0
ELECTRICAL METHODS OF ANALYSIS' A. 1. MARTIN E. I. d u Pont d e Nemours & Company, Inc., Polychemicals Department, Wilmington, Delaware
INDUSTRIAL analytical groups
require a variety of electrometric techniques in conducting general analytical research and in developing methods to aid in the work of other research and production groups. An analytical group must be versatile with these methods,
' Presented as part of the Symposium an An Analysis Group in an Industrial Research Organization before the Divisions of Chemical Education and Analytical Chemistry a t the 130th Meeting of the American Chemical Society, Atlantic City, September, 1956.
recognizing the merits and limitations of each. The chemist in analytical research must have a fundamental understanding of the techniques if he is to adapt them logically to analytical problems. Since the fundamental theories of electroanalytical chemistry are amply presented in the literature (1, 2, 8, 4), no effort will be 'made to cover this aspect. The objective of this paper is to cite some general and specific applications of electrometric techniques in an industrial laboratory.
IOURNAL O F CHEMICAL EDUCATION
POTENTIOMETRY
Direct potentiometry is employed to determine activities rather than concentrations (1). The activity measurements are useful in studying reaction rates or equilibria and in making thermodynamic calculations. The technique is used for pH and other half-cell measurements. For example, electrode potentials of several ferric-ferrous complex ion systems were measured. The potentials were then related empirically to the rate of a reaction in which the ferric complex was reduced to the ferrous complex. A simple potentiometric measurement on a new ferric-ferrous complex could then be used to predict the rate of ferric complex reduction in the subsequent reaction. It must be remembered that for such electrode potentials t o be meaningful the oxidation-reduction system must be thermodynamically reversible. Potentiometric titrations are much more widely applicable for analysis than direct potentiometry and can be employed wherever one or more species will influence the potential of an indicating electrode. The potentiometric method has been applied to titrations involving acid-base, oxidation-reduction, precipitation, and formation of complexes (1). The choice of solvent and titrant is important in these analyses (8). Solvents can be used to enhance the properties which are to he measured but, of course, all species must be compatible in the system. I n general, potentiometric acid-base titrations are straightforward. Typical analyses involving titration of acids and amines are cited in Table 1. Where applicable an aqueous or alcoholic system is recommended for potentiometric titrations, as for example, in differentiating between many mineral and carboxylic acids. Titration of phenols, as weak acids, however, requires a basic solvent such as dimethylformamide (DMF) or ethylenediamine (EDA). Weak bases require an acidic solvent such as glacial acetic acid. TABLE 1 Potentiomatric Titration Methods for Acids a n d Arnines Determination Organic aaids in anhydrides Traces of acetic acid in polymer Inorganic and carboxylie acids, phenols Carboxylic acids, phenols
Titrant
Solvent
Acetone
Tripropylilminc in acetone PyridineNaOH in methbutanol ( l + l ) anal Dimethylform- KOCHI in amide methanolbenzene EthylenediSodium aminonmme ethoxide in ethanalamineEDA Carbaxylic acids, EthylenediKOH in isoprophenols amine Amines Methanol m methan01 Amines, aminoGlacial acetic HCIOn in acetic acids acid acid Secondary plus Methanol HCL in isaprotertiary aminova pan01 Tertiary aminesQet,hsnol HC1 in methand Aliphatic, aroma- Acetonitrile HCI04 in dioxtic amine8" ane
H,y?'
Reference (5) (6)
(4 6, 7) ( 7 , 8 , 9)
(7, 8) (10) (11) . , (12)
(12) . , (IS)
Primary amines removed by reaction with salicylddehyde. Primary and fiecondary amines removed with acetic anhydride. " Method will diverentiate het,ween diphattic and some aromatic and heterocyclic aminpa.
VOLUME 35, NO. 1, JANUARY, 1958
The industrial analyst must be aware of any anomalies associated with titration curves of complex acid mixtures. For example, when an approximately equimolar mixture of p-toluene sulfonic acid, o-phthalic acid and phenol was titrated in water, two sharp breaks and one very poor break were observed. The first sharp break represented titration of the sulfonic acid plus one carboxyl group of phthalic acid, the second represented the second carboxyl group of phthalic acid, and the very weak break, phenol. I n methano1 or D M F three hreaks were observed: the first was equivalent to the toluene sulfonic acid; the second, the first carboxyl group of phthalic acid; and the third, the second carboxyl group of phthalic acid. Phenol in this four-component acid mixture gave no significant break. However, when phenol was titrated alone or in the presence of one or two other acid functional groups in DMF, it was resolved. Titration of thet mixture in EDA showed two breaks. The first included the sulfonic acid and both carboxyl groups of phthalic acid. The second corresponded to phenol (8). An interesting application of the potentiometric method involves titration of organic peracids and
V O L U M E TITRANT
Figur. 1.
