EDTA COMPLEXOMETRIC TITRATIONS

EDTA COMPLEXOMETRIC TITRATIONS.

EDTA COMPLEXOMETRIC TITRATIONS
Instructions
2.Experimental – approximately half a typed page. You are to include experimental steps
you have taken in conducting the experiment, such that the reader can repeat the
experiment by following instructions given in this section. However, there is no need to
include lengthy description on preparation of solutions provided to you.
3. Results and Discussion (you are to present results in the Results section, followed by a
discussion and/or comments on the results in the Discussion section).
(a) Tabulate all titration results in a table with appropriate units shown in each column.
(b) Using your own results, calculate the mean, standard deviation and relative standard
deviation of the titres in each of the three assays above. Based on an appropriate
stoichiometric ratio, you are able to work out the number of moles of Al3+, Ca2+ and
Mg2+ present in the titration experiments. You should subtract moles of Al3+, Ca2+ and
Mg2+ present in the blank solutions from the total to obtain the true number of moles
present in the sample. Next, calculate the number of moles of Al3+, Ca2+ and Mg2+
present in the sample weighed out in the beginning of the experiment. A sketch shown
for Ca2+ assay below may help you analyse steps required in your calculations.
Next, determine the percentage weight (i.e. concentration) of Al3+ and that of individual Ca2+
and Mg2+ in the antacid sample. You can ignore calculating the standard deviations of the
percentage weights.
(c) Use the mean and standard deviations from part (b) to calculate the 95% confidence
intervals for the mean titres Al3+, Ca2+ and Mg2+ of future antacid batches.
(d) For each assay above, how many titrations should you perform if you wish to be 95%
confident to have an estimated bias within ±0.05?
(e) Based on the stipulated values, estimate the percentage weight of Al3+ and that of
individual Ca2+ and Mg2+ (note we are to calculate the percentage weight of ions
here, not the compounds) in the antacid sample.
Compare and comment on the results estimated in (b) and (e). [Note: when you are asked
to make a comparison, do not simply report that one result is higher/lower than the other,
which is obvious when we examine two numbers. Instead, you should comment on what
could have happened during the experiment, giving rise to such a difference.] You can
use the percentage difference between your experimental value and the stipulated value
for each ion as a guide in making your comparison.
Comment on all relevant errors encountered in this experiment. If necessary, classify
them under the headings “Determinate Errors”, “Indeterminate Errors” and “Gross
Errors”. Caution: Be careful of the distinctions among “errors”, “mistakes” and
“carelessness”.
4. Conclusions – no more than half a typed page.
There are many possible ligands capable of forming complexes with metal ions. It should be
kept in mind that water itself will act as a ligand by donating electron-pair density on the
oxygen atom to a metal ion in solution. Thus, all complexation reactions in aqueous solution
are really a competition between H2O and the ligand, L, for the metal ion, as shown in Equation
(1).
!
M(H2O)6
n+
+ 6Ll”
!
ML6n”6l + 6H2O (1)
For example, the formation of hexacyanoferrate(II) ion can be written as
!
Fe(H2O)6
2+
+ 6CN”
!
Fe(CN)6
4″
+ 6H2O (2)
Ligands, such as H2O, CN-, NH3, and halide ions, which donate one pair of electrons, are called
monodentate ligands. There are a few important analytical applications using these inorganic
monodentate ligands. However, the great rise in popularity of complexometric titrations has
been due to the development of organic complexing agents that donate more than one electron
pair. For example, ethylenediamine, with the formula NH2-CH2-CH2-NH2, donates two
Experiment 2 EDTA Complexometric Titrations 2
electron pairs – one on each nitrogen atom, to form a chelate, and is called a bidentate ligand.
Ligands donating three, four, five or six electron pairs are termed tri-, tetra-, penta-, and
hexadentate.
For analytical purposes, by far the most widely used complexing agent for complexometric
titrations is ethylenediaminetetraacetic acid, also known as ethylene-dinitrilotetraacetate. It
has the structure
and is hexadentate because (i) all four hydrogen atoms on the carboxylic acid groups can be
replaced by the metal ion, and (ii) the two nitrogen atoms donate electron pairs. The ionised
anion of this complexing agent is often given the abbreviation EDTA and the structure of an
EDTA complex is shown in Figure 1. EDTA belongs to a group of exceptionally strong
ligands called aminopolycarboxylic acids.
