 3, calculate the bicarbonate concentration. Finally, calculate the
concentration of carbonate in the sample. Express the wt% compositions of
11
Na2CO3 and NaHCO3 in the solid unknown.
References
1. Skoog, West, Holler and Crouch, Fundamentals of Analytical Chemistry, 8th ed.,
Chap. 14 & 16
2. D.C. Harris, Quantitative Chemical Analysis, 7th ed., Chap. 11
12
Data
B. Standardization of HCl against Na2CO3
Titration
I
II
Weight of Na2CO3
g
g
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of HCl
mL
mL
Volume of HCl used in blank titration
mL
mL
Volume of HCl used in this titration
mL
mL
Molarity of HCl
M
M
Average Molarity of HCl
M
Calculation
Weight of Na2CO3
1
 MHCl  VmL, HCl 
Molar mass of Na2CO3
2000
13
D. Standardization of NaOH against potassium hydrogen phthalate (KHP)
Titration
I
Weight of KHP
II
g
g
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of NaOH
mL
mL
Volume of NaOH used in blank titration
mL
mL
Volume of NaOH used in this titration
mL
mL
Molarity of NaOH
M
M
Average Molarity of NaOH
M
Calculation
Weight of KHP
1
 MNaOH  VmL, NaOH 
Molar mass of KHP
1000
14
E. Analysis of a mixture of carbonate and bicarbonate
1. Total alkalinity
Unknown No.
Concentration of sample
g/mL
Initial buret reading
mL
Final buret reading
mL
Volume of HCl
mL
Total alkalinity
mol
Calculation
Total alkalinity (mol)  MHCl  VmL, HCl 
1
1000
2. Bicarbonate content
Initial buret reading
mL
Final buret reading
mL
Volume of HCl
mL
3. Blank titration of Ba2+
Initial buret reading
mL
Final buret reading
mL
Volume of HCl
mL
15
5. Data calculation
the wt% content of NaHCO3 in the sample
%
Accurate value of %NaHCO3
%
Relative error
%
the wt% content of Na2CO3 in the sample
%
Accurate value of %Na2CO3
%
Relative error
%
Calculation
wt % of NaHCO3 
wt % of Na2CO3 
molar mass of NaHCO3
1000
100%
Concentration of sample (g/mL)  VmL
MHCl  (V3, mLl - V2, mL) 
molar mass of Na2CO3
2000
100%
Concentration of sample (g/mL)  VmL
[(MHCl  V1, mL) - MHCl  (V3, mL - V2, mL)] 
Discussion
16
Experiment 3
Potentiometric Titration Using a pH Electrode
Introduction
Titrations are most commonly performed either to find out how much analyte
is present or to measure equilibrium constants of the analyte. We can obtain the
information necessary for both purposes by monitoring the pH of the solution as the
titration is performed. In this experiment you will use a pH electrode to follow the
course of an acid-base titration. You will observe how pH changes slowly during
most of the reaction and rapidly near the equivalence point. You will compute the
first derivatives of the titration curve to locate the end point. From the mass of
unknown acid or base and the moles of titrant, you can calculate the molecular mass
of the unknown.
Reagents
1. Standardized 0.1 M HCl
2. Standardized 0.1 M NaOH
Procedure
1. Your instructor will recommend an appropriate mass of unknown (5~8 mmol)
for you to weigh accurately and dissolve in distilled water in a 100-mL
volumetric flask. Dilute to the mark and mix well.
2. Make a rough test of the pH of your unknown solution to ascertain whether it is
an acid or a base. If your unknown is a base, follow Step 3 below. Otherwise,
follow Step 4 below.
17
3. Preparation and standardization of 0.1 M HCl – follow Procedure A and B of
Experiment 2.
4. Preparation and standardization of 0.1 M NaOH – follow Procedure C and D of
Experiment 2.
5. The first titration is intended to be rough, so that you will know the approximate
end point in the next titration. For the rough titration, pipet 20.0 mL of unknown
into a 250-mL Erlenmeyer flask. If you are titrating an unknown acid, add 2
drops of phenolphthalein indicator and titrate with standard 0.1 M NaOH to the
pink end point, using a 50-mL buret. If you are titrating an unknown base, add 3
drops of bromocresol green indicator and titrate with standard 0.1 M HCl to the
green end point.
6. Following instructions for your particular pH meter, calibrate a meter and glass
electrode, using buffers with pH values near 7 and 4. Rinse the electrodes well
with distilled water and blot them dry with a tissue before immersing in any new
solution.
7. Now comes the careful titration, pipet 40.0 mL of unknown solution into a
100-mL beaker containing a magnetic stirring bar. Position the pH electrode in
the liquid so that the stirring bar will not strike the electrode. If a combination
electrode is used, the small hole near the bottom on the side must be immersed in
the solution. This hole is the salt bridge to the reference electrode. Allow the
electrode to equilibrate for 1 min with stirring and record the pH.
8. Add 1~2 drop of indicator and begin the titration. The equivalence volume will
18
be two times greater than it was in Step 5. Add ~2.0 mL aliquots of titrant and
record the exact volume, the pH, and the color 30 s after each addition. When
you are within 1.0 mL of the indicator end point, add titrant in 0.1 mL
increments. Continue with 0.1 mL increments until you are 1.0 mL past the
indicator end point. The equivalence point has the most rapid change in pH.
Then, add five more 2.0 mL aliquots of titrant and record the pH after each.
Data analysis
1. Construct a graph of pH versus titrant volume. Make on your graph where the
indicator color change was observed.
2. Compute the first derivative (the slope,
ΔpH/ΔV)
for each data point within
±
0.5 mL of the indicator end point volume. From your graph, estimate the
equivalence volume (Ve) as accurately as you can, as shown below.
3. Compare the indicator end point volume to the equivalence volume (Ve)
estimated from the first derivative.
4. From the equivalence volume (Ve) and the mass of unknown, calculate the
19
molecular mass of the unknown.
5. From the pH value at halfway the equivalence volume (1/2Ve), find the
approximate pKa value of the unknown acid (or the conjugate acid of the
unknown base).
References
1. Skoog, West, Holler and Crouch, Fundamentals of Analytical Chemistry, 8th
ed., Chap. 14 & 15
2. D.C. Harris, Quantitative Chemical Analysis, 7th ed., Chap. 11
20
Data
1. Sampling
Unknown No.
Weight of sample
g
3. Standardization of HCl against Na2CO3
Titration
I
II
Weight of Na2CO3
g
g
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of HCl
mL
mL
Volume of HCl used in blank titration
mL
mL
Volume of HCl used in this titration
mL
mL
Molarity of HCl
M
M
Average Molarity of HCl
M
Calculation
Weight of Na2CO3
1
 MHCl  VmL, HCl 
Molar mass of Na2CO3
2000
21
4. Standardization of NaOH against potassium hydrogen phthalate (KHP)
Titration
I
Weight of KHP
II
g
g
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of NaOH
mL
mL
Volume of NaOH used in blank titration
mL
mL
Volume of NaOH used in this titration
mL
mL
Molarity of NaOH
M
M
Average Molarity of NaOH
M
Calculation
Weight of KHP
1
 MNaOH  VmL, NaOH 
Molar mass of KHP
1000
22
5. Rough titration
Initial buret reading
mL
Final buret reading
mL
Volume of NaOH or HCl
mL
8. Careful titration
Estimated volume of end point for 40 mL unknown sol’n
V(mL)
pH
color
0
2.00
4.00
6.00
V(mL)
pH
color
V(mL)
pH
color
V(mL)
pH
color
V(mL)
pH
color
V(mL)
pH
color
V(mL)
pH
color
23
mL
Data analysis
1. Construct a graph of pH(y-axis) versus titrant volume(x-axis). Make on your
graph where the indicator color change was observed.
