The lab report below was submitted as part of the coursework for
CM2142 Analytical Chemistry. Please do not plagiarise from it as
plagiarism might land you into trouble with your university. Do note
that my report is well-circulated online and many of my juniors have
received soft copies of it. Hence, please exercise prudence while
referring to it and, if necessary, cite this webpage.
1.
Aim
To determine the concentration of cupric ions in aqueous solutions
using ion selective electrodes by calibration curve and titration with ETDA.
2. Results
and calculations
By
calibration curve
Preparation
of 1000pmm standard solution
Concentration of Cu2+
solution used= 1000ppm = 1 g L-1Mass of Cu2+ in 50mL of
solution = = 0.05g
No. of moles of Cu2+
in 50mL of solution = = 7.868 x 10-4 mol
Concentration x
volume = no. of moles 0.1 x v = 7.868 x 10-4
Volume of 0.1 M Cu2+
used = = 0.007868 L = 7.87mL 7.90mL
Table 1: electrode potential at various concentrations of Cu2+
Volume
used to prepare standard
|
Concentration
of Cu2+ /ppm
|
Log [Cu2+]
|
Electrode
Potential /mV
|
7.90 mL of 0.1 M Cu2+
|
1000
|
3
|
272
|
5.00 mL of 1000ppm
|
100
|
2
|
241
|
5.00 mL of 100ppm
|
10
|
1
|
211
|
5.00 mL of 10ppm
|
1
|
0
|
181
|
5.00 mL of 1ppm
|
0.1
|
-1
|
151
|
-
|
Unknown
|
-
|
272
|
Graph 1: Graph of electrode potential against log [Cu2+]
From the graph, y= 30.2 x +181, where y
is the electrode potential (mV) and x is the log [Cu2+].
Since the unknown Cu2+ solution gave an
electrode potential reading of 244 mV,
log[Cu2+] in unknown solution =
(244-181) ÷ 30.2 = 2.086
Concentration of unknown Cu2+ solution =
10 2.086 ≈ 121.89 ppm ≈ 122mg/L =0.122 g /L
Molarity of unknown Cu2+ solution =
0.122 / 63.546 ≈ 0.00192 mol/L
By titration
Given the molecular mass of EDTA = 372.24 g/mol,
Theoretical mass of EDTA needed to make 0.01M EDTA solution
= (100/1000)L x 0.01 mol/L x 372.24 g/mol = 0.37224 g
Mass of EDTA used to make EDTA solution = 0.3720 g
\Molarity
of EDTA solution prepared = (0.3720 g ÷ 372.24 g/mol) ÷ (100 ÷ 1000)L » 0.0100 M
Volume of EDTA added (ml)
|
Electrode potential
Reading (mV)
|
Volume of EDTA added (ml)
|
Electrode potential
Reading (mV)
|
0.00
|
235
|
4.70
|
165
|
0.50
|
233
|
4.80
|
118
|
1.00
|
230
|
4.90
|
103
|
1.50
|
229
|
5.00
|
98
|
2.00
|
227
|
5.20
|
95
|
2.50
|
224
|
5.40
|
93
|
3.00
|
220
|
5.60
|
90
|
3.20
|
219
|
5.80
|
86
|
3.40
|
215
|
6.00
|
82
|
3.60
|
213
|
6.20
|
82
|
3.80
|
212
|
6.40
|
81
|
4.00
|
207
|
6.90
|
80
|
4.10
|
204
|
7.40
|
78
|
4.20
|
205
|
7.90
|
76
|
4.30
|
200
|
8.40
|
74
|
4.40
|
197
|
8.90
|
73
|
4.50
|
191
|
9.40
|
71
|
4.60
|
185
|
9.90
|
71
|
Table 2: electrode potential with addition of EDTAGraph 2: graph of electrode potential against volume of EDTA added
From
the graph above, the equivalence point occurs at = 150mV with volume of EDTA added=
4.80mL
EDTA4-
+ Cu2+ è [Cu(EDTA)]2-
No.
of moles of EDTA = 0.00999 × ≈ 0.00004796 ≈ 0.0000480 mol
No.
of moles of Cu2+ = 0.0000480 mol
Molarity
of Cu2+ solution = = 0.00100
M
3. Discussion
Ion-selective
electrodes
Ion-selective
electrodes (ISE) are membrane electrodes that measure the electric potential of
a specific ion in the presence of other ions. They are used in biochemical,
biophysical and environmental analysis for determining the concentration of
various ions in aqueous solution.
In this experiment, a cupric ISE is used. This
instrument consists of a thin solid-state crystal membrane (right diagram) which
specifically permits the movement and transport of Cu2+ from a high
concentration to a low concentration and generates a potential difference which
can be measured by a voltmeter.
Increasing Cu2+
concentration results in more Cu2+ ions attracted towards the
electrode, hence producing a greater current flow. This measurement is done at
equilibrium, where the rate of exchange of Cu2+ across the membrane
is the same. The electrodes can be calibrated by measuring the electrode
potentials in standard solutions of various concentrations and the
concentration of unknown ion in solution can then be determined from the
calibrated curve obtained.
Advantages of using a ISE
Due to the above reasons, ISE are frequently employed to
characterise reactions.
Limitations of ISE
The accuracy of ISE measurements may be decreased due to the
activities of other ions in the same solution[2]. ISEs are not
ion-specific; they are sensitive to ions with similar physical properties to some
extent. For many applications these interferences are insignificant and can be
ignored.