Figure 1. Structure of a metal-EDTA complex. EDTA forms hexadentate chelates.
Electron pair donation is from oxygen and nitrogen to the metal centre.
Solutions of EDTA are particularly valuable as titrants because the reagent combines with
metal ions in a 1:1 ratio regardless of the charge on the cation. For example, formation of the
silver and aluminium complexes is described by the equations
Ag+ + (EDTA)4- Ag(EDTA)3-
Al3+ + (EDTA)4- Al(EDTA)-
EDTA is a remarkable reagent not only because it forms chelates with many cations, but also
because most of these chelates are sufficiently stable to form the basis for a titrimetric
method. This great stability undoubtedly results from the several complexing sites within the
molecule that give rise to a cage-like structure in which the cation is effectively surrounded
and isolated from solvent molecules, as shown in Figure 1.
It may appear that EDTA titrations would not be very useful because essentially all metal ions
form complexes. In other words, the method lacks specificity. However, the use of pH, as
N CH2 CH2 N
CH2
CH2
HOOC
HOOC CH2
CH2 COOH
COOH
O N
O N
M
2+
CH2
C
O
CH2
CH2
O
C C
O
C
O
CH2
C
O
CH2
O
H2
2-
Experiment 2 EDTA Complexometric Titrations 3
shown in Figure 2, offers a means of titrating one metal ion (of higher conditional formation
constant, KMY) in the presence of another (of lower conditional formation constant).
EDTA TITRATION TECHNIQUES
As so many elements can be analysed by titration with EDTA, there is an extensive literature
dealing with many variations of the basic procedure.
Direct Titration
In a direct titration, analyte is titrated
with standard EDTA. The analyte is
buffered to an appropriate pH at which
the conditional formation constant for the
metal – EDTA complex is large enough to
produce a sharp end-point. As most
metal ion indicators are also acid-base
indicators, they have different colours at
different values of pH (see Figure 2). An
appropriate pH must be one at which the
free indicator has a distinctly different
colour from the metal-indicator complex.
Figure 2. Minimum pH needed for
satisfactory titration of various cations with
EDTA.
Back Titration
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
when no suitable indicator is available, if the analyte precipitates in the absence of EDTA, or
if it reacts too slowly with EDTA under titration conditions. The metal ion used in the back
titration must not displace the analyte metal ion from its EDTA complex.
Displacement Titrations
For metal ions that do not have a satisfactory indicator, a displacement titration may be feasible.
In this procedure the analyte usually is treated with excess Mg(EDTA)2- to displace Mg2+,
which is later titrated with standard EDTA.
(3)
Hg2+ is determined in this manner. The formation constant of Hg(EDTA)2- must be greater
than the formation constant for Mg(EDTA)2-, or else Reaction (3) will not occur spontaneously.
There is no suitable indicator for Ag+. However, Ag+ will displace Ni2+ from the
tetracyanonickelate ion:
(4)
The liberated Ni2+ can then be titrated with EDTA to find out how much Ag+ was added.
Experiment 2 EDTA Complexometric Titrations 4
Indirect Titrations
Anions that form precipitates with certain metal ions may be analysed with EDTA by indirect
titration. For example, sulfate can be analysed by precipitation with excess Ba2+ at pH 1. The
BaSO4 precipitate is filtered and washed. Boiling the precipitate with excess EDTA at pH 10
brings the Ba2+ back into solution as Ba(EDTA)2-. The excess EDTA is back-titrated with
Mg2+.
Alternatively, an anion may be precipitated with excess metal ion. The precipitate is filtered and
washed, and the excess metal ion in the filtrate is titrated with EDTA. Anions such as
!
CO3 2″,
!
CrO4
2″, S2-, and
!
SO4 2″ can be determined by indirect titration with EDTA.
Masking
A masking agent is a reagent that protects some component of the analyte from reaction with
EDTA. For example, Al3+ reacts with F- to form the very stable complex
!
AlF6
3″. The Mg2+ in
a mixture of Mg2+ and Al3+ can be titrated by first masking the Al3+ with F-, leaving only the
Mg2+ to react with EDTA.
