stants, derived from the expression given by Dewar (2’) for substituted naphthalenes and from the electronic atom-atom polarizabilities calculated by Coulson and Daudel ( I ) , did not correlate with the shifts nor did plots of the form (1 - 3 co!j2B)/R3, a formula representing anisotropy effects, correlate with the chemical sh[fts. We feel that any attempt to combine quantitatively the effects giving rise t o the shifts would not be meaningful in view of the large number of parameters involved; we therefore prefer to interpret these shifts on a qualitativv basis. The abnormally low chemical shifts for 1,s- and 4,5-methyl groups may be interpreted in terms of a steric effect. Reid ( 1 2 ) and MacLe:tn and Mackor (6) have observed that protons which are sterically hindered give NMR signals at lower fields than thcse which are unhindered. They also found that 4,saromatic protons in lI4-dimethylnaphthalene give lower chemical shifts than the corresponding protons in naphthalene. The downfield shift resulting from steric hindrance may be satisfactorily explained in terms of Pople’s treatment of nuclear magnetic deshielding in an electrostatic field (9), since distortion of the electron clouds in hindered groups is always expected. The relatively large negative contribution (A64 = -0.20) to the 1,s-or 4,5-
methyl shifts is probably due to the strong steric interaction between the methyl groups a t the 1 and 8, and a t the 4 and 5 positions. One may note that this interaction can be represented by Hirschfelder-Taylor molecular models. The strain is probably sufficient to distort the naphthalene nucleus into a nonplanar configuration. Intramolecular distortions of this kind have been revealed by both x-ray diffraction (3) and ultraviolet spectrometry ( 8 ) . Where three or four methyl groups are substituted on one ring, as in 1,2,3- and 1,2,3,4-methylnaphthalenes,large differences between the calculated and measured chemical shift values are observed. These also are probably due to steric interaction and distortion of the ring (IO). [X-ray measurements show, though not conclusively, that adjacent methyl groups in 1,2,4,5tetramethylbenzene cause a slight distortion of the ring (C).] If negative contributions to chemical shift ( - 0.06 p.p.m. for lI2,3-suhstitution and -0.13 p.p.m. for 1,2,3,4-substitution) are attributed to this increasing steric effect; the calculated chemical shift values are in satisfactory agreement with the experimental values (values in the parentheses in Table I). Methyl group shifts for additional compounds are needed to determine this effect with greater accuracy.
ACKNOWLEDGMENT
Grateful acknowledgment is made to G. Dana Johnson for providing most of the samples used in this investigation. LITERATURE CITED
(1) Coulson, C. A,! Daudel, It., “Diction-
ary of Values of 1Iolecular Constants,” Vol. 11, p. 9, Centre de Chemie Theorique de France, Paris. (2) Dewar, M. J. S., Grisdale, P. J., J . rlm. Chem. SOC.84, 3548 (1962). (3) Donaldson, D. SI., Robertson, J. M., J . Chem. SOC.1953, p. 17.
(4) Harnik, E , Herbstein, F. H., Schmidt, G. AI. J., Ibzd., 1954, p,. 3288. ( 5 ) Jackman, L. ?*I., “Application of Nurlear Magnetic Resonance Spectroscopy in Organic Chemistry,” p. 47, Pergamon Press, Xew York, 1959. ( 6 ) MacLean, C., hlackor, E. L., J l o kcular Physics 3, 223 (1960). ( 7 ) Zbid., 4 , 241 (1961). (8) Mosby, IT. I,.>J . .4m. Chem. SOC.75, 3348 (1953). ( 9 ) Pople, J. .4.,M o l . Phys. 1, 199 (1958). (10) Pople, J. A., Srhneider, W. G., Bernstein, H. J., “High R,ysolution Xuclear Magnetic Resonance, p. 426, McGraw-Hill Xew York. 1959. (11) Ibzd., p. 263. (12) Reid, C., J . Mol. Spectr. 1, 18 (1957). RECEIVED for review Xovember 13, 1963. ilccepted ,January 2, 1964. This research was supported by a grant from the American Chemical Society Petroleum Research Fund.
