The Electrician's Multimeter in the Chemistry Teaching Laboratory Part 2: Potentiometry and Conductimetry Fortunato Sevilla Ill, Rafael L. Alfonso, and Roberto T. Andres Research Center for the Natural Sciences, University of Santo Tomas Espana, Manila 1008, Philippines Instrumentation has become integral with chemical measurements. This trend has to he introduced in the chemistry laboratory course, if the course is to he relevant and up to date. However, the high cost ofmost instruments poses a restriction to the adoption of this approach. The multimeter is a basic equipment of an electrician. It also could be a basic equipment in the teaching laboratory for a foundation course in chemistry. Coupled with sensors and transducers, the multimeter could provide a means for the measurement of chemical parameters. Colorimetric and thermometric systems based on a multimeter have been described earlier ( I ) . Further applications of the multimeter in the chemistry laboratory are discussed in this paper. High-Impedance Vonrnetsr
I
Electrodes
LtW
I
Potentiometry Potcntiometry is the basis of a p e a t number of measurements in the chemistry laboratory. It is an eiectmchemical technique that measures the ele&romotive force in a galvanic cell. Potentiometric measurements require two electrodes exhibiting different reduction potentials and a voltage-measuring device with a high input impedance (Fig. 1A). ---,.
The multimeter can be used to measure the eledromotive force (ern0 of galvanic cells. This measurement is carried out by wnnecting the multimeter, which is set in the voltmeter mode, to the two-electrode system. This measurement should be carried out under reversible conditions wherein no current is withdrawn from the cell. Adigital multimeter fulfills this requirement, because i t possesses a high input impedance. In order to yield emf data, an analog multimeter has to he interfaced to an operational amplifier circuit (Fig. 1B)arranged in a voltage follower mode. This circuit imposes a high input impedance and, therefore, will not allow a cnrrent flow across the electrodes. Reduction Potentials and Cell EMF The reduction potential of electrodes could he determined by coupling electrodes with a reference electrode. An easily assem-
Test Solmion
Operational Amplnier
A
Analog Vomneter
Cdled Copper Wre
saturated NaN03)
-I
Lo., m Figure 1. Electronic circuit for measuring electromotive force (EMF)(A) using a high-impedance voltmeter, and (6) using an analog voltmeter with a voltage follower operational amplifier Figure 2. CuiCuSO, (0.10 M) reference electrode. circuit.
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the cell can be calculated fmm the reduction potentials of the electrodes (2).
Table 1. Determinationof Electrode Potentiala
Half Cell
!?(theor),!F
~(ex~tl),
Zfln(ll), (0.10 M Pb/Pb(ll),(0.10 M) C/Fe(ll),(Fe(lll), (0.10 M)
4.743
4.762
4 . 1 55
4.126
0.648
0.771
0.727 0.800 Ag/Ag(l),(0.10 M) 'Electrode potentials of some half-cellsdetermined using the 0.10 M CUICUSOIreference electrode. The standard reduction potential of the haifcells ate given forcornpatison. b~(expl)= E(cel1)+ Qtheor,for CuiCu(l1)couple) 'Standard reduction potential bled and inexpensive reference electrode is a CulCu2+electrode, which consists of a copper wire (1.0 mm dia.) immersed in a 0.10 M CuSOa kept in a tapered glass tubing (Fig. .. 2). The bottom of the electrode is sealed with cotton. previously soaked in a saturated solution of ammonium nitrate. which serves as a salt bridrre. The too ofthe electrode is sealed with wax to prevent seepage of the solution from the electrode. as well as to control evaporation and prevent . any change & the concentration of thc solution. The test electrode muld involve f a ,a mersl immersed in a solution of its ion or, (b) a carbon rod immersed in a solution containine ions in two oxidation states. The cell is set up by imrnersi;lg the reference electrode in the solution of the test electrode and connectine the two to the terminals of the multimeter. The results obtained for some test electrodes are given in Table 1,together with the theoretical values. The arrangement of the different metals in the list could be recognized as the electromotive series. A cell could be constructed by coupling two test electrodes, and the emf of the cell could be measured in the same manner with the multimeter. The expected emf of
-
-
-1
-0.5
0
0.5
The values measured for some cells are not significantly different From rhe values calculated thrnueh the redunion potentials of the electrodes. The agreemes of these values indicates the sat is fa do^ behavior of the measurine svstem used. Nernst Equation The dependence of the electrode potential on the concentration of the electrolvte solution can be investieated usine the mult~meter.~ g o d test d electrode for this purposeis th; svstem involvinrr the solution ot'FeL*and Fe'* ions and a c&bon rod. several solutions containing varying concentrations of Fez+and Fe3+ions are prepared and a carbon rod is immersed in it to complete thk test electrode system. The half cell is coupled with a CUICU'~ reference electmde a s before, and the eel1 emfis measured. The potential of the ~ e ~ + - F electrode e~+ is obtained by subtracting the reduction potential of the reference from the measured potential. The electrode potential is then Figure 3 illustrates the plotted against log ([Fe2+ll[Fe3+l). linear relationship existing between the electrode potential and the electrolyte concentrations, a s described by the equation,
E," = Eo + 0.05916 log [ F ~ ~ ' V [ F ~ ~ + I
where 0.05916 is the slope of the curve and E, is the y-intercept of the potential exhibited by an equimolar mixture of ions. Measurement of pH pH is a vem important chemical parameter whose value has to be measured often. An eas& constructed electrode for pH measurement i~ the Sb4b,O3 electrode (3,: ~
~
1
log [Fe(lll)l/[Fe(ll)l
Fig~re3. Plot of the var~alionofcell EMF aga nst the ogaritnm of tne ra1.o of tne concentrations of Fe(llJand Fe(li ). The ce I consists of a ~ e " - ~ ee~ectroae ' ano a CLCU ' referenceelectroae.