24
2. Compute the first derivative (the slope,
ΔpH/ΔV)
for each data point within
±
0.5 mL of the indicator end point volume. From your graph, estimate the
equivalence volume (Ve) as accurately as you can.
V(mL)
pH
△pH
△V
△pH/△V
Vavg
( Graph: △pH/△V v.s Vavg )
25
3. Compare the indicator end point volume to the equivalence volume (Ve)
estimated from the first derivative.
4. From the equivalence volume (Ve) and the mass of unknown, calculate the
molecular mass of the unknown.
5. From the pH value at halfway the equivalence volume (1/2Ve), find the
approximate pKa value of the unknown acid (or the conjugate acid of the unknown
base).
Discussion
26
Experiment 4
Complex Titrations with EDTA
Introduction
In a direct titration, analyte is titrated with standard EDTA. The analyte is
buffered to a pH at which the conditional formation constant for the metal-EDTA
complex is large and the color of the free indicator is distinctly different from that of
the metal-indicator complex.
An auxiliary complexing agent, e.g. NH3, may be employed to prevent the metal
ion from precipitating in the absence of EDTA. For example, the direct titration of
Pb2+ is carried out in ammonium buffer at pH 10 in the presence of tartrate, which
complexes the metal ion and does not allow Pb(OH)2 to precipitate. The Pb-tartrate
complex must be less stable than the Pb-EDTA complex, or the titration would not
be feasible.
In a back titration, a known excess of EDTA is added to the analyte. The excess
EDTA is then titrated with a standard solution of a second metal ion. A back titration
is necessary if the analyte precipitates in the absence of EDTA, if it reacts too
slowly with EDTA under titration conditions, or if it blocks the indicator. The metal
ion used in the back titration must not displace the analyte metal ion from its EDTA
complex.
A solution of the Mg(EDTA)2- complex is useful for the titration of some cations
which form more stable complexes with EDTA than the Mg(EDTA) 2- complex but
for which no indicator is available. Mg2+ ions in the complex are displaced by a
27
chemically equivalent quantity of the analyte cations. The liberated Mg2+ ions are
then titrated with standard EDTA solution. Eriochrome black T, or Calmagite can
serve as an indicator.
Mn+ + Mg(EDTA)2- → M(EDTA)n-4 + Mg2+
Note that the concentration of the Mg(EDTA)2- solution is not important; all that
is necessary is that the molar ratio between Mg2+ and EDTA be exactly unity in the
reagent.
Reagents
1. Standard 0.01 M EDTA solution
2. 0.01 M MgSO4 solution
3. Eriochrome black T indicator solution
4. NH3/NH4Cl (pH 10) buffer
5. standard 0.01 M Mg2+ solution
6. 12 M HC1
7. 6 M NaOH
8. Methyl red indicator solution
9. 0.1 M Mg(EDTA)2- solution
Procedure
A. Preparation of standard 0.01 M EDTA solution
Dry Na2H2EDTA‧2H2O (F.M. 372.25) at 80 oC for 1 hour to remove superficial
moisture and cool in a desiccator. Accurately weigh out 0.37~0.38 g and put into a
250-mL beaker with 50 mL of distilled water. Stir periodically until the EDTA is
28
dissolved and then cover with a watch glass. It takes a while to dissolve EDTA
completely. Examine the bottom of the flask for undissolved salt; if there is any
remaining, continue the periodic shaking until solution is complete. Transfer all the
solution to a 100-mL volumetric flask. Dilute to the mark and mix well. Calculate
the molarity of EDTA solution.
B. Standardization of 0.01 M Mg2+ solution by direct titration
1. Dissolve 0.25 g of MgSO4‧7H20 in a 250-mL beaker. Dilute with 100 mL of
distilled water and mix well.
2. Pipette two 15.0 mL aliquots of MgSO4 solution to 250-mL Erlenmeyer flasks,
add 2 mL of pH 10 buffer, 2 drops of Eriochrome black T indicator and titrate
with 0.01 M EDTA until the solution turns from red to blue (Note).
3. Calculate the molarity of the Mg2+ solution.
Note
The color change tends to be slow in the vicinity of the end point. Care must be
taken to avoid overtitration.
C. Complexometric determination of Ca2+ by back titration
1. Weigh 0.10~0.12 g of the Ca2+ contained sample into a 250-mL beaker (Note 1).
Cover with a watch glass, and carefully introduce 3.5 mL of 12 M HCl. After the
sample has been dissolved, add 50 mL of distilled or deionized water and boil
gently for a few minutes to remove CO2.
29
2. Cool the solution, and neutralize with 6 M NaOH. (Note 2)
3. Transfer the solution to a 100-mL volumetric flask, and dilute to the mark. Mix
well.
4. Take two 15.0 mL aliquots solution for titration. To each 15.0 mL aliquot add
about 2 mL of pH 10 buffer and 2 drops of Eriochrome black T indicator.
5. Run in an excess of 0.01 M EDTA solution from a buret, and record the volume
taken. The color of the solution changes from red to blue. Titrate the excess
EDTA with standard 0.01 M MgSO4 solution until a color change from blue to
red. Calculate the wt% of Ca in the sample.
Notes
1. The sample taken should contain about 150 to 160 mg of calcium ion.
2. To neutralize the solution, introduce a few drops of methyl red, and add base
until the red color is discharged.
D. Complexometric determination of Ca2+ by displacement titration
Take two 15.0 mL aliquots of the above Ca2+ solution for titration and treat each
as the followings: add 2 mL of pH 10 buffer, 1 mL of the Mg(EDTA)2- solution, and
2 drops of Eriochrome black T indicator. Titrate with standard 0.01 M EDTA
solution until a color change from red to blue. Calculate the number of milligrams
of Ca in the sample.
30
E. Determination of water hardness
1. Acidify two 100.0 mL aliquots of tap water with a few drops of 12 M HCl, and
boil gently for a few minutes to remove CO2.
2. Let the solution cool and add 2 drops of methyl red, and neutralize the solution
with 6 M NaOH.
3. Introduce 2 mL of pH 10 buffer and 4 drops of Eriochrome black T indicator,
and titrate with standard 0.01 M EDTA until the color changes from red to blue
(Note).
4. Report the results of the analysis in the units of milligrams of calcium carbonate
per liter of water.
