Chelometric Titrations with Potentiometric End Point Detection Mercury as p M Indicator Electrode CHARLES N. REILLEY and R. W. SCHMID Deparfmenf o f Chemisfry, University of North Carolina, Chapel Hill, N. C. ,The mercury electrode may b e applied under certain conditions as a pM indicator electrode for numerous metal ions. The theoretical principles concerning the use o f this electrode in chelometric titrations are presented. The effects of pH, buffer, chelon, and metal ion on the potentiometric titration curve have been studied b y means of potential-pH diagrams. Interference levels of competitive redox systems, such as mercurous halides or oxygen, have been evaluated similarly. Potential-pH diagrams allow a prediction of the extent of the end point break and thus a choice for optimal titration conditions. The mercury electrode may also b e used to characterize a new chelon for its interaction with various metal ions and its applicability as a titrant.
B
THE CHELOSS (a term used to describe complexing agent.. such as E D T A and the polyamines) form stable 1 to 1 complexes with most iiietal ions, standard solutions of them can be used as the basis for the volumetric determination of these metal ions. I n order to have a successful method, however. certain solution conditions-selection of pH, proper buffer, masking agents-must be established according to the nietal ion titrated and tlie particular chelon used. A compatible and sensitive end point detector inust also be found. ECAUSE
FACTORS INFLUENCING METAL-CHELONATE FORMATION
The reaction of the metal ion with the chelon should, of course, be extensive for the metal ion to be titrated and slight for the other foreign iiietal ions. The following diagram, illustrating the reaction between a bivalent iiietal ion and EDTA, is helpful. AIetal
)IT+
+
IliZ
+ chelon metal-chelonate + 1--4 + A t ? - +
+
nH +
Complex pH
effect
effect
1 -+Hh.IT-
4h4 Y OH - - -
Metal chelonate derivatives
pH Effect. T h e diagram shons t h a t t h e hydronium ion competes with t h e metal ion for t h e chelon. For this reason metal ions nhich form only weak chelonates (such as alkaline earths) can be titrated effectively only in alkaline solution. On the other hand, metal ions n hich forin very stable chelonates (such as mercury, bismuth, and copper) can be titrated even in acid solution. The minimum pH a t IT hich metal ions of giren stability can be effectively titrated n i t h E D T A \\ as calculated and the results are plotted in Figure 1. This calculation is based on the soine\I hat arhitrary restriction that a minimum effectiT e stability constant of 106 is necessary for a satisfactory titration of 0.01Jf solution of metal ion, [see also is)]. The effective stability constant, K E , is given by the following cupression:
from the tlierniodyiinmic Ftaiiclpoiiit, but the kinetics of the reaction between chelons and precipitates are usually slow. Consequently, hetter results are obtained if tlie precipitate is first eliminated through use of the complex effect. M eta1-C helonat e Derivatives. Tlie ability of t h e metal-chelonate t o form acid derivatives (HMY-) tends t o increase t h e effectiveness of t h e titration of certain metal ions in acid solutions. This means that certain metals may be titrated effectively in solutions more acid than one would expect solely from consideration of Figure 1. The formation of the MYOH--- in alkaline solutioiis means that certain metal ions react more extensively n ith the chelon under these condition.. The formation of derivatives of the type 1\/IYNH3-- also points to more exteiisire reaction in the presence of am-
n here Ii is the stability constant of the metal-clielonate and
K 4 ,K 3 , K 2 . and k', are the acidity constants of the chelon. Complex Effect. T h e presence of a complexing or precipitating species, 2, shifts t h e metal-chelonate equilibrium t o t h e left. Purposeful addition of selective complexing agents or precipitants often allows masking of certain metals, permitting titration of others. Most buffer chemicals are also complexing agents. Their preence can present certain disadvantages -for example, the p l I break a t the end point is decreased. The complex effect of buffer constituents also can elert a useful purpose. Most metal ions, for example, precipitate in alkaline solutions unless they are complexed to some extent v i t h a buffer or some purposefully added complexing agent. The precipitate itself is not injurious
moiiia. [The latter reaction forms t h e basis of a method for the volunietric analysis ( 2 ) of ammonium salts without the usual distillation.] METHODS FOR END POINT DETECTION
Even though suitable conditions are established (by use of the above principles) so that a desired chelometric titration reaction will occur. such effort is in vain if a suitable end point detection device is not available. Titration of Liberated Acid. This particular method is based upon t h e titration of t h e acid liberated when excess chelating agent is added to a solution of t h e metal ion.