Figure 4. pH calibration curve forthe antimony electrode. Volume 70 Number 7 July 1993
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whose reduction potential is described by the equation,
The Sb-SbzOs electrode can be constructed by dipping a carbon rod in molten antimony metal. The measurement of pH is carried out by immersing the electrode in a test buffer solution and coupling it with a Cu/CuU.reference electrode that also is dipped in the same solution. The relationship between the measured cell potential and the pH of buffer solutions is shown in Figure 4. The linear plot can be employed as a calibration curve to convert the measured voltage of an unkown solution into pH. Potentiometric Acid-Base Titration
The above pH-sensing system could be used to monitor the course of an acid-base titration (4).This time, the test solution is a titration mixture that consists ofthe unknown acid solution mixed with the added amount of the base solution. Voltage readings are taken during the course of the titration. The titration curve obtained for the solution of ~ o t a s sium hydrogen phthalaw reacted with 0.10 M ~ a 0 H ' i des oicted in Fimre 5. The eauivalence wint is indicated bv the abruptvchange in the measured potential and thk eauivalence ooint can be estimated from the inflection p& in the &we. Potentiometric OxidatiowReduction Titration
A redox titration could be monitored using the multimeter. The test electrode to be used in the measurements is a carbon rod immersed in the titration mixture and coupled with the Cu/CuS04(0.10 M) reference electrode. The voltaee of the cell is recorded durine the course of the titration. ~Yheresults obtained for the titration of Fez+solution with K9Cr9O7is shown in Firmre 6. A siemoid behavior similar to-that acid-base titgtion also is obtained.
_I
2
0
4
8
8
10
12
rnL 0 2 0 7 Figure 6. Potentiometric titration curve of Fe(ll) versus 0.10 M K2cr24. Conductimetry
Conductimetry is the measurement ofthe ability of solutions to conduct electricity. Solutions differ in their conductance depending on the number of ions and their mobilities in these solutions. This difference is comrnonlv demonstrated using the familiar conductivity apparatusthat consists of a pair of electrodes in series with a light bulb and the main power supply. The intensity of the light emitted bv the bulb indicates the demee of conductance of the solutfon between the electrode; The multimeter enables a quantitative measurement of conductance. Asimple set-up is shown in Figure 7. A transformer reduces the mains supply (220V, 60Hz) to 12V The low voltage makes the experiment safe to perform and the alternating current prevents the occurrence of faradaic or electrolytic processes. A conductimetric probe can be constructed using carbon rods from spent dry cell batteries and securing the rods on a rubber stopper held by a PVC tubing. The multimeter is connected in series with the probe and is set in the went-measuring mode. The magnitude of the current measured by the meter provides a measure of the conductance of the solution.
Prim.
Sec.
II PVC Tube
Rubber Stopper
mL NaOH Figure 5. Potentiometrictiirationcutveofpotassium hydmgen phthalate (KHP)versus 0.10 M NaOH. 582
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14
Carbon Rods
Figure 7. Diagram of set-upfor conductimetric measurement.