Note
The color change is sluggish if Mg is absent. In this event, add 1 to 2 mL of the
0.1 M Mg(EDTA)2- solution before starting the titration.
References
1. Skoog, West, Holler and Crouch, Fundamentals of Analytical Chemistry, 8th ed.,
edition, Chap. 17
2. D.C. Harris, Quantitative Chemical Analysis, 7th ed., Chap. 12
31
32
Data
A. Preparation of standard 0.01 M EDTA．2Na solution
g
Weight of EDTA．2Na
Volume of EDTA．2Na solution prepared
mL
Molarity of EDTA．2Na
M
Calculation
Molarity of EDTA  2Na =
Weight of EDTA  2Na
1000

Molar mass of EDTA  2Na Volume (mL) of solution
B. Standardization of 0.01 M Mg2+ solution by direct titration
Titration
I
II
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of EDTA．2Na
mL
mL
Average volume of EDTA．2Na
mL
Average Molarity of Mg2+
M
Calculation
M EDTA．2Na x VmL, EDTA．2Na =
MMgSO4 x VmL, MgSO4
33
C. Complexometric determination of Ca2+ by back titration
Unknown No. _________
Weight of sample: ___________g
Titration
I
II
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of EDTA．2Na
mL
mL
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of MgSO4
mL
mL
the wt% content of Ca in the sample
%
%
the average wt% content of Ca in the sample
%
Accurate value of %Ca
%
Relative error
%
Calculation
[(MEDTA  2Na  VmL,EDTA  2Na)-(MMgSO4  VmL,MgSO4)] 
wt % of Ca =
Weight of sample 
34
Atomic mass of Ca
1000
100%
VmL,used
VmL,total
D. Complexometric determination of Ca2+ by displacement titration
Titration
I
II
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of EDTA．2Na
mL
mL
Average volume of EDTA．2Na
mL
the wt% content of Ca in the sample
%
Accurate value of %Ca
%
Relative error
%
Calculation
Atomic mass of Ca
1000
100%
VmL,used
Weight of sample 
VmL,total
MEDTA  2Na  VmL,EDTA  2Na 
wt % of Ca =
35
E. Determination of water hardness
Titration
I
II
Volume of tap water
mL
mL
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of EDTA．2Na
mL
mL
Average volume of EDTA．2Na
mL
ppmCaCO3
ppm
Calculation
Atomic mass of CaCO3
1000mg
1000
1 L
Volume of sample(mL) 
1000 mL
MEDTA  2Na  VmL,EDTA  2Na 
ppm CaCO3 =
Discussion
36
Experiment 5
Precipitation Titrations with Silver Nitrate
Introduction
Silver nitrate (AgNO3) can be obtained in primary standard purity. It has a high
equivalent weight and dissolves readily in water. The solid as well as its solution
must be scrupulously protected from organic matter and from sunlight; elemental
silver is the product of reduction by the former and photodecomposition by the latter.
The reagent is expensive, and every effort should be made to avoid waste. Unused
solution should be collected rather than discarded; similarly, appreciable amount of
silver chloride should also be collected.
The Volhard titration is actually a procedure for titrating Ag +. To determine Cl-,
a back titration is necessary. First, Cl- is precipitated by a known, excess quantity of
standard AgNO3.
Ag+ + Cl- → AgCl(s)
The AgCl is isolated, and excess Ag+ is titrated with standard KSCN in the presence
of Fe3+.
Ag+ + SCN- → AgSCN(s)
When all Ag+ has been consumed, SCN- reacts with Fe3+ to form a red complex.
Fe3+ + SCN- → FeSCN2+
The appearance of red color is the end point. Knowing how much SCN- was
required for the back titration tell us how much Ag + was left over from the reaction
37
with Cl-. Because the total amount of Ag+ is known, the amount consumed by Clcan then be calculated.
In the analysis of Cl- by the Volhard method, the end point slowly fades because
AgCl is more soluble than AgSCN. The AgCl slowly dissolves and is replaced by
AgSCN. To eliminate this secondary reaction, you can filter off the AgCl and titrate
only the Ag+ in the filtrate. Alternatively, a few milliliters of the water-immiscible
nitrobenzene can be added to the solution to cover the AgCl precipitate with a layer
of organic agent, thus avoiding its contact with SCN-. Br- and I-, whose silver salts
are less soluble than AgSCN, can be titrated by the Volhard method without
isolation the silver halide precipitate.
Because the Volhard method is a titration of Ag+, it can be adapted for the
determination of any anion that forms an insoluble silver salt.
In the Fajans titration, an adsorption indicator is applied along with the electrical
phenomena associated with precipitate formation. In the early addition of Ag + to the
solution of Cl-, there will be excess chloride ions in solution prior to the equivalence
point. Some Cl- is selectively adsorbed on the AgCl(s) surface, imparting a negative
charge to the crystal surface (Figure (a) below). After the equivalence point,
adsorption of the excess silver cations on the crystal surface creates a positive
charge on the particles of precipitate. The abrupt change from negative charge to
positive charge occurs at the equivalence point.
38
Common adsorption indicators are anionic dyes, which are attracted to the
positively charged particles of precipitate produced immediately after the
equivalence point (Figure (b) above). The adsorption of the negatively charged dye
on the positively charged surface changes the color of the dye by interactions that
are not well understood. The color change signals the end point in the titration.
Because the indicator reacts with the precipitate surface, it is desirable to have as
much surface area as possible. The indicator most commonly used for AgCl is
dichlorofluorescein.
In all ‘argentometric titrations’, but especially with adsorption indicators, strong
light should be avoided.
Reagents
1. AgNO3(s)
2. NaCl(s)
3. Dextrin
39
4. Dichlorofluorescein indicator
5. 16 M HNO3 solution
6. Indicator: 0.1 M iron(III) ammonium sulfate
7. Standard 0.1 M potassium thiocyanate solution
Procedure
A. Preparation of 0.1 M AgNO3 solution
Dry AgNO3(s) at 110 oC for 1 hour but not much longer (Note 1), cool down to
room temperature in a desiccator. Weigh approximately 1.70 g of the dry AgNO3 to
a clean and dry weighing bottle. Weigh the paper and its contents to 0.001 g.
Quickly transfer the contents to a 100-mL volumetric flask. Weigh the weighing
paper again and record the weight of AgNO3(s). Dissolve the AgNO3 in distilled
water and dilute to the mark. Mix well before use (Note 2).
B. Standardization of 0.1 M AgNO3 solution
Dry a quantity of standard NaCl(s) for 1 hour at 110 oC, and cool in a desiccator.
Weigh 0.05~0.06 g NaCl into a 250-mL Erlenmeyer flask and dissolve it in 20 mL
of distilled water. Add about 0.05 g of dextrin and 4 drops of dichlorofluorescein
indicator. Titrate the solution with AgNO3 solution to the first permanent
appearance of the pink color of the indicator.
Notes
1. AgNO3 is perceptibly decomposed by prolonged heating. Some discoloration
40
may occur, even after 1 hour at 110 oC. The effect on the purity is ordinarily
negligible.
2. AgNO3 solution should be stored in a dark place when not actually in use.
C. Determination of Cl- by the Volhard method
1. Preparation of 0.1 M potassium thiocyanate
Dissolve 0.45~0.50 g of KSCN and dilute it to 50 mL with distilled water.
2. Standardization of 0.1 M potassium thiocyanate
Accurately transfer 10 mL (to 0.01 mL) of standard AgNO3 solution into a
250-mL Erlenmeyer flask, followed by the addition of 1.5 mL of iron(III)
ammonium sulfate indicator. A 1.0 mL of 16 M HNO3 is added into the solution.