The Potsnti*mstric Titration C u m of Wed, Acid. in Ethylenediamine
(The potential at half-neutralization is shown by the solid oircle.)
hydroperoxides as weak acids in EDA using antimony electrodes (14). Figure 1 shows titration curves obtained with several peroxides and with phenol used as reference. As in aqueous titrations, the potential a t half-neutralization, marked by a black dot on the curves, appears to be characteristic of a particular species and is related to its acid strength. These potentials are also related to the structure of the peroxide. Primary hydroperoxides studied have potentials a t half-neutralization of 0.55 to 0.57 volt. Secondary hydroperoxides show potential less than 0.54 volt. Tertiary hydroperoxides show no inflection points. These correlations are valuable in interpreting titration curves of unknown hydroperoxides or peracids. The analysis showed a standard deviation of about 2% relative. Some hydroperoxide mixtures were resolved by this technique. Potentiometric analysis of trace amounts of a weak base was developed for the determination of amine end groups in polyamides. The sample was dissolved in phenol and titrakd with methanolic perchloric acid (15). Since the original development, m-cresol has been found to be a better solvent than phenol. The physical problems encountered in performing this titration which involves 75 ml. of viscous, corrosive solution requiring about 0.1 meq. of perchloric acid can he
TABLE 2 Typical Polarographic Analyses Deteminotion
Samples
Inorganic Cu, Zn, Fe, Pb, Cr, Ca, Ni, Sn Molecular oxygen
Bisulfite ion Oraanic Peroxide groups Carbonyl groups Sulfur containine functions
I
0
Figure 2.
6 8 10 12 14 16 MlLLlMOLES SILVER N I T R A T E Potnntiornetric Titration 0of an Eqvirnolar HalidMixturn
Z
4
I
readily appreciated. Standard deviation of 1.0% relative was calculated from 26 analyses of a reference sample over a period of a year. Another useful potentiometric analysis involves direct titration for low levels of alkali metal salts of organic acids as weak bases (8). For example, a polymer containing a few tenths of a per cent of sodium acetate is analyzed asfollows: a 2-g. sample isdissolved in 100 ml. of glacial acetic acid, the solution is titrated with 0.1 N perchloric acid in acetic acid. Colored solutions require potentiometric titration vhile waterwhite solutions may be titrated using crystal violet as indicator. By performing the determination in this way, tedious and time-consuming ashing procedures are avoided and a less ambiguous answer is obtained. The feasibility of determining iodide, bromide, and chloride ions, separately or in mixtures, by titration with silver nitrate in dilute nitric acid solution has been demonstrated (16). Figure 2 shows a titration curve in which an equimolar mixture of halides was titrated. The potential measurements were made with silver versus calomel electrodes. It was expedient to determine the end points by measuring the intersection of straight lines as shown, rather than by determining the inflection points. Data from the titration curves of halide mixtures show that there are substantial errors in the individual halide determinations but that the total halide analysis is quite accurate. These errors are attributed t o mixed salt formation. However, the errors are reproducible and predictable and therefore subject to correction. POLAROGRAPHY
The polarograph can be used for low level determinations of most cations, some anions, and a number of organic materials (see Table 2). An interesting analytical research program was undertaken to investigate the polarographic behavior of a number of organic peroxides. The work was done in aqueous and nonaqueous media. The buffered aqueous solutions were made to cover the pH range of 2 to 11. The nonaqueous solutions were made up of 3 parts of methanol to 1 part of benzene and contained buffering constituents t o control the acidity. The half-wave potentials of the peroxides studied were essentially independent of pH. The half-wave potentials of classes
Reagents, metals, polymers Organic solvents, polymer solutions Reagent solutions Polymers, solvents, additives Polymers, solvents, additives Polvmers. solvents
of peroxides are shown in Figure 3. This information, along with diffusion current data, has been catalogued for future reference. The polarograph has been used to investigate the reaction of peroxides with various reagents and the stability of peroxides a t high pH as shown in Figure 4. These data are valuable in developing analyses for peroxide mixtures. The fifth reagent, DeeO,= is an enzyme system which reacts specifically with Hz02to form HzO and Oz. Use of the enzyme provides a convenient means for determining organic peroxides in the presence of hydrogen peroxide. Preliminary treatment with DeeO removes only the PEROXIDE FORM
RANGE OF E l,2VALUES
HYDROPEROXIDES PERACIDS D I A C Y L PEROXIDES DIALKYL PEROXIDES PERESTERS
AOUEOUS
SOLUTION
0
0.5
HYDROPEROXIDES PER ACIDS DIACYL PEROXIDES DIALKYL PEROXIDES PERESTERS
E .,I
1.0
1.5
NON A O U E O U S SOLUTION
latter. Hydrogen peroxide can be estimated by diierence by analysis for total peroxide before and after the DeeO reaction. Polarographic principles are also used to advantage in amperometric and dead-stop techniques. These methods combine the accuracy of titration with the sensitivity of polarography. The techniques have the advantage of being readily adaptable to automatic or continuous operation. Some typical amperometric and dead-stop titrations are shown in Table 3. TABLE 3 Typical Amprometric and Dead-Stop Titrations
9
Determination
Titrant
Sulfate Fluoride Halides Cations Water Formaldehyde Merca~tms
Pb(NO& Pb(NOa)n AeNO, Oiine, "Versene" Karl Fisoher Reagent KaHSOa AeNO?