Titrimetric Determination of Zinc with Tetracyanomo no - 1,I (I-phena nt hro I ine-f e rra te (II) ALFRED A. SCHILT and ANTHONY V. N O W A K ’ Department o f Chemistry, Northern lllinois University, DeKalb, Ill.
b
Potassium tetracyano-mono-(1 ,I Ophenanthro1ine)-ferra te(ll) is found to be superior to potassium hexacyanoferrate(l1) as a reagent for the titrimetric determination (of zinc. An amperometric procedure is described that enables titration of 1Oe5M concentrations of zinc. Measurements of stoichiometry, solubility, sensitivity, interferences, and relative accuracy and precision are reported.
T
was undertaken to evaluate the eflectiveness of the divalent complex anion tetracyanomono - (1,lO - phenanthroline) - ferrate(11) as a precipitimetric reagent for the determination of zinc. The [Fe ~ h e n ( C N ) 4 ] -complex ~ ion merits study in this regard since i t is closely related to the hexacyanoferrate(I1) ion, [Fe(CN),j]-4, a well known and useful HE PRESENT WORH
precipitimetric reagent for zinc and certain other heavy metal ions. It was believed that several characteristics of the 1,lO-phenanthroline-substituted ferrocyanide derivative should prove advantageous in comparison to [Fe(CX)b]-4: the anion is intensely colored, thus small amounts may be more readily detected; its formula weight is greater, hence its metal saltq may exhibit lower solubility; and since i t has a divalent rather than a tetravalent charge, i t should exhibit a simpler, more analytically favorable stoichiometry in its precipitation with Zn+Zions. Various disadvantages associated with the classical zinc-ferrocyanide titration have been cited in the literature (1-3, 6, 6). These can be summarized as follows: Precipitate composition depends upon pH. The sohtion must be distinctly acidic to form KIZns [Fe(CN)6]2. Vari-
ations in p H can cause formation of the normal zinc ferrocyanide Zn2[Fe-
(CW, 1.
Attainment of equilibrium is slow unless a n elevated temperature (70” C.) is employed. The use of higher temperature is not only inconvenient but also decreases the sensitivity of the method by increasing the solubility of the precipitate. Use of internal oxidation-reduction indicators requires careful control of solution conditions, and false end points due to adsorption of excess of either reactant are troublesome. External indicators, such as uranyl nitrate, are neither sufficiently sensitive nor convenient for many purposes. Micro amounts of zinc are not satisfactorily determined using the ferrocyanide method. Present address, Department of Chemistry, University of Illinois, Urbana, Ill. VOL. 36, NO. 4, APRIL 1964
845
Many substances interfere. Titration conditions must be carefully controlled to obtain precise results. Solutions of ferrocyanide decompose and must be standardized frequently. An important issue to be resolved by the present study is whether or not substitution of 1,lO-phenanthroline in the ferrocyanide reagent will impart improved properties to its application as a reagent for zinc. If such substitution proves advantageous, then perhaps other or similar modifications in analogous systems might be profitably explored. EXPERIMENTAL
Apparatus. Spectrophotometric measurements were made using a Beckman Model DU spectrophotometer and 1.00-cm. cells. Potentiometric titrations were carried out using a Beckman Model G pH meter equipped with platinum and saturated calomel electrodes (S.C.E.). Amperometric titrations utilized a Sargent Model I11 Manual Polarograph, a stationary platinum microelectrode, a large nonpolarizable saturated calomel electrode, and a n agar-KC1 salt bridge as a junction between the S.C.E. and titration cell. During amperometric titrations, the solution titrated was stirred at a constant rate by means of a Teflon-coated stirring bar and magnetic stirrer. The platinum microelectrode was placed near the side of the titration vessel, as far as possible from the vortex of the solution, to locate it in the region of maximum solution velocity. Titrant was delivered, in all cases, from a 10-ml. microburet with 0.02-ml. graduations. Preparation of Kz[Fe p h e n ( C N ) ~ ] . 