Titrate the solution with the KSCN solution, swirling the flask vigorously until the
red-brown color of FeSCN2+ is present for 1 minute. Find the concentration of the
KSCN solution.
3. Determination of ClDry the unknown at 100 ~ 110 oC for 1 hour. Weigh two 0.05~0.06 g of samples
(to 1 mg) into numbered 250-mL Erlenmeyer flasks. Dissolve each sample in 10 mL
of distilled water, and acidify with 1.0 mL of 16 M HNO3. Introduce an excess of
standard AgNO3 solution, and be sure to note the volume taken. Shake
vigorously and filter. Add 1.5 mL of iron(III) indicator in the filtrate. Titrate the
excess silver ions that are in filtrate with standard KSCN solution until the red color
of FeSCN2+ is present for 1 minute. Calculate the wt% content of Cl in the
41
unknown.
Notes
1. With the concurrence of the instructor, a larger quantity of unknown can be
weighed into a volumetric flask and diluted to known volume. The
determination can then be made upon aliquot portions of this solution.
2. To obtain an approximation of the volume of standard AgNO 3 that constitutes
an excess, calculate the amount that would be required for one of the samples
assuming that it is 100% NaCl. When actually adding silver ions solution to the
sample, swirl the flask vigorously, and add 3 to 4 mL in excess of the volume
required to cause the AgCl to coagulate.
3. HNO3 is introduced to improve observation of the end point. Since the lower
oxides of nitrogen tend to attack SCN-, the acid should be freshly boiled.
4. At the outset of the back-titration, an appreciable quantity of silver ions is
adsorbed on the surface of the precipitate. As a result, there is a tendency for a
premature appearance of the end point color. Since success of the method
depends upon an accounting for all excess silver ions, thorough and vigorous
agitation is essential to bring about desorption of these ions from the precipitate.
A magnetic stirrer is helpful for this purpose.
D. Determination of Cl- by the Fajans method
Dry the unknown for 1 hour at 110 oC, and cool in a desiccator. Weigh
accurately 0.05~0.06 g sample into a 250-mL Erlenmeyer flask and dissolve in 10
42
mL of distilled water. Add about 0.05 g of dextrin and 4 drops of
dichlorofluorescein indicator, titrate with AgNO3 to the first permanent appearance
of the pink color of the indicator. Calculate the wt% content of Cl in the unknown.
References
1. Skoog, West, Holler and Crouch, Fundamentals of Analytical Chemistry, 8th
ed., pp. 353-363
2. D.C. Harris, Quantitative Chemical Analysis, 7th ed., pp. 127-134
43
Data
B. Standardization of 0.1 M AgNO3 solution
Titration
I
II
Weight of NaCl
g
g
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of AgNO3
mL
mL
Molarity of AgNO3
M
M
Average Molarity of AgNO3
M
Calculation
44
C. Determination of Cl- by the Volhard method
2. Standardization of 0.1 M potassium thiocyanate
Titration
I
II
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of KSCN
mL
mL
Average volume of KSCN
mL
Molarity of KSCN
M
Calculation
45
C. Determination of Cl- by the Volhard method
3. Determination of ClTitration
I
II
Unknown No.
Weight of sample
g
g
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of AgNO3
mL
mL
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of KSCN
mL
mL
the wt% content of Cl in the unknown
%
%
the average wt% content of Cl in the
%
unknown
Accurate value of %Cl
%
Relative error
%
Calculation
46
D. Determination of Cl- by the Fajans method
Titration
I
II
Unknown No.
Weight of sample
g
g
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of AgNO3
mL
mL
the wt% content of Cl in the unknown
%
%
the average wt% content of Cl in the unknown
%
Accurate value of %Cl
%
Relative error
%
Calculation
47
Discussion
48
Experiment 6
Analysis of Copper in Ore by Iodometric Method
Introduction
Iodometry is a method based on the reducing properties of I-:
2I- → I2 + 2eThe reaction product, I2, is then titrated with a standard Na2S2O3 solution, with
starch serving as the indicator:
I2 + 2S2O32- → 2I- + S4O62The following directions will permit the determination of Cu in samples of ore
iodometrically.
Cu(II) is quantitatively reduced to Cu(I) by I-. The reaction can be expressed as
the following:
2Cu2+ + 4I-
→
2CuI(s) + I2
The importance of CuI formation in forcing this reaction to completion can be seen
from the following standard electrode potentials:
Cu2+ + eI2 + 2e-
⇄
⇄
Cu+
Eo = 0.15 V
21-
Cu2+ + I- + e-
⇄
Eo = 0.54 V
Eo = 0.86 V
CuI(s)
The first two potentials suggest that I- should have no tendency to reduce Cu(II); the
formation of CuI, however, favors the reduction reaction. The solution must contain
at least 4% excess I- to force the reaction to completion. Moreover, the pH must be
49
less than 4 to prevent the formation of basic Cu species that react slowly and
incompletely with I-. The acidity of the solution cannot be greater than about 0.3 M,
however, because of the tendency of I- to undergo air oxidation, a process catalyzed
by Cu salts. Nitrogen oxides (NOx) also catalyze the air oxidation of I-. A common
source of these oxides is the HNO3 ordinarily used to dissolve metallic Cu and other
Cu-containing solids. Urea is used to scavenge NOx from solutions:
(NH2)2CO + 2HNO2
→
2N2 + CO2 + 3H2O
The titration of I2 by S2O32- tends to yield slightly low results when CuI is
present, owing to the adsorption of small but measurable quantities of I2 upon the
solid. The adsorbed I2 is released only slowly, even when S2O32- is in excess;
transient and premature end points result. This difficulty is largely overcome by the
addition of SCN-. The sparingly soluble Cu(I) thiocyante replaces part of the CuI at
the surface of the solid:
CuI(s) + SCN-
→
CuSCN(s) + I-
Accompanying this reaction is the release of the adsorbed I2, which thus becomes
available for titration. The addition of SCN- must be delayed until most of the I2 has
been titrated to prevent interference from a slow reaction between the two species,
possibly
2SCN- + I2
→
2I- + (SCN)2
Some samples require the addition of HCl to complete the solution step. The Clmust then be removed by evaporation with sulfuric acid because I- will not reduce
Cu(II) quantitatively from its chloro-complexes.
50
Of the elements ordinarily associated with Cu in nature, only Fe, As, and Sb
interfere with the iodometry. Fortunately, difficulties caused by these elements are
readily eliminated. Iron is rendered unreactive by the addition of such complexing
agents as fluoride or pyrophosphate; because these ions form more stable complexes
with Fe(III) than with Fe(II), the potential for this system is altered to the point
where appreciable oxidation of I- cannot occur. Interference by As and Sb is
prevented by converting these elements to the +5 state during the solution step.