VOL. 30, NO. 5, MAY 1958
e
947
A p H meter or suitable chemical indicators can be used. This method is only of academic interest a t the present time. Use of Metal-Indicator Dyes. T h e application of metal-indicator dyes in chelometric titrimetry has focused considerable favorable attention to the practicality of the chelon approach to metal ion analysis. I n principle, the dye complexes with the metal ion to form a metal-dye coniples differing in color from the free dye: hl+"
+ dye-"
e.g., blue
28 26
24. 22,
'"i 18
+ Mldyeu-n
e.g., pink
Upon titration with chelon, the metal ion is removed from the metal-dye complex a t the equivalence point with a resulting color change. This method, when applicable, is the easiest and most practical approach. I t s limitations are as follows.
I4 -
IO 12
The metal ion must form a complex with the dye of just the correct stability. If the metal-dye complex is too tight, late or no end points are obtained. This effect can only occasionally be remedied by changing the chelon titrant employed. If the metal-dye complex is too loose, early end points are obtained and the presence of the buffer frequently accentuates the difficulty. For this reason, a given metal-indicator dye is generally useful for only a few metal ions, and then only under limited conditions of p H and buffer type. Another severe limitation is encountered in cases where the metal ion or metal chelonate is highly colorede . g , with chromium, vanadyl, strong copper, or iron solutions. The visual detection of the color change of the indicator is then difficult. Physicochemical Methods. Conductometric titrations are limited in scope because of the large concentrations of buffer which must ordinarily be employed. T h e amperometric method does not suffer in this respect and is generally applicable, especially when the appearance of the anodic wave of excess chelon is used to detect each end point. The purpose of this study was to develop a potentiometric end point system n-hich could be applied with several types of chelons, most metal ions, in most buffer solutions, and o ~ e r a maximum range of pH. Because the negative logarithm of the free metal ion concentration ( p l I ) will vary in the course of a chelometric titration in a way perfectly analogous to the variation of p H in an acid-base titration, an electrode system which would respond to -log [,II-"]would be desirable. The mercury electrode can furnish such an electrode system (S). Siggia ef al. (9) employed a mercuryplated platinum electrode in an empirical manner for detecting end points
948
ANALYTICAL CHEMISTRY
Figure 1. Minimum pH for effective titration of various metal ions with EDTA
8-
2
0
4
6
8 1 0 i 2 1 4 PH
2
4
0
6
0.5
10
1.0
15 a
PH
Figure 2. Prediction of potentiometric titration curve from potential-pH diagram (neglecting buffer effect)
Left. Potential-pH diagram for EDTA system a t 25" C. and ionic strength of 0.1 (NaC104) 0.00111.1 EDTA I. 0.001M HgEDTA:2 0 H - + HgO H20 2e11. Calculated potential for Hg Horizontal l/nes, 0.001M HgEDTA-0.001M MEDTA-0.001M M + + Right. Titration curves derived from potential-pH diagram
+
+
in chelometric titrations. The present paper describes in detail the functioning of this type of electrode system under a wide variety of solution conditions. MERCURY AS pM INDICATOR ELECTRODE
A mercury electrode in contact with a solution containing metal ions (to be titrated) and a small added quantity
+
+
++
of mercury(I1)-chelonate, HgY*-", exhibits a potential corresponding to the folloa-ing half-cell: Hg 1 HgY2-", MYz--, M + + where Y-" stands for the completely dissociated chelon. The potential for this indicator electrode may be found by combining the h'ernst equation for a mercury electrode E = EOee 0.0296, log [Hg++] (3)
+
with the equations for the stability constants of a 1 to 1 mercury(I1)-chelonate,
and a 1 to 1 metal-chelonate,
This yields, a t 25' C., Eag = EOH,
-+
Equation 6 shows that the potential of tlie electrode bears a linear relationship with log [JIT*],and consequently, the electrode may serve as a pl\I indicator electrode in the presence of fixed concentrations of the metal-chelonate and mercury(I1)-chelonate. Although other metal electrodes could lie used, several features render tlie mercury electrode system superior. First, the mercury electrode is highly reversibleeand it has a high hydrogen overvoltage. Secondly, mercury chelonates generally are of high stabilitj- compared to the chelonates of other metals. This is desirable because the electrode responds properly only to metal ions: M liich form cheloiiates weaker than the mercury-chelonate. Furthermore, this high stability permits the use of the electrode for titrations carried out in acid solutions. Finally, the relatively positire potential of mercury niininiizes the interference of oxygen with the elrcatrode reaction, as well as contaniination of the surface n i t h oxide coatings. Chelon Characterization with Mercury Electrode. The niercuiy elrctiode is useful not only for detecting end points in clielonietric titrations. but also for determining in a simple \I ay t h e chelating properties of a given chelon under a wide variety of conditions. Reillq- and Porterfield describe this aaalication for ammoniacal solutions (Si. From Equation 6, the potential of the mercury electrode is seen to depend linearly also on the log K of the particular metal chelonate involved. The log Z i can thus be calculated from experinirrital data by means of Equation 6, using the observed potential of the niercury electrode, the concentrations of the compounds occurring in Equation 6, and the stability constant of the niercury(I1)-chelonate, K H g y . The latter also can be determined n i t h the mercury electrode. The potential of a mercury electrode in a solution containing niercury(I1)-chelunate and excess chelon is given by
