Determination of Iron(111) with a Copper Selective Electrode Y. S. Fung and K. W. Fung* Department of Chemistry, University of Hong Kong, Hong Kong
The application of a cupric ion-selective electrode with a membrane of mixed Ag2S-CuS to measure the activity of Fe(ll1) is presented in this paper. The method is simple and to M of Fe(ll1). The rapid with a working range from interference of some common cations and anions has been studied.
Colorimetric method is commonly employed for quantitative determination of dilute solution of Fe(II1) in the presence (1-3) and absence (3)of Fe(I1). However, it is time-consuming, destructive, and not quite suitable to certain applications. It is considered that potentiometry is a more effective method for some applications if an electrode selectively sensitive to Fe(II1) is available. Despite the vast number of different types of ion-selective electrodes produced in recent years ( 4 ) ,there is no commercial electrode made available for measuring the activity of Fe(II1). Attempts for measuring the activity of Fe(II1) with an electrode device have been reported (5-8); however, they are still in a development stage. The use of an electrode device to measure the activity of an ion, which is not incorporated in the matrix of the membrane, has been reported. The determination of cyanide with a halide-selective electrode (9, l o ) ,and citrate and tartrate with a cupric ion electrode ( 1 1 ) are well known. Although the interference of Fe(II1) on solid state mixed sulfide electrodes is well known (9,12,13), such phenomenon has not yet been used for the determination of Fe(II1). In this paper, we report the use of an Orion cupric ion electrode, which contains a mixed Ag2S-CuS membrane, for the determination of Fe(lI1) in aqueous solution and the effect of cations and anions which may interfere with the determination. EXPERIMENTAL Apparatus. An Orion Model 94-29A cupric ion electrode and z Model 90-02 double junction reference electrode were used with an Orion Model 701 digital meter for the potential measurement. A 10% KNO? solution was used in the outer chamber of the reference electrode. A Radiometer 26 with its associate glass microelectrode (Type G2222C) and saturated calomel electrode were used for the pH measurement. Reagents. All chemicals used in this study were analytical grade and were used without further purification. All solutions were made up with double glass-distilled water. The stock solution of Fe(II1) was prepared by using ferric nitrate or 99.9% iron wire and was adjusted to pH less than 2. Standard solutions of Fe(II1) were prepared fresh daily. The salicylaldoxime reagent ( 1 4 ) was prepared by dissolving 1 g of salicylaldoxime in 5 mL of alcohol, and then it was added drop by drop to 95 mL of water at 30 "C. A clear solution was obtained after shaking and filtration. Procedure. Potentiometric measurements were made in a 100-mL glass or PVC beaker (for solutions below 10W M) at 25 OC in darkness to prevent photovoltaic effect. Before each measurement the electrodes were washed thoroughly with double glass-distilled water and then dried. The solution (50 mL) was stirred constantly with a magnetic stirrer embedded in PVC and its ionic strength was adjusted to a constant value with sodium perchlorate. For the investigation of the interference of cation and anion, salts of nitrate and potassium were used, respectively. Ferrous nitrate solution was prepared by mixing ferrous sulfate and barium nitrate solutions under an inert atrnosphere. The timing of the experiment was started as soon as the electrodes were immersed in the solution, and readings were taken a t 1-min intervals for 8 min in a concentrated solution and 15 min for the dilute ones.