Ordinarily, the hot HNO3 used to dissolve the sample will convert them to the
desired oxidation state, although a small amount of Br2 water can be added in case
of doubt; the excess Br2 is then expelled by boiling. As pointed out, As in the +5
state does not oxidize I-, if the provided solution is not too acidic. Sb behaves
similarly. Thus, the pH of the solution maintains at 3 or greater, interference by
these elements can be avoided. We have seen, however, that oxidation of I- by
copper is incomplete at pH values greater than 4. Thus, when Cu is to be determined
in the presence of As or Sb, it is essential to control the pH between 3 and 4. The
anion of the salt dissociates as follows:
HF2HF
→
→
HF + F-
K = 0.26
H＋ + F-
K = 7.2x10 -4
-
The first dissociation provides equa1 quantities of HF and F- which then buffer the
solution to a pH somewhat greater than 3. In addition to acting as a buffer, the salt
also serves as a source of F- to complex any Fe(III) that may be present.
51
Reagents
1. Sodium thiosulfate
2. Sodium carbonate
3. Primary standard grade potassium iodate
4. Potassium iodide
5. 12 M HCl
6. Starch indicator solution
7. Concentrated nitric acid
8. Concentrated ammonia solution
9. 5% (w/v) urea solution
10. Ammonium bifluoride
11. Potassium thiocyanate
Procedure
A. Preparation of 0.1 M sodium thiosulfate solution
Boil about 100 mL of distilled water for at least 5 minutes. Cool, and add about
2.5 g of Na2S2O3∙5H2O and 0.1 g of Na2CO3. Stir until solution is complete; then,
transfer to a clean stoppered bottle (glass or plastic), and store in the dark.
B. Standardization of sodium thiosulfate against potassium iodate
1. Dry primary standard grade potassium iodate for at least 1 hour at 110 oC and
cool in a desiccator. Weigh 0.2 g (to 1 mg) of potassium iodate sample into a
52
100-mL volumetric flask, dissolving in water, and diluting to the mark.
2. Pipet two 25.0 mL aliquots KIO3 solutions into two separate 250-mL
Erlenmeyer flasks. Add about 7 mL of 20% KI to each sample. (Treat each
sample individually from this point in order to minimize errors resulting from air
oxidation of iodide in the acidic solution).
3. Add 0.7 mL of 12 M HC1, and titrate immediately with sodium thiosulfate until
the color of the solution becomes pale yellow. Add 2 mL of starch solution, and
titrate till the disappearance of the blue color. Calculate the concentration of
Na2S2O3 solution.
C. Determination of Cu in an ore
1. Accurately weigh two 0.2~0.25 g of dried and finely ground ore samples into
100-mL beakers. Treat two samples individually. Add 5 mL of concentrated
HNO3 to each beaker, and cover with a watch glass. Carefully heat in a hood
until all Cu is dissolved. If the volume becomes less than 3 mL, add more nitric
acid. Continue the heating until only white or slightly gray siliceous residue
remains (Note 1). Evaporate to about 3 mL.
2. Add 25 mL of distilled water and 5 mL of 5% urea, and boil to dissolve soluble
salts and to expel all nitrogen oxides. If the residue is small and nearly colorless,
no filtration is necessary. Otherwise, filter the suspension, and collect the filtrate
in a 250-mL Erlenmeyer flask. Wash the paper with several small portions of
1:100 hot HNO3 solution; then, discard the filter paper. Evaporate the filtrate and
53
washings to about 5 mL, then let it cool. Add concentrated ammonia solution
slowly to the first appearance of the deep blue Cu(NH3)42+ complex. A faint odor
of ammonia should be detectable over the solution. If it is not, add another drop
of ammonia and repeat the test.
3. From this point, treat each sample individually. Add 0.5 g of ammonium
bifluoride (CAUTION! Note 2), and swirl until completely dissolved. Then, add
5 mL of 20% KI (The color of the solution should be brown, Note 3) and titrate
immediately with 0.1 M standard sodium thiosulfate. When the color of the
iodine is nearly discharged, add 5 mL of 1 M KSCN and 2 mL of starch solution.
Swirl vigorously for several seconds. Continue the titration with vigorous mixing
until the blue color disappears for several minutes.
Notes
1. If the ore is not readily decomposed by the nitric acid, add 5 mL of concentrated
HC1 and heat in a hood until only a small white or gray residue remains. Do not
evaporate to dryness. Cool, add 10 mL of concentrated sulfuric acid, and
evaporate in a hood until copious white fumes of SO3 are observed. Cool and
carefully add 15 mL of water and 10 mL of saturated Br2 water. Boil the solution
vigorously in a hood until all of the Br2 has been removed. Cool and proceed
with the filtration step in the second paragraph.
2. Ammonium bifluoride (or ammonium hydrogen fluoride) is a highly toxic and
corrosive chemical. Avoid contact with the skin. If exposure does occur, rinse
immediately the affected area with copious amounts of water.
54
3. If too much ammonia is added, neutralize the excess with 3 M sulfuric acid
solution.
References
1. Skoog, West, Holler and Crouch, Fundamentals of Analytical Chemistry, 8th ed.,
p.576 and Chap. 37H
2. D.C. Harris, Quantitative Chemical Analysis, 7th ed., pp.340-343
55
56
Data
B. Standardization of sodium thiosulfate against potassium iodate
Weight of KIO3
g
Volume of KIO3 solution prepared
mL
Molarity of KIO3
M
Calculation
Titration
I
II
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of Na2S2O3
mL
mL
Average volume of Na2S2O3
mL
Molarity of Na2S2O3
M
Calculation
57
C. Determination of Cu in an ore
Titration
I
II
Unknown No.
Weight of sample
g
g
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of Na2S2O3
mL
mL
the wt% of Cu in the sample
%
%
the average wt% of Cu in the sample
%
Accurate value of %Cu
%
Relative error
%
Calculation
Discussion
58
Experiment 7
Determination of Ascorbic Acid in Vitamin C Tablets by
Titrations with Potassium Bromate
Introduction
Primary-standard potassium bromate is available from commercial sources and
can be used directly to prepare standard solutions that are stable indefinitely. Direct
titrations with KBrO3 are relatively few. Instead, the reagent is a convenient and
widely used stable source of Br2. In this application, an unmeasured excess of KBr
is added to an acidic solution of the analyte. On introduction of a measured volume
of standard KBrO3, a stoichiometric quantity of Br2 is produced.
BrO3- + 5Br- + 6H+ → 3Br2 + 3H2O
This indirect generation circumvents the problems associated with the use of
standard Br2 solutions, which lack stability.
The primary use of standard KBrO3 is the determination of organic compounds
that react with Br2. Few of these reactions are rapid enough to make direct titration
feasible. Instead, a measured excess of standard BrO3- is added to the solution that
contains the sample plus an excess of KBr. After acidification, the mixture is
allowed to stand in a glass-stoppered vessel until the Br2/analyte reaction is judged
complete. To determine the excess Br2, an excess of KI is introduced so that the
following reaction occurs:
2I- + Br2 → I2 + 2BrThe liberated I2 is then titrated with standard Na2S2O3 solution, with starch serving
59
as the indicator:
I2 + 2S2O32- → 2I- + S4O62The standard BrO3- solution is employed to generate a known amount of I 2 that
is used to standardize the solution of Na2S2O3.
BrO3- + 6I- + 6H+ → Br- + 3I2 + 3H2O
Ascorbic acid is cleanly oxidized to dehydroascorbic acid by bromine:
An unmeasured excess of KBr is added to an acidified solution of the sample. The
solution is titrated with standard KBrO3 to the first permanent appearance of excess
Br2; this excess is then determined iodometrically with standard Na2S2O3. The entire
titration must be performed without delay to prevent air oxidation of the ascorbic
acid.