where
the coniplexing agent, and [I'll is the stoichiometric concentration of the chelon. K H ~ Y can thus be calculated from the experimental potential-pH function. Curve I in Figure 2 shows this function for a solution containing eaual amounts of free chelon (EDTA) ind'the corresponding mercury-chelonate. For practical determination of stabilitJ- constants, a potential-pH diagram is constructed froin experimental data obtained with solutions containing the chemicals involved in Equation 6 in definite concentrations. I n Figure 2, left, such data are plotted (horizontal lines) for the case of EDT,I, the concentrations in this case being 0.001M for the metal ion, the metal-chelonate, and the mercury-chelonate. For these concentrations the potentials corresponding to various stability constants xere calculated by means of Equation 6 and plotted on the potential-pH diagram. The stability constant for a particular metal ion is then directly read from the potential-independent region of the experimental points. This method of determination of stability constants of metal-chelonates has been described in detail ( 6 ) . Potential during Chelometric Titrations. I n t h e course of a chelometric titration the potential of the mercury electrode varies in a way analogous t o t h e potential of a glass electrode in a n acidimetric titration.
E
=
E%,
+
Curre I indicates the potential past the end point in the presence of a definite excess (where [H,Yn-4] = 1HgY2-n]) of chelon: E = E'E,
+ 0.0296 log K H ~ Y (15) ~
On the right side of Figure 2 are given some theoretical titration curves which follolv from a consideration of the potential-pH diagram. The points on the right side are derived from corresponding points on the left side. Consider the E D T A titration of manganese a t a p H of 5, where a n amount of mercuryE D T A equivalent to one half of the total manganese concentration has been added. It is seen from the potential p H diagram that A represents the most positire potential the electrode can assume. The mercury ions which establish this potential are furnished through displacement from the HgY-- complrx by manganese ions. Point B is reached a t half titration, point C a t 50% overtitration. The potential difference between B and C is a measure for the extent of the end point break. At p H 7.3 the potentials would be different. The most positive potential Acidimetric titration: possible in this case is D. This potential is even less positive than the one correB H + + HB+ sponding to half titration at lower p H values. The potential of the electrode E E'H, is therrfore leveled a t 250 mv. u p to the 1 - 0.059 pK (9) vicinity of the end point. The end 0.059 log _[B _ IBH-I point extends until G, and the break is much larger than a t p H 5 . I n the presChelometric titration: ence of a n equivalent quantity of barium hI + Y -n + NY"-" ion, the potential would drop only to [?*I+"]approximately F after the end point. E = EoaA 0.0296 log [ 11T -n ] A useful second end point, after the titration of barium, would not occur a t this pH, for the potential difference, F - G, is too small. I n the presence of a 10 times larger amount of barium, The concentration of the niercuryF would be 29 niv. more positive (Equachelonate is kept constant by addition tion 6), and a smaller end point break of a small amount of this compound to for the manganese titration ( D - F ) the solution prior to titration. would be obtained. Finally, a titration The potentials expected prior to and of calcium a t p H 7 would follow the after the end point-and thus the extent path, DEG. of the end point break-can be predicted with help of a potential-pH diaI n this way the potential-pH diagram. Curve I1 in Figure 2, left, repregram readily reveals the expected insents an upper limit imposed on the fluence of p H and of the nature of the potential by formation of mercuric ion titrated on the end point. It must oxide. be emphasized that Figure 2 represents Hg 2 0 H - - 2 e - ;r: HgO + H 2 0 (11) the case in n hich buffers with complexing properties are absent. I n the presE Eong 0.0296 log _ K H g_ o (12) ence of coniplexing buffers a different [OH-1 potential-pH relation, and consequently where (1) different end point characteristics, may be obtained. K H ~ =O [Hg++] [OH-]' = IO-*j,5 (13)
+
+
+
~~
+
+
and
STUDY OF SOLUTION CONDITIONS =
K J stands for the acidity constants of
indicate the potential expected halfaay toward the end point for the various metal ions :
0.612 volt
US.