4H@. Except for minor modifications, the procedure employed was the same as that previously published (8). A mixture of 5 grams of dicyanobis-(1,lO-phenanthro1ine)-iron(I1) dihydrate and 400 ml. of an aqueous solution containing 40 grams of KCN was heated under reflux conditions for approximately 20 hours. (The reaction time is much less in this case than the 5-day period reported for reaction a t the steam-bath temperature.) Subsequent treatment, to isolate the desired product, was the same as previously reported. Four consecutive recrystallizations from water were made to improve product impurity. The results of various titrations, described below, indicated a purity of 98% for the air-dried product (based upon the formula weight of the tetrahydrate), The disparity from 100% purity proved to be a result of nonstoichiometry in water content of the air-dried product. Conditions were sought that might provide a product of definite composition. Heating at 110" C. in vacuo over MgC104 resulted in nearly complete loss of water; however, the resultant product was very hygroscopic. Drying at room temperature over hfgC104 yielded a hydrate with approximately 2.6 moles of water per mole of complex; this form also absorbed moisture rapidly on exposure
846
ANALYTICAL CHEMISTRY
to air. The air-dried product contains approximately 4 moles of water per mole of complex. Thus a primary standard grade product was not obtained, and it appears doubtful that such can be conveniently prepared. Other Reagents. All chemicals used were of reagent grade and not further purified. Solutions were prepared using deionized distilled water to minimize trace metal contaminations. Standard zinc chloride solutions were prepared using pure zinc metal (Baker's Analyzed 100.0% Zn sheet) and a slight excess of hydrochloric acid for dissolution. The more dilute solutions (lO-3-lO-5M) were prepared fresh before use by accurate dilution of the more concentrated standard solutions (10-2M) with pure water. Standard cerium(1V) sulfate was prepared 0.01Jd in cerium(1V) and 1121 in sulfuric acid using ammonium ceric sulfate. drsenious oxide was employed for standardization. Sodium acetate-acetic acid buffer solution was prepared by dissolving 1 mole of each component in sufficient water to give 1 liter of solution. Solutions and reagents employed for the colorimetric dithizone determination of zinc were prepared and treated in accordance with published procedures (7). Titrant solutions of K2[Fe phen(CN)a] and of K4[Fe(Cs)6] were prepared in approximate concentrations and standardized us. standard zinc chloride solution by amperometric titration. Influence of p H on Precipitation. Solutions of various p H and 0.002M in zinc chloride were treated with less than equivalent amounts of 0.002.V K2[Fe phen(CN)d] to determine appropriate p H conditions for precipitation of zinc. The presence of unprecipitated [Fe phen(CK)h]-* could be readily detected by the color it imparted to the clear supernatant solution. Similarly, the precipitate could be distinguished by its characteristic orange color. Various buffer systems were employed for control of pH: hydrochloric acid, pH 0-2; phthalic acid-potassium acid phthalate, p H 2-4; acetic acid-sodium acetate, pH 4-6; ammonium acetate, pH 7 ; sodium dihydrogen phosphate-disodium hydrogen phosphate, pH 6-8; boric acidsodium borate, pH 8-10; and sodium hydroxide, p H 10-12. The