Reagents
1. Na2S2O3∙5H2O(s)
2. Na2CO3(s)
3. KBrO3(s)
4. KI(s)
60
5. 3 M H2SO4
6. Starch indicator solution
7. KBr(s)
Procedure
A. Preparation of 0.05 M sodium thiosulfate solution
Boil about 100 mL of distilled water for at least 5 minutes. Cool, and add about
1.25 ~1.30 g of Na2S2O3∙5H2O and 0.1 g of Na2CO3. Stir until solution is complete;
then, transfer to a clean stoppered bottle (glass or plastic), and store in the dark.
B. Standardization of 0.05 M sodium thiosulfate solution
1. Dry reagent grade KBrO3 for 1 hour at 100 ~110 oC, and cool in the desiccator.
Weigh out 0.14 g accurately and transfer to a 100-mL volumetric flask. Dilute to
the mark, and mix thoroughly. Calculate the concentration of BrO3-.
2. Pipet two 20.0 mL aliquots of the KBrO3 solution into 250-mL Erlenmeyer
flasks. Treat each sample individually. Add about 10 mL of 20% KI and 1 mL
of 18 M H2SO4 to one of the flasks, and then titrated the liberated I 2 immediately
until the solution is just faintly yellow. Add 5 mL of starch indicator solution,
and continue the titration until the blue color disappears. Repeat the same
process for the other flask of KBrO3. Calculate the concentration of the Na2S2O3
solution.
61
C. Determination of ascorbic acid in vitamin C tablets by titration with KBrO3
solution
1. Weigh (to the nearest milligram) 3 to 5 vitamin C tablets (Note 1). Pulverize
them thoroughly in a mortar, and transfer the powder to a dry weighing bottle.
Weigh individual 0.20~0.21 g samples (to the nearest 1 mg) into two dry
250-mL Erlenmeyer flasks. Treat each sample individually beyond this point.
2. Dissolve the sample (Note 2) in 20 mL of distilled water, add 5 mL conc.
H2SO4; mix well and add about 12 mL of 20% KBr. Titrate immediately with
standard KBrO3 to the first faint yellow due to excess Br2. Record the volume of
KBrO3 used. Add 7 mL of 20% KI and 2 mL of starch indicator; back-titrate
(Note 3) with standard 0.05 M Na2S2O3.
3. Calculate the average mass (in milligrams) of ascorbic acid (176.12 g/mol) in
each tablet.
Notes
1. This method is not applicable to chewable vitamin C tablets.
2. The binder in many vitamin C tablets remains in suspension throughout the
analysis. If the binder is starch, the characteristic color of the complex with I2
appears on addition of KI.
3. The volume of Na2S2O3 needed for the back-titration seldom exceeds a few
milliliters.
62
References
1. Skoog, West, Holler and Crouch, Fundamentals of Analytical Chemistry, 8th
ed., Chap. 20, pp.577-580
2. D.C. Harris, Quantitative Chemical Analysis, 7th ed., Chap. 16
63
Data
B. Standardization of 0.05 M sodium thiosulfate solution
Weight of KBrO3
g
Volume of KBrO3 solution prepared
mL
Molarity of KBrO3
M
Calculation
Titration
I
II
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of Na2S2O3
mL
mL
Average volume of Na2S2O3
mL
Molarity of Na2S2O3
M
Calculation
64
C. Determination of ascorbic acid in vitamin C tablets by titration with KBrO 3
solution
Titration
I
II
Weight of Vit. C tablet
g/tablet
Weight of Vit. C powder
g
g
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of KBrO3
mL
mL
Initial buret reading
mL
mL
Final buret reading
mL
mL
Volume of Na2S2O3
mL
mL
mg Vit. C / tablet
mg
mg
Average mg Vit. C / tablet
mg
Accurate value of Vit. C / tablet
mg
Relative error
%
Calculation
65
Discussion
66
Experiment 8
Determination of Iron in an Ore by Potentiometric Titration
Introduction
The common iron ores are hematite (Fe2O3), magnetite (Fe3O4), and limonite
(3Fe2O3∙3H2O). Volumetric methods for the analysis of iron samples containing
these substances consist of three steps:
(1) dissolution of the sample
(2) reduction of the iron to the divalent state
(3) titration of iron(II) with a standard oxidant.
Iron ores are often completely decomposed in hot concentrated HCl. Because
iron(III) tends to form stable chloride complexes, HCl is a much more efficient
solvent than either sulfuric or nitric acid.
Many iron ores contain silicates that may not be entirely decomposed by
treatment with HCl. Incomplete decomposition is indicated by a dark residue that
remains after prolonged treatment with the acid. A white residue of hydrated silica,
which does not interfere in any way, is indicative of complete decomposition.
Part or all of the iron will exist in the trivalent state in the dissolved sample;
pre-reduction must therefore precede titration with the oxidant. Many reductants
have been used in determinations of iron, including sulfites, SO2, H2S, SnCl2, and a
variety of amalgamated and free metals. Reduction by zinc in the well-known Jones
reductor is avoided because of the inconvenience and expense of maintaining the
columns and because of the mercury usually included. The popular method of
67
reduction with excess SnCl2 and elimination of the excess with HgCl2 is also
environmentally objectionable. Zinc is the reductant of choice because of the
conceptual and practical simplicity of its use.
In this experiment, the reduction of iron(III) to iron(II) is carried out in three
stages in a small volume of solution strongly acidified with HCl. A high
concentration of HCl accentuates the yellow color of the iron(III) to make the
progress of reduction easily observed. Since the HCl is exhausted in each stage by
reaction with zinc, it is replenished in each following stage. The reintensified color
provides a visual indicator until reduction is complete. Unreacted zinc is
subsequently eliminated by reaction with H2SO4.
Potassium permanganate, KMnO4, is widely used as an oxidizing agent in
volumetric analysis.
MnO4- + 8H+ + 5e-  Mn2+ + 4H2O
E0 = 1.51 V
In 1 M HClO4, the formal potential is just 1.70 V.
In 1 M HNO3, it is 1.61 V.
The half-reaction shown above for the permanganate ion occurs only in
solutions that are 0.1 M or greater in strong acid. In less acidic media, the product(s)
may be Mn (III), Mn (IV), or Mn (VI) depending on conditions.
The oxidation of iron(II) with permanganate is based on the reaction:
5 Fe2+ + MnO4- + 8 H+ → 5 Fe3+ + Mn2+ + 4 H2O
Often, this reaction is performed in the presence of moderate concentrations of HCl.
But the permanganate ion cannot be used with hydrochloric acid solutions for the
slow oxidation of chloride ion, this can be eliminated through removal of the
68
hydrochloric acid by evaporation with sulfuric acid or by introduction of
Zimmermann-Reinhardt reagent, a mixture of manganese(II), concentrated
sulfuric and phosphoric acids.
Because permanganate ion is violet in color and manganese(II) ion is nearly
colorless, no additional indicator is needed for this titration, when one drop excess
of potassium permanganate has been added to the sample, the endpoint can be taken
as the first pale red/pink color. Alternatively, reactions can be monitored with Pt and
calomel electrodes.