The cupric ISE will not give an accurate reading if Ag or S
are present in the solution. Mercury ions also have very high interference and
can only be tolerated in low concentration compared to the Cu.
Bromide and chloride ions both have selectivity with the
cupric ISE and will cause a significant error if they are present in
concentrations greater than one tenth of that of copper ions.
The accuracy of the results
is affected by the presence of interfering ions. Should there be contaminants,
the results may be inaccurate.
Experimental techniques: by calibration
curve
Before use,
the electrodes must be calibrated by measuring a series of known standard
solutions, made by serial dilution of the 1000ppm solution. For a full
calibration, 100 ml of solutions containing 100, 10, 1 and 0.1 ppm was
prepared. To prepare the various concentrations of solutions, successive
dilutions were carried out carefully.
This must be done with utmost precision: should the concentration of a
preceding solution be wrongly prepared, this will result in a propagation of
errors in the concentrations of successive solutions, thereby further leading
to inaccurately measured conductance for all these successive solutions.
To each standard solution, 0.9 mL of ISA (5M
NaNO3) was added. The Cu electrode works most reliably when samples
and standards are mixed with ISA to give a background matrix of around 0.1M
NaNO3. This keeps the total
ionic strength of the sample and standards constant, therefore ensuring that the electrode potential reading increase proportionally with
increment in Cu2+ concentration (as observed in graph 1).
After determining the electrode potential of
the Cu2+ sample and with the aid of the prepared calibration graph,
the concentration of the sample can then be derived.
Experimental
techniques: by titration with EDTA
Ethylenediaminetetraacetate (EDTA) is a polydentate ligand where the 2 N and
4 O atoms can chelate to the Cu2+ ion to form an octahedral complex[3].
Figure 1: reaction of EDTA with Cu2+
As seen from the equation,
EDTA4- + Cu2+è
[Cu(EDTA)]2-, EDTA reacts with Cu2+ in a 1:1 ratio to form a
[Cu(EDTA)]2- complex. As the volume of EDTA added increases, the
amount of free Cu2+ ions in the solution decreases. EDTA forms a cage-like
structure containing Cu2+, thereby isolating it from the solvent
molecules.
The cupric ion selective electrode which measures the
electrode potential of the remaining Cu2+ ions in the solution will
register a drop in the electrode potential. The initial addition of EDTA
produced a slow and steady decrease in electrode potential reading until when
4.80 ml of EDTA was added. This change marks the equivalence point of the
titration. After that, the addition of excess EDTA continued to result in slow
and steady decrease in electrode potential reading because there was little
free Cu2+ present in solution.
Using titration with EDTA to determine the Cu2+
concentration was less tedious as the
electrodes need not be washed with deionised water so many times. On the other
hand, this method may be time consuming should there be many data points to be
taken; time is required for the reaction to be react thoroughly before the ISE
could register a stable electrode potential value.
Precautions
Prior to each measurement, the electrodes on the dip cell are
rinsed a few times with a dropper containing the solution to be tested. This
displaces any residual ions that may be on it. For the same reason, the beaker
is also rinsed several times with the solution for which conductance is to be
measured.
The probe was then blotted dry. This prevented an
accumulation of static charges (which might occur if it was rubbed dry) and
ensured that the readings obtained were more accurate. A clean, dry tissue was
used each time to prevent cross-contamination.
Time is allowed for the solution to equilibrate before its
conductance is measured. Also, a magnetic stirrer was placed in the beaker
containing the solution to be measured; this ensures even mixing such that a
more representative electrode potential may be recorded. The stirrer may generate sufficient heat to change the
solution temperature; to counteract this effect, a piece of insulating
material, such as Styrofoam sheet, could be placed between the stirrer and
beaker.
The standard Cu2+ solutions were measured starting
with the lowest concentration of 0.1ppm to the highest concentration of
1000ppm; this minimizes the error incurred particularly if there were
contaminants from previous tests.
The experiment was carried out at an environment with
relatively constant temperature as variation in temperature can lead to
measurement error. When not in use, the probe has to be kept moist at all times
to maintain its sensitivity.
Alternative method of measuring concentration of an unknown sample
In this experiment, the concentration of copper ions is derived
from a calibration curve and titration with EDTA. These results may be double-checked
by recording the absorbance of the blue-colored copper-containing sample and
calculating its concentration with Beer-Lambert’s law. According to this law, A
= ε c l, where A is the absorbance, ε is the molar absorptivity, c is the
concentration of the absorbing species and l is the path length of the sample[4].
Once the absorbance of the sample is known, its concentration of copper can
then be calculated.
4. Conclusion
The concentration of the Cu2+ in the unknown solution
determined using the calibration graph is 122 ppm or 0.00192M. Using titration,
the concentration of the copper solution determined is 0.00100 M.
5. References
[1]
Rundle, Chris. Advantages of Ion-Selective Electrode Measurements. Article retrieved on 25 Mar 2012: <http://www.nico2000.net/Book/Guide3.html>
[2] Prince George’s Community College. Structure of EDTA. Article retrieved on 25 Mar 2012: <http://en.wikipedia.org/wiki/EDTA>
[3] Rundle, Chrs. Cupric Ion-Selective
Electrodes. Article retrieved on 26 Mar 2012: <http://www.nico2000.net/analytical/copper.htm>
[4] Sheffield Hallam University. Beer Lamber Law. Article retrieved on 27 Mar 2012: < http://teaching.shu.ac.uk/hwb/chemistry/tutorials/molspec/beers1.htm>
Comments
Post a Comment