p H of each test solution, after addition of all reagents, was measured using a Beckman Model G p H meter equipped with glass and saturated calomel electrodes. Solubility of Zn[Fe phen (CN),]. A sample of the zinc compound was prepared, and its solubility in pure water a t 25" C. was determined by measuring the concentration of zinc ion in solution in equilibrium with the solid. Zinc was determined colorimetrically as the dithizonate, following a literature procedure (7). To prepare the sample, 100 ml. of 0.01M ZnClz and 100 ml. of 0.01M Kz[Fe ~ h e n ( C N ) ~were ] mixed slowly with vigorous stirring, After aging for 1 day the precipitate was collected by suction filtration, washed generously
with hot water, and dried in air. The filtrate was found to contain only a trace amount of zinc and no detectable amount of [Fe ~ h e n ( C N ) ~ ] ions, -~ hence the formula of the compound was deduced from the stoichiometry of the reaction. Portions of the solid sample were added to pure water (free of heavy metal ions) in two glass-stoppered flasks. One mixture was heated to boiling; the other was not. Both were then placed in a mechanical agitator device in a room thermostated a t 25" rt 2' C. After 24 hours of agitation, each mixture was filtered, and the pH and zinc ion concentration for each was measured. The filtrate from the heated mixture was found to have a pH of 4.65 and to be 3.0 X 10-6.W in zinc ion. Corresponding data found for the unheated sample were p H 5.00 and 3.4 x 10-6J1 zinc ion. The results indicate the molar solubility of Zn[Fe phen (CN),] to be 3.2 X 10-6 a t 25' C. in pure water; on this basis the solubility product is 1.0 X 10-11. Amperometric Titrations. h stationary platinum microelectrode in conjunction with stirred solutions (a "convection electrode") (4) was used. The titration vessel was a regular-form 50-ml. beaker, and the Teflon-coated magnetic stirring bar was 5 , ' ~ inch in diameter and 1 inch in length. ii current-voltage curve was obtained for a solution 10-3.W in Kz[Fe ~ h e n ( C N ) ~ ] and 0.131 in NH4Cl over the range 0-2 volts us. the S.C.E. and using a moderately fast stirring rate. The curve for the oxidation of the iron(I1) complex anion to the corresponding iron(II1) form was well defined, exhibiting a decomposition voltage of about 0.25 volt, a half-wave potential of 0.35 volt, and nearly constant limiting current in the region from 0.7 to 1.1 volts. Oxidation of chloride ions and corresponding increase in convection current began a t approximately 1.2 volts. Known amounts of zinc were titrated amperometrically using rapid, steady stirring rates and an applied potential of 1.0 volt for the platinum microelectrode us. the S.C.E.Solutions were prepared for titration us. standard Kz[Fe ~ h e n ( C N ) ~by] mixing either 5 ml. of 191 ?JH4Cl or 5 ml. of the acetic acidsodium acetate buffer with a measured volume of standard zinc solution and sufficient water to provide a total volume of 25 to 35 ml. The p H of each of the various solutions titrated was well within the range of 1 to 7 previously demonstrated to be optimum for complete precipitation of Zn[ Fe phen(CWa]. For titrations employing standard KFe(CN)e] titrant, known zinc solutions of total volume 25 to 35 ml. were prepared to contain 5 ml. of 1M "&I and sufficient hydrochloric acid to provide a pH of 2. All titrations were performed a t room temperature. Stoichiometry of the Precipitation Reaction. A stock solution of a weighed sample of K2[Fe phen(CX)l] was prepared in a known volume and standardized by two different titration
Table 1. Comparison of Standardization Results from TWO Procedures to Establish Exact Stoichiometry of Precipitation IReaction
Concn. KP [Fe phen(Ch 141 stock so1n.O 0.5513 gram per 100.00 ml.