Referenee
(17) (18)
(19)
(4)
($0) (91) (Sf)
Available from Takamine Laboratories, Clifton, New Jersey.
JOURNAL OF CHEMICAL EDUCATION
PEROXIDE VR R S VS
= very rapid
-
reaction
rapid reaction slow reaction = very slow or no reaction
3
t-Bu ACETATE 1-Bu HYDRO D i - t - Bu PERACETIC ACID DIACETYL HYDROGEN F'b4rS 4.
OTHER TECHNIQUES
Constant cathode potential or electrodeposition techniques are used to perform gravimetric determinations of copper, cobalt, nickel, antimony, and lead (3). Electrodeposition has also been used t o remove major quantities of Cu before polarographic analyses for traces of Zn. Now that high-precision current-integrating motors are available, the constant potential technique will become more widely applicable in the analytical laboratory. Constant current and coulometric techniques are being explored. Automatic analyzers are easily designed for a particular problem by coupling a constant current source to a potentiometric, amperometric or photometric end point detection system and a timing mechanism. Continuous analysis of flowing streams should also be possible by applying constant current intermittently as called for by the detection system. The per cent of the time in which current flowed would be an absolute measure of the oxidized or reduced species. A novel application in which a definite need exists is the determination of water by the Karl Fischer method, employing electrical generation of iodine. When this system is clearly understood, the basis for a very sensitive and accurate water analyzer will be in hand. Analyzers are available commercially employing coulometric generation of silver ion for determining active sulfur compounds. Conductance techniques were used a t one time for the indication of the end point in the original titration of amine end groups in polyamides. The solvent was phenol and the titrant, perchloric acid in methanol. Since then the potentiometric method has been shown to be a more effective end point detector. The standard deviation of the conductometric analysis was calculated as 2.0y0 relative from the results of 21 analyses of a representative sample run over a period of 6 months. The evolution of this analysis to the potentiometric method, and then t o the use of m-cresol
VOLUME 35, NO. 1, JANUARY, 1958
Reaction. of Puoaid..
.e*h some Reagent.
as solvent demonstrates an irnportdnt aspect of maintaining high caliber analytical data. LITERATURE CITED (1) DELAHAY,P., "New Instrumental Methods in Electro-
chemistry," Interscience Publishers, Inc., New York, 1954. (2) FRITZ,J . S., "Acid-Base Titrations in Nonaqueous Sol-
vents!'
The G. Fredriek Smith Chemical Co.. Columbus,
0hi0,1952. (3) LINGANE, J . J., "Electromalytioal Chemistry," Interscienre Publishers, Ino., New York, 1953. (4) KOLTHOFF, I. M., AND J. J. LINGANE, "Polarography," 2nd ed., Intersoienee Publishers, Inc., New York, 1952. (5) SIGOIA,S., AND N. A. FLORAMO, Anal. Chem., 25, 797 (1953). (6) PIFER, C. W., E. G. WOLLIGH, AND M. SCHMALL, J . Am. Pharm. Assoe., Sei. Ed., 42, 509 (1953). (7) DEAL,V. Z., AND G. E. A. WYLD,Anal. Chem., 27,47 (1955). (8) MITCHELL, J., JR., B. A. MONTAGUE, AND R. H. KINSEYin
"Organic Anslysis," Vol. 111, Interscience Publishers, Inc., New York, 1956, p. 1. (9) Moss, M. L., J. H. E~LIOTP, AND R. T. HALL,Anal. Chert., 20,784 (1948). (10) HILLENBRAND, E. F., JR., AND C. A. PENT^ in "Organic (11) (12) (13) (14) (15)
Analysis," Vol. 111, Interscience Publishers, Inc., N e r York, 1956, p. 143. FRITZ,J. S., A d . Chem., 22, 1028 (1950). WAGNER, C. D., R. H. BROWN, A N D E. D. PETERS,J . Am. Chem. Soe., 69,2609,2611 (1947). FRITZ,J. S., Anal. Chem., 25,407 (1953). MARTIN,A. J., Anal. C h m . , 29,79 (1957). WALTZ,J . E., AND G. B. TAYLOR, Anal. Chem., 19, 448
(1947). (16) . . SHRINER.U. J.. AND M. L. SMITH.Anal. Chem., 28, 1043 (1956): (17) KOLTHOFF, I. M., AND Y. D. PAN,J . Am. Chem. Soe., 62, 3332 (1940). (18) PETROW,H. G., AND L. K. NASH,Anal. Chern., 22, 1274 cioiinl \----I. (19) LAITINEN, H. A,, AND I. M. KOLTHOFF, J . Phys. Chem., 45, 1079 (1941). (20) MITCHELL, J., JR. AND D. M. SMITH,"Aquzmetry," Interscience Publishers. Inc.. New York. 1948. (21) MARTIN, A. J., unpublished work. (22) KOLTHOFP. I. M.. AND W. E. HARRIS.I d . Ena. Chem.. Anal. ~ d . 18, , kl (1946).