Cyanide is a common masking agent that forms complexes with Cd2+, Zn2+, Hg2+, Co2+, Cu+,
Ag+, Ni2+, Pd2+, Pt2+, Fe2+, and Fe3+, but not with Mg2+, Ca2+, Mn2+, or Pb2+. If cyanide is
first added to a solution containing Cd2+ and Pb2+, only the Pb2+ is then able to react with
EDTA. Fluoride can mask Al3+, Fe3+, Ti4+ and Be2+. Triethanolamine masks Al3+, Fe3+ and
Mn2+, and 2,3-dimercaptopropanol masks Bi3+, Cd2+, Cu2+, Hg2+, and Pb2+.
Demasking refers to the release of a metal ion from a masking agent. Cyanide complexes can
be demasked by treatment with formaldehyde:
M CN ( )m
n !m + mH2CO + mH+ ”
+ Mn+ (5)
Thiourea reduces Cu2+ to Cu+ and masks Cu+. Cu2+ can be liberated from the Cu(I)-thiourea
complex by demasking with H2O2. The selectivity afforded by masking, demasking and pH
control allows individual components of complex mixtures of metal ions to be analysed by
EDTA titration.
In this experiment, we will determine the concentration of Al3+, Ca2+ and Mg2+ in “Quick-
Eze” samples by EDTA complexometric titration. Ca2+ is then analysed separately by
precipitating Mg(OH)2.
Reagents (to be prepared by laboratory technicians)
1. Buffer (pH 4.5) – add 14 mL of glacial acetic acid to 19.3 g of ammonium acetate and
dilute to 250 mL.
mH2C
CN
OH
Experiment 2 EDTA Complexometric Titrations 5
2. Buffer (pH 10) ammonia-ammonium chloride solution – 54 g NH4Cl dissolved in
approximately 200 mL of deionised water. To this solution, 350 mL of ammonia solution
(not less than 25%) is added and then made up to 1 L with deionised water.
3. 0.01 mol L-1 Zinc Sulfate standard – accurately weigh approximately 0.72 g of
ZnSO4•7H2O and dissolve in a 250 mL volumetric flask with water. Dilute to the mark.
4. Dithizone indicator – (to be prepared freshly) dissolve 0.3 g of dithizone
(diphenylthiocarbazone) in 100 mL of absolute ethanol.
5. 3 mol L-1 HCl.
6. Hydroxynapthol blue indicator (Calcon or Solochrome Dark Blue).
7. Antacid – tablets or suspension.
8. Eriochrome black T indicator. Dissolve 0.2 g of indicator in 15 mL of triethanolamine,
and 5 mL of absolute ethanol.
Preparation of Samples
A sample of antacid powder (“Hardy’s indigestion powder”) is supplied containing aluminium
hydroxide, calcium carbonate and magnesium carbonate. Students should note the stipulated
percentage weights and be aware that quality control regulations for pharmaceuticals of this
nature allow ±10% deviation from the stipulated concentrations. Weigh approximately 0.65 g
of the sample accurately (THINK! Which type of balance should you use?) in a 250 mL
conical flask. Record this accurate mass. Add 5 mL of water and then 10 mL of 3 mol L-1 HCl
(Do we need accurate volumes here?) Heat the mixture (microwave for 30 seconds) to boiling
and shake until dissolved. Add about 100 mL of water (does this mean we need to measure out
100 mL accurately?) and then quantitatively transfer the solution to a 250.0 mL volumetric
flask and dilute to volume.
PROCEDURE
Assay for Al3+
Pipette 25.00 mL of the prepared sample into a 250 mL conical flask and add 20 mL of water.
Then, with swirling add, by pipette, 25.00 mL of 0.01 mol L-1 EDTA (you need to copy down
the EXACT concentration (it therefore should not have just one or two significant figures!) of
EDTA provided). Continue stirring, and add 20 mL of pH 4.5 buffer. (NOTE: The order of
these additions is important). Heat the solution in a microwave oven for 2 minutes. Cool,
add 50 mL ethanol and then 2 mL of dithizone indicator (be careful not to use any undissolved
dithiozone).
1. Titrate with 0.01 mol L-1 zinc sulfate standard (copy down the EXACT concentration)
until the colour changes from green-violet to rose-pink. You should estimate the titre to 2
decimal places.