tlolarity found CR standard Zinc Cerium(1V) chloride sulfate 0 .(11183 0.01187 O.Cl1183 0.01184 O.CIl187 0.01186 0.C11173 0.01178 0.01188 Averages 0.01183 0.01184 0.4578 gram per O.OO982 0.00984 100.00 ml. 0 .Cl0985 0.00978 0 .C'O981 0,00981 0,00980 0.00978 Averages 0.00982 0,00980 a Prepared using product dried over MgC104 a t room temperature. procedures: amperometric titration us. standard zinc solution, and potentiometric titration us. standard cerium(1V) sulfate. The entire procedure was carried out a second time using a second sample of the complex s a l t t o gain greater reliability in t h e interpretation of results. Interference Studies. Solutions of various salts, in concentrations of 0.0002 t o 0.1M, wvre examined for precipitate formation on addition of 0.01 M K 1[ F e phen ( CX)4]. Further tests for interferences were performed by actual titration of known quantities of zinc ir the presence of known concentrations of various salts. The concentration of salt in each rase was t h a t which first failed t o give a visible precipitate in the preliminary tests. Recommended Procedure for Zinc Determination. An accurately measured sample is p u t into solution and freed of interfering substances by means appropriate o t h e nature of t h e sample. Pipet a n aliquot, containing 0.01 t o 1 mg. of zinc and of volume 1.00 t o 25.00 ml., into a 50-ml. beaker containing a magnetic stirring bar inch X 1 inph). Add several drops of methyl red indicator solution (1%) and adjust the p H to approximately 5 by appropriate addition of either ammonia water or hydrochloric acid-Le., until the indicator color just turns from red to yellow or vice versa. h d d 5 ml. of the acetic acidsodium acetate buffw and sufficient water to make a total volume of 25 to 35 ml. With the electrodes in place a n d using a stirring rate as rapid as possible without danger of loss by swirling action of the solution, carry out the amperometric titration employing standard 0.002111 K J F e phen(CN),] titrant and a 10-ml microburet. A potential of +1.0 volt us. S.C.E. is applied to the platinum microelectrode. After adding each increment of titrant, sufficient time (usually 1 t o 2 minutes) should be allowed for establishment of equilibrium before making current measurements; this is especially important
in the immediate neighborhood of the equivalence point. The end point is determined from a graph of current us. volume of titrant in the usual manner. For larger amounts of zinc (1 to 10 mg.), a more concentrated titrant (0.01M) may be employed. Titrant solutions are standardized us. known amounts of pure zinc using the same procedure as that for the unknowns. RESULTS AND DISCUSSION
On the basis of the known ionic charges on the Zn+2 and [Fe phen(CK),]-2 ions, clearly the stoichiometry of their combination should be 1:l. This premise is confirmed by the experimental results compiled in Table I. Statistical analysis reveals no significant difference in either the means or the precisions of the two procedures at a 95% confidence level. The stoichiometry of the precipitation is the same as that of the oxidation-reduction reaction, within experimental error limits. It can therefore be concluded that 1 gram-ion of the iron(I1) form of the complex will combine in precipitate formation with exactly 1 gram-ion of zinc ions. This is a n important result, since it demonstrates that there is no si,nificant tendency for coprecipitation to adversely affect titration accuracy. Studies of the influence of p H reveal that precipitation of Zn [Fe ~ h e n ( C N ) ~ ] from most solutions is quantitative in the p H range 0 to 9.5. In sodium hydroxide solutions of pH 10 or greater, complete precipitation does not occur due to complexing of zinc by hydroxide ions. Ammonia, above a pH of 8, or phosphate species, above a p H of 7 , also prevent complete precipitation by complexation of zinc. Borate, phthalate, and acetate species display no adverse effects. Results obtained on titrating known amounts of zinc with standard K9[Fe phen(CN)4] solutions and with standard K4[Fe(CN)6] solutions are listed in Table 11. The purpose of these titrations was to enable comparison of the two titrants with regard to relative precision