David R. Williams
University of St. Andrews st. Andrews, Scotland, U.K.
The Uranyl-Acetate System Studied by pH Potentiometry
coordination chemists tend to choose the simpler complexing systems when selecting p H potentiometry experiments for their teaching laboratories. This usually involves metal ions chosen from the second half of the first transition series; occasionally, the lanthanide series may be invoked. This paper reports a simple experiment that embodies the following novelties: (1) it employs an oxycation from the actinide series (the uranyl ion, U022+), (2) it has acetate ions acting as bidentate ligands, and ( 3 ) it introduces the student to two varieties of ligand; those that never leave the U(V1) (the oxide ligands) and those that compete for the vacant coordination positions (the water molecules and acetate ions). Absorption and Raman spectra have shown that U(V1) exists in aqueous solution as UOZ2+ions, the U-0 bonds being inert to substitution and on opposite sides of the ion. This oxycation has six coordination positions arranged in a plane (see Fig. 1). These positions are usually occupied by the water molecules of aquation but under favorable conditions these may be replaced by other oxygen ligands. Ahrland studied the uranyl-acetate system 20 years ago and used a potentiometric approach (1). His results showed that a maximum of three (not six) acetate ions could he accommodated by the six coordination positions and this indicated bidentate bonding involving both of the oxygens on each acetate ion. Crystallographic evidence later confirmed this structural conclusion that required a four-n~en~bered chrlr~trring (2'). The, cxnrrimt~utuscs :r ritration m ~ r r i l u wthus , :xllo\r.ing many measurements to be made in a shirt time. I n theory, as long as the uranyl or acetate concentration is monitored during this titration, the formation curve for the uranyl acetate complexes may he constructed. However, ion selective electrodes for either of these two ions are difficult to make and are somewhat erratic in their response. Hence, the competing cation approach is adopted instead. The nucleophilic oxygens of the acetate ion can either bond to a uranyl ion or to a proton
Figure 1. The structure of the [ ~ 0 z l C H ~ C 0 0 1 a ] - ion. oxygens are in a planor hexagon.
480 / Journal of Chemical Education
The six acetote
(expressed quantitatively, for I
=
lM, T
=
20°,
pKca,cooa = 4.59). If this pK is known, the [H+] of
the complexing solution may be related mathematically to the [CH,COO-1. Then, the formation curve may be calculated from the total concentrations of ionizable protons, TH,uranyl ions, Tu, and acetate ions, TA,and the free concentration of acetate ions, [CHaCOO-I. The experiment involves varying the ratio TH:Tu: TA to construct formation curves and to see whether polynuclear complexes or hydrolysis reactions are present. The Experimental Measurements
A potentiometer is connected to the electrodes of the following cell Ag,AgCI I Cl- ref. s o h 11M NaCIO, ITitrate Solution S I Glass eleotrode
or a NaCl calomel, porous plug bridge ref. electrode (KC104 is insoluble and will block the plug so NaCl calomels must he used). Titrate Solution S has the initial composition TuM UOsZ+ 1M NaClO,
and is titrated with a titrant solution T of composition TN.AMCHsCOONa TEAMCHICOOH (1 - TN.A)MNaClO,
The titrant solution T may be added stepwise from a piston buret (Metrohm Herisau E274) and the titrate solution S may be conveniently thermostatted (20°C) in a magnetically stirred Pye-Ingold 604 titration vessel but, of course, any air-tight constant temperature vessel will suffice. The electrode system must be calibrated by measuring the emf, E, for perchloric acid solutions ( I = lilt)of known hydrogen ion concentration, [H+], and calculating the system's Eo from E = EO (RTIF) In [H+]. Thence, using this Eovalue, E readings for each step of the complexing titration may he converted directly into their corresponding [H+]. Each titration gives a formation curve but to prove the absence of complexes other than UOs(CHsrequires studying a series of different concentrations. A suitable pattern of compositions of solutions S and T is shown in the table, but there is little point in a student studying all these combinations as they have already been fully reported in the literature (1). Nevertheless, to illustrate the principle, we ask the student to attempt two or three differenttitrations and, in the table, the S and T solutions in boldface print are
+
method were log 61 = 2.38, log ,% = 4.36 and log p8 = 6.34 and from half % method used on his curve gives log PI = 2.6, log 82 = 4.7 and log 83 = 6.3. The student curve in Figure 2 gives log 81 = 2.7, log p, = 4.9 and log a = 6.3 by this latter method.
A Convenient Concentration Pattern for the Titration of the Uranyl-Acetate System
Tu (mM)
TEA(mM)
T N ~=A lOOOnM
Any Ta value may be combined with any TEAhut the combinations eivina the most satisfying results are shown in bold
less extremist and are more likely to give the student convincing results. Finally, the student calculates his formation curves from (1) the pK of the ligand (a value that corresponds to the ionic background medium used may be taken from the literature (5) or determined as described in reference (4)), (2) the E, value for the electrode system, and (3) a table of IC ( a log[H+]) versus ml solution T added.
