### Lab Report on Ion Selective Electrodes

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.
ISE are frequently used for analysis because[1]:
·      ISE are responsive in a linear manner to the log of activity of analyte.
·      It does not consume the sample being tested.
·      It has negligible contamination.
·      It is unaffected by turbidity and color.
·      It has a rapid respond time.
·      It is relatively inexpensive as the basic setup requires only a meter capable of reading millivolts, a probe for the analyte of interest, and chemicals to adjust the ionic strength of the solution.
·      With careful use and frequent calibration, ISEs can produce accurate results reliably.
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>