and accuracy, applicable concentration rrtnge, and general reproducibility and rate of equilibrium attainment. The various titrations revealed that Kk[Fe ~ h e n ( C N ) ~is] a superior titrant for zinc, especially for small amounts of zinc. A very significant feature of this titrant in comparison with K4[Fe(CN)6] is that individual measurements during the course of any given titration could be obtained with much greater reproducibility and in a shorter period of time. Replicate titration gave very well defined, linear, and almost superimposable titration curves. In the case of the ferroryanide titrant, particularly when the more dilute solutions of zinc were involved, individual experi-
Table II. Comparison of K4[Fe(CN)6] as] Titrants for and KZ[Fe ~ h e n ( C N ) ~ Amperometric Titration of Zinc
Zinc found using std. titrant of Zinc taken, K2 [Fe mg. phen(CN)41 Kd[Fe(CN)c] A. Concentrations of titrants of order 10-zM 2.50 2.50a 2.51 1.25 1.26 1.25 1 .oo 1.00 0 62.i 0.622* 0 500 0 515 B. Concentrations of titrants of order l O - 3 A f 0,500 0.488 0.250 0 24Ev 0 . 267d 0 125 0 124 0.0124 0 011 Average of 6 detns. with relative standard deviation (R.S.D.) of 0.367,. *Average of 6 with R.S.D. of 4.2y0. Average of 6 with R.S.D. of 1.87,. d Average of 5 with R.S.D. of 5.0%. 0
Table 111. Interference Studies Precipitation0 of various metal ions by Kt[Fe ~ h e n ( C N ) ~ ] Molarity of salt Salt 0 . 1 0.02 0.002 0.0002
hig(NO3),
CaC12 SrCll Ba(S03)z Cr(X01)3
hln(NO3)n FeS04 Fe(SO3h Co(x03)2
Ni(NOa)2 CU(;h;Od? AgNO3 Cd(S03)2 Hp?(N03)2 Hg(N03)2 AI C11
0 $
0 0
0 0
+
O
O
++
++ ++ ++ ++ ++ ++0 +0 +
0 0
++ ++ ++ +++ +0 +++ +0
0 0
++ ++ ++ +++ +0 +++ +0
0 0 0 0 0 0 0 0
+0 ++ ++ 0 0
SnCI2 0 SnCll 0 Pb(N03)z 0 NatH.4~04 0 U0dNOn)t 0 The symbol denotes formation of visible, colored precipitate and symbol 0 represents absence of precipitate forrnation on adding reagent.
+
mental points in the titration curves tended to be erratic; titration times were quite long; and galvanometer readings tended to drift, reaching only moderately steady values very slowly. These differences in behavior are mirrored in the differences in relative precision and accuracy of the results given in Table 11. Use of K2[Fe phen(CN)r] as a titrant not only gives rise to better precision but smaller amounts of zinc can be more satisfactorily determined. The smallest concentration successfully titrated was 0.0124 mg. in 25 ml. I n VOL. 36, NO. 4, APRIL 1964
847
Table IV.
Interference Studies
Titration of 0.33 mg. of zinc in presence of known concentrations of various salts Salt Molarity Zn found, mg. Comments 0.32 ”defined titration curve 0.10 0.32 Well defined curve 0 020 0 020 0 32 0 020 0 33 0 00020 0 69 High residual current 0 00020 0 47 Hiah residual current 0,00020 0.34 0 00020 0 33 0.00020 0.00020
0.32
0.34 0.63 0.65
0,00020 0.00020 this case, the titration curve exhibited considerable curvature in the neighborhood of the equivalence point, and the difference in slope of the two lines, measured well before and well beyond the end point, was nearly beyond the limit of experimental discernibility. Thus the level of concentration below which zinc cannot be detected by the titration is approximately 8 X 10-6Jf. An estimate of the for Zn[Fe phen(CN)a] can be made on the basis of this lower limit; the result agrees within a factor of 7 with the value 1.0 x 10-11 found by direct quantitative measurement. Results of interference studies are given in Tables I11 and IV. As would be expected, many of the same heavy metal ions that are known to interfere when K4[Fe(C?;)6] is employed also interfere when K2[Fe phen(CN)r] is used as the titrant. It appears, however, that the latter titrant can tolerate higher concentrations of certain ions than can the former. Relatively large of Mgf2 or concentrations (-0.1.11) AI + 3 and moderate concentrations (
-