2. Perform a blank, using 25.00 mL of deionised water instead of the prepared sample.
3. Repeat the titration to obtain three concordant values.
Experiment 2 EDTA Complexometric Titrations 6
Assay for Mg2+ + Ca2+
Pipette 10.00 mL of the prepared sample into a 200-250 mL conical flask. Add 34 mL of an
aqueous triethanolamine solution and mix. Add 20 mL of pH 10 buffer and 2 drops of
Eriochrome black T indicator. Cool the solution to approximately 3-4°C in an ice bath (do not
place a thermometer in the solution or you would have lost a small volume of the solution after
removing the thermometer) and then titrate with standard 0.01 mol L-1 EDTA to a pure blue
colour (not violet). Perform a blank. Repeat the sample titration to obtain three concordant
values.
Assay for Ca2+
Pipette 10 mL of the prepared sample into a 200-250 mL conical flask. Add 70 mL of water
and then 1 mL of 50% NaOH solution and 10 mg of hydroxynapthol blue. Titrate with 0.01
mol L-1 EDTA until a distinct blue colour is achieved. Perform a blank. Repeat the titration
to obtain three concordant values.
Record the results obtained in the three assays above and the respective masses used, on an
Excel sheet so that results from the entire class can be pooled. Each of you is then to analyse
the results independently.
LABORATORY REPORT
Please check the list of dos and don’ts dealt with in Week 1
before preparing the laboratory report!
1. Introduction – In CBMS208/CBMS608, only a one-page Introduction is required in each
laboratory report. You need to carefully examine the model Introduction to learn how to
construct a concise and coherent account.
2. Experimental – approximately half a typed page. You are to include experimental steps
you have taken in conducting the experiment, such that the reader can repeat the
experiment by following instructions given in this section. However, there is no need to
include lengthy description on preparation of solutions provided to you.
3. Results and Discussion (you are to present results in the Results section, followed by a
discussion and/or comments on the results in the Discussion section).
(a) Tabulate all titration results in a table with appropriate units shown in each column.
(b) Using your own results, calculate the mean, standard deviation and relative standard
deviation of the titres in each of the three assays above. Based on an appropriate
stoichiometric ratio, you are able to work out the number of moles of Al3+, Ca2+ and
Mg2+ present in the titration experiments. You should subtract moles of Al3+, Ca2+ and
Experiment 2 EDTA Complexometric Titrations 7
Mg2+ present in the blank solutions from the total to obtain the true number of moles
present in the sample. Next, calculate the number of moles of Al3+, Ca2+ and Mg2+
present in the sample weighed out in the beginning of the experiment. A sketch shown
for Ca2+ assay below may help you analyse steps required in your calculations.
Next, determine the percentage weight (i.e. concentration) of Al3+ and that of individual Ca2+
and Mg2+ in the antacid sample. You can ignore calculating the standard deviations of the
percentage weights.
(c) Use the mean and standard deviations from part (b) to calculate the 95% confidence
intervals for the mean titres Al3+, Ca2+ and Mg2+ of future antacid batches.
(d) For each assay above, how many titrations should you perform if you wish to be 95%
confident to have an estimated bias within ±0.05?
(e) Based on the stipulated values, estimate the percentage weight of Al3+ and that of
individual Ca2+ and Mg2+ (note we are to calculate the percentage weight of ions
here, not the compounds) in the antacid sample.
Compare and comment on the results estimated in (b) and (e). [Note: when you are asked
to make a comparison, do not simply report that one result is higher/lower than the other,
which is obvious when we examine two numbers. Instead, you should comment on what
could have happened during the experiment, giving rise to such a difference.] You can
use the percentage difference between your experimental value and the stipulated value
for each ion as a guide in making your comparison.
Comment on all relevant errors encountered in this experiment. If necessary, classify
them under the headings “Determinate Errors”, “Indeterminate Errors” and “Gross
Errors”. Caution: Be careful of the distinctions among “errors”, “mistakes” and
“carelessness”.
250.0 mL
Antacid
Sample
Solution
10.0 mL
Ca2+
EDTA
Experiment 2 EDTA Complexometric Titrations 8
4. Conclusions – no more than half a typed page.
ADDITIONAL EXERCISES
1. Many elements can be analysed by titration with EDTA. Four important EDTA titration
techniques have been described under Introduction in this experiment. Which technique
has been employed in the determination of Al3+ in this experiment? Give three
circumstances in which such a technique might be necessary.