Reagents
1. Potassium permanganate
2. Concentrated hydrochloric acid
3. Concentrated sulfuric acid
4. Zimmermann-Reinhardt reagent
Procedure
A. Preparation of 0.02 M potassium permanganate solution
Measure about 60 mL of standardized potassium permanganate solution.
B. Determination of iron in an ore
(1) Dissolution of the sample:
1. Dry the ore sample for at least 3 hours at 105 to 110 oC, and cool in a desiccator.
Weigh accurately a 0.30 g portion of ore sample into a 250-mL Erlenmeyer flask.
69
2. Add 8 mL of concentrated HCl. Heat nearly to boiling in hood until the ore is
dissolved, about 5~15 minutes. Some white residue of silica may remain and can
be ignored. If evaporation reduces the volume significantly, add additional HCl;
that is, iron oxide is reprecipitated. Boil the solution gently for a few minutes to
expel as much HCl(g) as possible (Note 1).
(2) Reduction of the iron to the divalent state:
1. Cool briefly, and add slowly ~1.5 g of zinc metal in hood. Heat or cool the flask
as necessary while swirling it gently, to maintain a vigorous but not violent
reaction, until the reduction is complete (Note 2).
2. When the solution appears to be completely reduced, add another 0.75 g of zinc
and heat nearly to boiling in hood for ~ 5 minutes.
3. Add 15 mL of distilled water and 3 mL of concentrated H2SO4. Mix well. Add
12 mL of Zimmermann-Reinhardt reagent and stir the solution until the
remaining zinc is dissolved.
4. Transfer into a 100-mL volumetric flask. and wash the residue in the beaker
with distilled H2O. Dilute the solution to the mark, and mix thoroughly.
(3) Titration of iron(II) with a standard oxidant:
1. For the rough titration, pipet 20.0 mL aliquots of the sample solution into a
250-mL Erlenmeyer flask. Titrate immediately with standard ~0.02 M KMnO4
solution until the pink color persists for 15 s.
2. Then comes the careful titration, pipet 40.0 mL of the sample solution into a
100-mL beaker containing a magnetic stirring bar. Position the Pt combination
electrode in the solution so that the stirring bar will not strike the electrode. Set
70
the pH meter to measure potential rather than pH.
3. Titrate immediately with standard ~0.02 M KMnO4 solution. The equivalence
volume will be two times greater than it was in Step 1. Firstly add 2.00 mL of
MnO4- solution from the buret, after mix well, record the volume, potential and
color. When you are within 1 mL of the indicator end point, add titrant in 0.1 mL
increments. Continue with 0.1 mL increments until you are 1.0 mL past the
indicator end point. The equivalence point has the most rapid change in potential.
Add three more 2.0 mL aliquots of titrant and record the potential after each.
4. Make the following plots of the data:
(1) potential versus titrant volume
(2)
ΔE/ΔV
vs. average titrant volume.
From the equivalence point of the titration curve, determine the wt% of iron in the
ore sample.
Notes
1. The HCl concentration cannot drop much below 6 M, but that does reduce
somewhat the amount of zinc required for reduction.
2. The blue-green hue of iron(II) will be apparent in a solution that is completely
reduced; a solution that appears green with a trace of yellow will require an
additional portion of zinc.
References
l. Skoog, West, Holler and Crouch, Fundamentals of Analytical Chemistry, 8th ed.,
Chap. 20, pp.574-575
71
2. D.C. Harris, Quantitative Chemical Analysis, 7th ed., Chap. 16
3. Skoog, West, Holler, Crouch, and Chen, Introduction to Analytical Chemistry,
2011, Chap. 11, 12, 13A, & 13B
72
Data
A. Preparation of 0.02 M potassium permanganate solution
The Molarity of KMnO4: _____________________ M
B. Determination of iron in an ore
(1) Dissolution of the sample:
Unknown No.
Weight of sample (g)
g
(3) Titration of iron(II) with a standard oxidant:
1. Rough titration
Initial buret reading
mL
Final buret reading
mL
Volume of KMnO4
mL
73
3. Careful titration
Estimated volume of end point for 40 mL unknown sol’n
V(mL)
E (mV)
color
0
2.00
4.00
V(mL)
E (mV)
color
V(mL)
E (mV)
color
V(mL)
E (mV)
color
V(mL)
E (mV)
color
V(mL)
E (mV)
color
V(mL)
E (mV)
color
74
mL
Data analysis
1. Construct a graph of potential(y-axis) versus titrant volume(x-axis). Make on
your graph where the indicator color change was observed.
75
2. Compute the first derivative (the slope,
±
ΔE/ΔVmL)
0.5 ml of the indicator end volume. Construct graph of
V(mL)
E (mV)
△E (mV)
△VmL
△E/△VmL
Vavg, mL
( Graph: △E/△VmL v.s Vavg,mL )
76
for each data point within
ΔE/ΔVmL
v.s Vavg,mL .
3. From the equivalence point of the titration curve, determine the wt% of iron in
the ore sample.
Accurate value of %Fe
%
Relative error
%
Discussion
77
78
Experiment 9
Spectrophotometric Determination of Manganese in Steel
Introduction
Manganese may be determined by conversion to the permanganate ion
which has a large molar absorptivity at 525 nm. In this experiment, the periodate ion
is used as the oxidant and, if it is kept in excess, the permanganate solution will be
stable.
2Mn2＋＋ 5IO4－ ＋ 3H2O ⇆ 2MnO4－ ＋ 5IO3－ ＋ 6H+
(colorless) (colorless)
(violet)
(colorless)
In determinations such as this, it is important to ascertain that no other ions
are present which will absorb at the chosen wavelength. In this analysis, the only
other ion which can cause complications is the dichromate ion which is formed by
oxidation of chromium, a common ingredient of steel. We have chosen the steel
sample so that this will not be a problem.
Since most of the sample is composed of iron and since ferric iodate is only
slightly soluble, the ferric ion must be kept in solution by the addition of masking
agent. Phosphoric acid forms sufficiently strong soluble colorless complexes with
ferric ion to accomplish this; the identities of these complexes are not, however,
well known.
Instrument
Bausch & Lomb Spectronic 20 Spectrophotometer
79
Reagents
1. Concentrated nitric acid
2. 4 M nitric acid
3. Ammonium persulfate
4. Phosphoric acid
5. Potassium periodate
Procedure
A. Preparation of pure manganese solution
Dissolve 0.010 g (±10% but accurately weighed) of pure manganese metal
in 7 mL of concentrated nitric acid. The solution was diluted to exactly 10 mL in a
volumetric flask.
B. Dissolution of steel and oxidation of manganese
Weigh duplicate 0.010 g of steel sample and dissolve in 8 mL of 4 M HNO3.
Boil for 5 minutes. Cautiously add about 0.5 g of ammonium persulfate, (NH4)2S2O8,
and boil gently for 10 ~ 15 minutes. Do not allow the solution to boil dry. If the
solution is pink or contains a brown oxide of manganese, add approximately 0.05 g
of sodium hydrogen sulfite, NaHSO3, and heat for another minute. Cool and dilute
to exactly 10 mL in a volumetric flask.
80
C. Formation of permanganate and measurement of its absorbance
(1) pure manganese solution
1. Pipette 1.00 mL aliquots of the pure manganese solution (from A) into small
beaker, add 2 mL of conc. phosphoric acid which serves to complex ferric ion.