0.02M) of Ca+21Sr+*, or Ba+2 can be present without serious effect if the latter titrant is employed. I t is found that Fe+2, C O + ~C, U + ~and , C d f 2 are quantitatively precipitated, simultaneously with zinc ions. Interference due to Ag+ and Hg2+2 ions can be prevented by prior precipitation as the chlorides. Qualitative tests indicate that Pb+2 ions may be eliminated from interfering by adding sodium sulfate to precipitate lead sulfate. Interference by Mn+2 ions can probably be avoided by oxidation to Rho2prior to titration. The p H of the solutions prepared for titration is sufficiently low to prevent interference due to complexation of zinc by cyanide or phosphate species if present. Solution stability of KP[Fe phen(Chr)4] is greater than that for K4[Fe(Cn‘),]. Yo change in molarity was detectable for either a 0.00231 or a 0.01.1f solution of the former, after standing for 1 month under normal laboratory conditions. Under similar conditions, solutions of the latter undergo decomposition as evidenced by
formation of yellow color and a decrease in molarity. The various limitations and disadvantages associated with the use of K4[Fe(CN)6]as a titrant for zinc have been found to be appreciably mitigated through chemical alteration involving substitution of cyanide ligands by 1 , l O phenanthroline. Use of the substituted derivative K2[Fe phen(CS)r] enables titration of smaller amounts and concentrations of zinc, gives results with better precision, significantly shortens titration times, is somewhat less sensitive to certain interferences, and appreciably minimizes the necessity for strict control of titration conditions that are so necessary for precise results when K4[Fe(Cr\;)6]is used. LITERATURE CITED
(1) Hillebrand, W. F., Lundell, G. E. F., “Applied Inorganic Analysis,’’ p. 335, Wiley, New York, 1946. (2) Kolthoff, I. hl., Belcher, R., “Volumetric Analysis,” Yol. 111, p. 659, Tnterwirnce, X e w York, 1957. (3) Kolthoff, I. M.,Furman. S . H., “Potentiometric Titrations,” 2nd. ed.. pp. 323-33, R’iley, New York, 1931. ( 4 ) Kolthoff, I. hI., Jordan, J., J . Am. C h ~ mSOC. . 75, 4869 (1953). (5) Kolthoff, I. M.,Pearson, E. A., IND. ENG.CHEM..ANAT..ED. 4. 147 11932). (6) Richardson, 11. R., ‘Bryson, A., Analyst 78, 291 (1953). ( 7 ) Sandell, E. B., “Colorimetric Determination of Traces of Metals,” 2nd ed.. DD. 623-5, Interscience, Sew York, 1950. ‘
(8) Schilt, A. A., J . A m . Chem. SOC. 82, 3000 (1960).
RECEIVED for review November 18, 1963. Accepted December 30, 1963. Work taken in part from a thesis by A. T’. Nowak, submitted in partial fulfillment for the degree of Master of Science at Northern Illinois University.
Potentiometric Determination of Orthophosphate D. H. McCOLL’
and T. A. O’DONNELL
Department o f Chemistry, University of Melbourne, Parkville,
b Orthophosphate can b e determined potentiometrically in a borate buffered solution by precipitation as silver phosphate at a silver electrode. There is a slight departure from stoichiometric reacting proportions of silver and phosphate solutions as the phosphate concentration increases; but the precision and reproducibility of titration values at any particular concentration are such that a linear calibration curve can b e used with confidence. There is no interference from commonly occurring anions such as nitrate, sulfate, and acetate. Fluoride, which interferes in many phosphate analyses, has no effect on this method. The 848
ANALYTICAL CHEMISTRY
N.2., Victoria,
Australia
other halides not only cause no interference, but can b e determined in addition to phosphate in a single titration.
A
METHODS for macrodetermination of phosphate have been reviewed recently by Rieman and Beukenkamp (4). Most of these are based on precipitation reactions. The precipitate, such as magnesium ammonium phosphate or ammonium phosphomolybdate, may be determined by gravimetric procedures. A gravimetric method involving precipitation of silver phosphate from homogeneous solution has been reported lately (1). However, NALYTICAL
if volumetric procedures are to be used, the initial precipitate may be redissolved and one of its components determined titrimetrically-e.g, the molybdenum in ammonium phosphomolybdate may be converted to Mo(II1) in a Jones reductor and then titrated against permanganate. Alternatively, excess precipitant may be determined titrimetrically-e.g., excess magnesium may be determined with (et hylenedinitril0)tetraacetic acid (EDTA) after precipita1 Present address, Metallurgy Department, South Australian Institute of Technology, Xorth Terrace, Adelaide, South Australia.