Preparation of Solutions Weighed samples of analytical grades of glacial acetic acid, sodium acetate, and perchloric acid are diluted to the required concentrations. Sodium perchlorate monohydrate may also be used straight fmm the container as long as agood sample is chosen (we use Merck). The analytical grade uranyl nitrate is somewhat hygroscopic so a stock solution is prepared s t the beginning of the term and analysed by reduction to uranium (IV) and then titration against dichromate. These are not the most accurate ways of preparing solutions of known composition hut thanks to the fairly high buffer capacity inherent in the umnyl-acetate-H+ system, errors so introduced are masked and hence the student can save one or two days of analysis.
Mathematical Relationships and Computer Calculations
For each point in the titration the total ligand, TA
+
(= TNUL TEA),uranyl ion Tu, and ionizable proton, T n (= T E n any mineral acid present in the uranyl
+
solution), concentrations and [H+] are known. The formation curves are calculated by a computer program, the mathematics of which is as follolvs TE = [Ht]
+ [CHICOOHI
and so [CHBCOO-I = (TH- [H+I)/([H+]K)
Figure 2.
Hence - log[CH,COO-I is calculated. Similarily T A = [CH3COO-] [CH,COOH] [CHaCOO that is complexed to U022+]and if Z is the average number of CH3COO- per U 0 2 +
+
*_ = [CH.COO.
Uranyl-acetate formation curves obtained by three different
students.
+
that is complexed to UOXZI
Tu
Formation curves of 1L versus - log[CHaCOO-] may then be plotted. Clearly these relationships are simple and the computer program is really only a time saving luxury. Formation Constants
Since all fugacities are held constant by the presence of a background electrolyte, any formation constants derived from these formation curves wiU be defined as concentration constants
There are a myriad of methods for converting formation curves into constants; Ahrland used the Fronaeus method (5) and the necessary X functions are included in our computer output. Methods that can be used by students are discussed in reference (6). Perhaps the most rapid approximate method is to apply the Bjerrum half %approach(7). Ahrland's results from the Fronaeus
Discussion
Titrations taken from the lab reports of three different students are depicted in Figure 2. Several points are worthy of note: (1) Compared to Ahrland's work, our more rapid and much less rigorous approach gives formation curves that closely follow his. (2) The more exacting Fronaeus integral method gives different formation constants to those arising from the half E method. Nevertheless, the difference is small enough for the student to be able to report that he has reproduced Ahrland's constants k 0 . 4 log units. (3) During sessions set aside for problem solving studies, students can be requested to recalculate their constants using the Fronaeus integral method. This graphical approach is useful for highlighting the parts of the formation curve that have unreliable points (large errors). (4) Because titrations of different buffer compositions do not give a family of curves but rather a more or less superimposable set, the possibility of hydrolysis of metal ion or its complexes may be ruled out. Similarly, if curves of different Tu coincide, polynuclear complexes may be assumed absent. Hence, a clean experiment, executable within one day, in which a student meets the coordination chemistry of oxycations of the actinide series and is left with the interesting challenge at the end: why a maximum Volume 48, Number 7, July 1971
/
481
n of 3 when the U0?+ ion has six vacant coordinating positions? A query that, when answered from the viewpoint of a four membered ring (Fig. I), stresses once again that maybe too much coordination chemistry is taught from the viewpoint of the tetrahedra and octahedra of the first transition series. I am indebted to Docent Sten *hrland of the univer~ityof Lund for suggesting this experiment and for suggesting improvements to the manuscript.
482
/
Journal of Chemical Education
Literature Cited (1)
A,.,,,..
s.. acremam.soand.5, IQQ (1851).
~;,",";~;~:.~=,";~;2',.B~,I'~~!;stsbility
Complexes" (2nd. ~ d . )Chem. . SOC. Speciai P U ~NO. I . 17. ~~~d~~ (3rd Ed. in preaaration.) (4) GUENTHER.W. B.. "Quantitative chemistry:M~~~~~~~~~~ and Equilibrium." ~ddi~on-wesky Publ. CO. rno.. 1968. (5) ROSSOTT& F. J. C.. *ND ROSSOTTI, H.. "The Determination oi Stability Constants." MoGraw-Hill Book Co. Ino., London, 1961. (6) G U = N T H ~ RW. . B.. J. CREM. EDUC., 44.46 (1967). (7) B~ennuna,J.. "Metd Ammine Formation in Aqueous Solution:' &ase and son.Copenhagen, 1941.