Table 2. Effectof the Nature of the Solution on Conductance
Solution
Current (A)
0.10M NaCl
0.27
0.10M HCI
0.67
0.10 M CH3COOH
0.02
0.10M NHzCONH2
0.00
Conductance of Solutions Conductance experiments often are employed in the laboratory to demonstrate the difference between an electrolyte and a nonelectrolyte and between a strong and a weak electrolyte. I t also can illustrate the difference in the mobilites of different ions. All of these can be obtained by comparing the conductance of the solutions of a weak and strong electrolyte, a nonelectrolyte, and strong electrolytes that differ only in one of the ions. Table 2 shows the current readines when the robe was dipped in 0.10 M solutions of aodiumchloride, hidrochloric acid. acetic acid. and urea. The hfference in the conductance of hydrochloric acid and acetic acid shows the difference in the extent of ionization of the two acids. The difference in the readings between hydrochloric acid and sodium chloride demonstrates the ereater mobility of hydro~en . ions in comparison with sodiim ions. ~~
~~
~
~
~
Variation of Conductance with Concentration
The conductance of an electrolyte solution is dependent on the concentration of the electrolyte. The conducting power of an electrolyte solution is attributed to the movement of the ions in the presence of an electric field. The moving ions carry a portion of the current flowingthrough the solution. As the number of the ions increases, it is ex-
0
5
10
15
20
mL NaOH Figure 9. Conductimetrictiiration curve of HCI versus 0.10M NaOH. pected that the current flow in the solution should increase correspondingly. This trend is observed in Figure 8 wherein the current increases almost linearly with the square root of the concentration of hydmchloric acid (5). Conductimetric Acid-Base Titratlon
The course of an acid-base reaction can be monitored through the change in conductance caused by the replace-
0
5
10
15
20
mL NaOH Figure 8.Piot of the variation of conductance against the square root of the concentration of HCI.
Figure 10.Conductimetric tiiration curve of acetic acid versus 0.10M NaOH. Volume 70 Number 7 July 1993
583
ment of the highly conducting hydrogen and hydroxide ions by ions with lower conductivity. During the titration of a strong acid, such as HC1 by a strong base, such as NaOH, hydrogen ions are removed fmm the solution as water that is weakly ionized and is replaced by an equivalent number of the less mobile sodium ions. The conductance of the mixture, therefore, decreases during the course of the titration. Beyond the equivalence point, further addition of the titrant now introduces highly mobile hydroxide ions in the solution. Since no more neutralization can occur, the conductance of the solution increases with the addition of more titrant. These changes are depicted in the wnductimetric titration curve in Figure 9. In the titration of a weak acid, e.g., acetic acid, with a strong base, e.g., sodium hydroxide, the neutralization reaction results in the conversion of the poorly conducting acid into a highly ionizable salt leading to an increase in conductivity. Beyond the equivalence point, the conductance changes more steeply because of the addition of highly conducting hydroxide ions into the solution. Figure 10 shows the conductimetric titration curve of acetic acid with sodium hydroxide. Concluslon This paper cites only some of the experiments that can be ~erformedin the chemistrv laboratory with the help of a multirneter. in each ofthe m e a s ~ r i n g ~ ~ s t edescribed, ms the multimeter is combined with a sensing probe that de-
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tects an electrochemical property of the system. All of the components used in the set-up are available easily from an electronic supply shop. Modem technology is providing novel materials that can act as transducers in a measurement process. Numerous experiments can be developed to utilize these devices to illustrate imoortant chemical conceots in the teachina lab--oratory. ~ h i approach s will promite not only an understanding of chemistry but also an appreciation of the general philosophy of instrumentation. This "clear box" a ~.~ r o a to c hinstrumentation will lead to an awareness of the developments in material scienw and could encourage creatintv in the utilization of the outputs of modern technology.
.
Acknowledgment The authors gratefully acknowledge the financial support awarded by the United Nations Educational, Cultural, and Scientific Organizations (UNESCO). Literature Cited 1. Andrea, RoberbT; SevlllalII, Fmtunata. J Chem Educ. 18 , X X , m 2. Popiel, W J.Labomfory MonvnlofPhysimlChemiatry:Englieh l a n g u v B a o k Sodetv: London. ~, . 1964. 3. h s , D. J. G. In oxide, L b y ~ e n .and S u l f d e E k t m d e s , in RefpmfpfpEktmdpsi lvea. D. J 0.; Jsnz,G. J., Eds., AcaduoicPreas:NearYorL, 1961. 4. Fritz.J. 5.: Schenk. G.H. Q v o n l i m t i o p h I y f l m l C h e m i s m . 4 t h e d . , ~ y nandBaeon, In".: Boston, 1966. 5. Maran, S. H.; Land0, J. B. Fundomntols ofPhyslml ch2miatry; Macmillan Pubhhing Go. In=.: New York. 1974.