S.C E.
The horizontal lines, corresponding to solutions containing equal amounts of free metal ion and metal-chelonate,
Effect of pH. The stability of a metal-chelonate decreases with decreasing p H according t o Equation l . T h e extent of t h e end point break, VOL. 30, NO. 5 , MAY 1958
949
w
c;
vj ln
>
-I
a I-
z W
-
I-
O il
pH=9
M L. Figure 3. Potentiometric titration of calcium with EDTA a t various PH values
2
8
6
4
IO
PH
Figure 4. Effect of buffers which d o not form complex with HgEDTA on potential prior to end point
(25' C.) I. Potential past end point unaffected 11. See Figure 2 a . 0.001M Hg'+ in 0.1X triethanolamine b. O,0OlL44Hg++ in 0.1M acetate c. 0.001M Hg++ in 0.1M tartrate (1. 0.00131 Hg++ in 0.1M citrate
3001
I 2
I
I
4
I
I
I
I
I
8
6
1
\
I
IO
PH
Figure 5. Effect of buffers which form complex with HgEDTA on potentials prior to and past end point
(25' C.) I, 11. See Figure 2 A . 0.001JJ Hg++ in 0.lM hexamethylenetetramine B . 0.001M Hg+' in 0.1M pyridine C. 0.001M Hg++ in 0.1M ammonia D. 0.001M H g + + in 0.1114 tris(hydroxymethy1)aminomethane E. 0 O O l M HgEDTA-0.001.V EDTA in 0.1M hexamethvlenetetramine P . O O O l h HgEDTA-0001J.1 EDTA in 0.1M pyridine G. 0.001Jf HgEDT-4-; 0.001M EDTA in 0.1M trisH.
+ + + ( hydroxymethv1)aminomethane 0001J.1 HgEDTA-- + 0.00lM ammonia
950
ANALYTICAL CHEMISTRY
EDTA in 0 1 V
-1001 3
Figure
1
4
6.
I
5
I
6
Direct
7 PH
8
9
complex
1
0
effect.
(25' C.)
+
0.00134 EGTA I. 0.001M HgEGTA-in 0.1M KaC104 11. See Figure 2 a, b, c, d. O.OOliC.1 HgEGTA-- + 0.001il.I ZnEGTA-- + O.OOlJ.1 ZnT+ in O.liC.1 electrolyte Electrolytes. a, NaC104; b, pyridine; c, triethanolamine; d, citrate
therefore, is a function of p H . T h e potential-pH diagram (Figure 2) shows this effect clearly. The distance between the horizontal lines and curve I becomes smaller toward lon-er p H values until the lines merge at the p H where the metal-chelonate is no longer stable. I n order to obtain a useful end point break in a titration, a difference of a t least 120 m r . is necessary between the t n o lines. End point breaks obtained a t various pH values for the case of the potentiometric titration of calcium are shown in Figure 3. The potential-pH diagram can be consulted for selection of the proper p H for a titration. This feature is particularly useful for choosing the conditions for the analysis of a two- or multicomponent mixture on the basis of the p H effect. Thus for example, Figure 2 shows that zinc or manganese could be titrated a t p H 5 in the presence of calcium, magnesium, and barium. A t p H 9, all of the metals would be titrated with a single break a t the end point. Effect of Buffer. A buffer is necessary t o maintain t h e desired p H during t h e course of t h e titration. M a n y common buffers form stable complexes with mercury(I1) ions, with t h e mercury-chelonate, or with t h e metal ion to be titrated. The formation of such complexes produces corresponding changes in the potential-pH diagram. UPPER PONTEKTIAL LIMIT. I n t h e presence of a complexing agent t h e potential of a mercury electrode is no longer limited b y the formation of mercuric oxide, but b y the formation of a corresponding mercury complex. The limiting potential is then less positive. In Figures -1 and 5 these potentials are illustrated for a number of coninion buffers. I n the case of ammonia, for example, this potential corresponds a t p H values less than 9 to the reaction Hg(lJH3)2+-
+ 2 H + + 2eHgo
+ 2XH4+
was calculated with help of Equation 16. I n an analogous way, stability constants of complexes of the type HgBz between mercury and other buffer bases, B, were determined. These are listed below, together with the pK, of the buffer. logk' logK Ammonia Tris(hvdroxvmetLy1)aminomethane Pyridine Hexamethvlenetetramine Triethanolamine
PKa 9 3
(HgBZ)
(BHgY)
176
6 4
8.1 16.1 5 . 1 5 10.8
4.3
4 85 7 8 13'8
4 1