RESULTS AND DISCUSSION
Sodium perchlorate was used as the ionic strength adjustor because of its well-known weak tendency of forming a complex. The concentration of sodium perchlorate has no effect M Fe(II1) solution unless on the electrode response for a it is higher than 0.5 M as shown in Figure 1. Since the measurements were made a t low pH (between 2 and 3), all solutions were adjusted to constant ionic strength with 0.1 M sodium perchlorate. A t high pH, Fe(II1) tends to form hydroxo complexes or precipitates as hydrous ferric oxide (15);on the other hand, the interference of hydrogen ion is pronounced at very low pH. Thus, stringent control of p H of the sample solutions is essential to obtain a meaningful result. The influence of p H on M Fe(II1) solution is shown the electrode response for a in Figure 2. The upper p H limit of the solution for various Fe(II1) concentrations is given in Table I. The observed pH limit is lower than the calculated one; it is probably caused by the formation of hydroxo complexes. The lower pH limit of the solution is that the hydrogen ion concentration should be less than 1000 times that of Fe(I1L) concentration. The best M Fe(II1) solution is 2. pH for 10-5 to The potential of the cupric ion electrode vs. reference electrode in various concentration of Fe(II1) solutions ( to M) is shown in Figure 3. The change in potential of the electrode with different concentrations of Fe(II1) is probably caused by the oxidation of sulfide ion around the electrode by Fe(II1) as suggested by Smith and Manahan (13).The rapid oxidation of sulfide by Fe(II1) is a well known reaction (15). The more soluble copper sulfide is leached out from the membrane producing cupric ion and sulfur. Thus a porous silver sulfide matrix is left behind, forming a diffusion barrier for Fe(1II) to reach the region of the membrane containing solid copper sulfide (16).A steady state is quickly established in which the Fe(II1) concentration is very low, whereas the cupric ion concentration is nearly one half of the sample Fe(1II) concentration a t the electrode surface. As a result, an equilibrium potential is reached within minutes and the electrode is sensitive to Fe(II1). The slope of the straight line in Figure 3 is 25 mV per decade increase in concentration. It is less than the Nernst value (29.5 mV a t 25 "C) for cupric ion which controls the electrode potential indirectly ( I 7). This may be attributed to the fact that the reaction reaches different degrees of completeness at different concentration of Fe(II1). However, the exact mechanism for the electrode response is not known. The working range can be extended i o lop6 M of Fe(II1) if the p H is adjusted to 3. Since copper sulfide is continuously ieached out from the membrane of the electrode, especially in concentrated Fe(II1) solution, it is not advisable to use the electrode for high Fe(II1) concentration measurement. The potential drifts toward the negative direction after prolonged use in high concentrations of Fe(II1) solution; however, it can be restored to the original value by polishing the electrode surface with fine alumina paste. The life of the electrode is shortened compared to that used for the determination of Cu(I1). Reproducibility of the electrode can be improved by immersing it in dilute acid (pH 2) or 1 M copper(11) sulfate solution for an hour before use. The electrode ANALYTICAL CHEMISTRY, VOL. 49,
NO. 3,
MARCH 1977
497
1
180
1
0
2
3
4
A
/I
5
- log I Ionic Strength 1 Figure 1. Effect of ionic strength on 1 X
M Fe(lll)in pH 2-3
t
120
.7
5
6
3
L
2
- l o g [FelIIIl1
Figure3. Calibration curve for Fe(lll)in 1 X lo-* M HCIO4 and 1 X
lo-'
M NaC104
Table I. Upper pH Limit for Fe(II1) Determination
Fe(II1) concentration,
M 1x 1x 1x 1x
PHa Observed
10-2 10-3 10-4 10-5
2.25
2.70 3.20 3.65
Calculated 2.68 3.01 3.35
3.68
a Above which the activity of Fe(II1) will be lowered. Calculated from K,, of Fe(0H)B (1.1X
0
2
4
8
6
IO
PH
Flgure 2.
Effect of
pH on 1 X
M
Fe(lll) in 1 X lo-' M NaCIO4
has a marked photovoltaic effect for the determination of Fe(II1). The potential is lower under sunlight and the deviation varies from -10 mV at low concentrations (10-5-10-6 M) to -1 mV a t high concentrations M). Such an effect is much less pronounced for Cu(I1) determination with the same electrode. It is probably caused by the change in activity of Fe(II1) by sunlight (18).Therefore, it is recommended that the cell be kept in darkness during measurements for getting reproducible results. The response time of the electrode, though dependent on the age of the electrode, the stirring rate, the surface condition, and the direction of the activity change (19,20),is found to be fast in the determination of Fe(II1). It varies from less than 1 min in concentrated solution M) to about 10 min in dilute ones (10-5 to M). The response time just quoted represents the time for the potential to reach within 1mV of the equilibrium potential rather than the usual 90-99% of the equilibrium value (19) because the reproducibility is about 1mV in this work. Thus, the accuracy for this method is about 9%. The mixed solution method (12) was applied to study the interference of different cations and anions on Fe(II1) determination. The concentration of Fe(II1) M), ionic strength (0.1 M sodium perchlorate), and p H (2) of the solution were kept constant and the concentration of interferent was varied; then the selectivity ratio was estimated from the intercept of the two extrapolated straight lines of the normal and deviated parts of the E vs. [interferent] curve. The selectivity ratios for some common cations are summarized in 496
ANALYTICAL CHEMISTRY, VOL. 49, NO. 3, MARCH 1977
Table 11. Selectivity Ratio ( K F ~ , M M ;= Cation) of Some Cations (Fe(II1) = 1 X M)
Cation
Selectivity ratio
Cu2+ 1 mL salicylaldoxime reagent" Cu2+ 6 mL salicylaldoxime reagentn
3
CU*+
+ +
Ni*+ Pb2+
Fe2+ Na+ K+ NH4+ Ca2+ Ba2+ Cd2+
6 1 x 10-1
5 x 10-2 3 x lo-"