Add 0.2 g of potassium periodate and boil for 2 minutes. Cool and dilute to
exactly 10 mL in a volumetric flask.
2. From above pure manganese solution, pipette 0.50 mL, 1.00 mL, and 1.50 mL
aliquots into three 10-mL volumetric flasks, receptively. Dilute to the mark with
distilled water and mix well. Determine the absorbance of each sample against
the blank at 525 nm to obtain a Beer's law plot.
(2) steel sample solution
1. Pipette 1.00 mL aliquots of the unknown solution (from B) into a small beaker
and add 2 mL conc. phosphoric acid which serves to complex ferric ion . Add 0.2
g of potassium periodate and boil for 2 minutes. Cool and dilute to exactly 10
mL in volumetric flask.
2. From above unknown solution, pipette 2.00 mL aliquots into 10-mL volumetric
flask. Dilute to the mark with distilled water and mix well. Measure the
absorbance against the blank at 525 nm and determine the % by weight of Mn in
the original steel sample.
References
1. Skoog, West, Holler and Crouch, Fundamentals of Analytical Chemistry, 8th ed.,
Chap. 24, pp.718-724; Chap. 25, pp.771-772
81
2. D.C. Harris, Quantitative Chemical Analysis, 7th ed., Chap. 18, pp.378-387
3. Skoog, West, Holler, Crouch, and Chen, Introduction to Analytical Chemistry,
2011, Chap. 14 & pp.473-487
82
Data
A. Preparation of pure manganese solution
Weight of pure Manganese: ____________ g
B. Dissolution of steel and oxidation of manganese
Unknown No.
Weight of sample
g
C. Formation of permanganate and measurement of its absorbance
(1) pure manganese solution:
Concentration of pure Mn (g/ mL)
Absorbance
Obtain a Beer's law plot (Calibration curve) and find the linear regression equation
between absorbance and concentration.
83
(2) steel sample solution:
Absorbance of diluted unknown solution
Linear regression equation
Concentration of diluted unknown solution
g/mL
Concentration of original unknown solution
g/mL
% by weight of Mn in sample
%
Accurate value of %Mn
%
Relative error
%
Calculation
Discussion
84
Experiment 10
Simultaneous Determination of Binary Mixture by
Spectrophotometric Method
Introduction
The absorbance of a solution at any wavelength is the sum of absorbances of
all the species in the solution.
Atotal = Xb[X] + Yb[Y] + zb[Z] + ……
where is the molar absorptivity of each species at the wavelength in question and
b is the cell pathlength. If we know the spectra of the pure components, we can
mathematically disassemble the spectrum of a mixture into those of its components.
To analyze a mixture of two compounds, it is necessary to measure absorbance at
two wavelengths and to know at each wavelength for each compound.
For a mixture of compounds X and Y, if spectra of the individual components
are moderately resolved from each other, as at wavelengths ’ and ” in the figure
above, we can solve two simultaneous equations to find the concentrations in the
mixture. The absorbance of the mixture at any wavelength is the sum of
85
absorbances of each component at the wavelength. For wavelengths ’ and ”,
A’ = X’b[X] + Y’b[Y]
A” = X”b[X] + Y”b[Y]
where the  values apply to each species at each wavelength. The molar
absorptivities of X and Y at wavelengths ’ and ” must be measured in separate
experiments. 
Instrument
Bausch & Lomb Spectronic 20 Spectrophotometer
Reagents
1. 5.00 x 10-5 M Tartrazine (Yellow # 4)
2. 5.00 x 10-5 M Allura Red (Red # 40)
Procedure
A. Finding the wavelengths of maximum absorbance
Determine the spectrum of the standard solutions (5.00 x 10-5 M Tartrazine and
5.00 x 10-5 M Allura Red) from 400 to 510 nm at every 10 nm interval. (Adjust
0Abs when change wavelengths.) Analytical wavelengths are 425 and 505 nm. (If
your analytical wavelengths are other than 425 and 505 nm, use your own ones for
later operations.)
86
B. Preparing diluted solutions of Tartrazine and Allura Red
Pipette 2.00, 4.00, 6.00 and 8.00 mL of 5.00 x 10-5 M Tartrazine and 5.00 x
10-5 M Allura Red into different 10-ml volumetric flasks, and dilute to the mark
with distilled water, respectively. Mix well.
C. Beer's law plots and Absorption of unknown mixture
At the two analytical wavelengths where maximum absorbance occurred,
measure the absorbance of the four Tartrazine, the four Allura Red solutions, and
unknown solution.
Results
1. From procedure A, plot the measured absorbance, A, versus wavelength, , for
both standard Red and Tartrazine solutions. From the graphs plotted, determine
the wavelengths of the maximum absorbance.
2. From procedure B and C, (a) calculate the concentration of each of Red and
Tartrazine standard solutions. (b) plot the measured absorbance versus
concentration for both standard solutions at the characteristic wavelengths. (c)
from the graphs plotted, find the molar absorptivities () of Red and Tartrazine.
3. Calculate the concentrations of Red and Tartrazine in the unknown solution.
References
1. Skoog, West, Holler and Crouch, Fundamentals of Analytical Chemistry, 8th ed.,
Chap. 24, pp. 718-724; Chap. 25, pp. 771-772
87
2. D.C. Harris, Quantitative Chemical Analysis, 7th ed., Ch.19, pp.402-407
3. Skoog, West, Holler, Crouch, and Chen, Introduction to Analytical Chemistry,
2011, Chap. 14 & pp.473-487
88
Data
A. Finding the wavelengths of maximum absorbance
λ(nm)
400
410
420
425
430
440
450
460
470
480
490
500
505
510
Abs. of
Tartrazine
Abs. of
Allura Red
λ(nm)
Abs. of
Tartrazine
Abs. of
Allura Red
Plot the measured absorbance, A, versus wavelength, , for both standard Red and
Tartrazine solutions. From the graphs plotted, determine the wavelengths of the
maximum absorbance.
89
90
C. Beer's law plots and Absorption of unknown mixture
Abs. at 425 nm
solutions
2.00 mL
4.00 mL
6.00 mL
8.00 mL
Concentration(M)
Tartrazine
Allura Red
Plot the measured absorbance versus concentration for both standard solutions at the
characteristic wavelengths (Calibration curve) and find the linear regression
equations.
91
Abs. at 505 nm
solutions
2.00 mL
4.00 mL
6.00 mL
8.00 mL
Concentration(M)
Tartrazine
Allura Red
Plot the measured absorbance versus concentration for both standard solutions at the
characteristic wavelengths (Calibration curve) and find the linear regression
equations.
92
Linear regression
Tartrazine
equation
at 425 nm
Allura Red
Linear regression
Tartrazine
equation
at 505 nm
Allura Red
Unknown No.
Abs. of unknown solution at
425 nm
Abs. of unknown solution at 505 nm
Calculate the concentrations(M) of Red and Tartrazine in the unknown
solution.
93
Experimental result of Tartrazine
M
Accurate value of Tartrazine
M
Relative error
%
Experimental result of Allura Red
M
Accurate value of Allura Red
M
Relative error
%
Discussion
94