2. In many direct titrations, an auxiliary complexing agent is used. Give an account on what
this is and how it works in EDTA titration.
Questions 3 and 4 below refer to the results obtained by the entire group of students during
your laboratory session. You would have downloaded from the CBMS208 webpage an Excel
sheet consisting all the results from your group. If you have processed the results using
3. Select any one student’s results from the Excel sheet. Test if the mean Al3+
concentration in percentage weight (note: not the mean titre) from this set of results
differ significantly from that you estimated from your own results. [Why can’t we
compare the mean titre instead?]
4. Using Excel, determine the corresponding percentage weight of Ca2+ from each set of
results recorded. Treat these results as obtained by different laboratories. Using Excel
again, determine the within-laboratory sum-of-squares for each laboratory as well as the
between-laboratory sums-of-squares. [You are required to submit the Excel sheet
together with your report.] Based on these values, set up your own analysis-of-variance
table in the laboratory report. Then:
(a) Calculate an estimate of within laboratories standard deviation (sw).
(b) Calculate an estimate of between laboratories standard deviation (sb).
(c) Perform an appropriate hypothesis testing to determine if the between-laboratories
variance is significant.
(d) Perform an appropriate hypothesis testing to determine if the within-laboratories
variance is not significantly different.
(e) Perform an appropriate hypothesis testing to check that the mean determination for
each laboratory comes from the same Normal distribution.
(f) Estimate the least significant difference and then determine if any three percentage
weights of Ca2+ (one of them is your own result) differ from each other at the 95%
confidence level.
(g) Calculate an estimate of repeatability of your results.
(h) Calculate an estimate of reproducibility of the Ca2+ titration technique.
(i) Calculate a 95% confidence interval for the true mean Ca2+ concentration and
hence estimate the maximum bias for the titration method. [Note: maximum bias
= maximum variation – true (or stipulated) value].
(j) Calculate a 95% confidence interval for your laboratory and hence estimate its
maximum bias.
Experiment 2 EDTA Complexometric Titrations 9
5. Apart from EDTA, there are many other well-known ligands which readily complex with
metal ions. (i) Name and give the structure of the ligand that forms a brick red coloured
complex with Ni2+ and Co2+ (give the structure of this ligand). (ii) Name and give the
structure of the ligand that forms a brick red coloured complex with Fe3+. [You will need
to refer to some Inorganic Chemistry textbooks in the Library to search for the
appropriate information. Please write down the source of information used in answering
this question.]
REFERENCES
1. CBMS208/CBMS608 lecture notes on Statistics for Analytical Chemists (2017).
3. D.C.Harris, Quantitative Chemical Analysis, 9th edition, Chapter 13 (2016).
4. D.A.Skoog, D.M.West, F.J.Holler, S.R.Crouch Fundamentals of Analytical Chemistry, 9th
Edition, Chapter 17 (2014).
5. A.I.Vogel, A Textbook of Quantitative Inorganic Analysis. (1962).
Experiment 2 EDTA Complexometric Titrations 10
Marking scheme for Experiment 2
Section Assessment Weighting
Pre-lab Work Introduction 10
Flow diagram and MSDS 10
Lab Notebook Quality of information; ease in understanding information 10
Introduction: Content:
General Introduction 5
Specific Introduction
Aim(s) of Experiment
Layout:
Coherent 2
Correct length
Experimental: Presented in Correct tense 5
Correct use of scientific words to describe actions
Sufficient information given without giving perfunctory information
Results: Presentation:
Results neatly presented/tabulated 3
Correct number of significant figures in both data and uncertainties
Values presented with appropriate units
Tables:
Tables have self explanatory legends and are numbered 3
Column headings Identify Data
Units of measurement indicated in column headings
Values in a column presented with the same precision
Statistics:
Correctly presented hypothesis testing 3
Data statistically processed in an appropriate manner 5
Quality:
Quality of experimentally obtained data 5
Data Processing:
Requisite processing of experimentally obtained data 15
Discussion: Adequate discussion of Results obtained 15
Comparison of results to literature/stipulated values with
adequate discussion of any differences
3
Major sources of error are identified in both procedure
and data
2
Conclusion: Summary of the aims of the experiment and whether aims
have been achieved
3
Suggestions for addressing main sources of error 1
Exercises: Adequate attempt at Final Exercises 10
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