isopoly species originally added is thc
MOTanion, and on the basis of general isopoly chemistry, it is presumed at these acid concentrations that the original species is aggregated through hydrolysis to a higher complex. “Red” stands for reduction, and refers t o the addition of sulfuric acid and chlorostannous acid in the reduction step. ‘%lueJJmerely means that the species under consideration is reducible t o a blue hue. These series of equilibria predict that since only one reducible species is present at pH 2.0, zirconium only serving to bleach the isopoly color a t lower temperatures, then the “blue” spectra of blanks and zirconium-containing solutions should be identical in shape. Similarly, since two reducible species are proposed as being available for reduction at pH 1.0 depending on the temperature, the spectra of blanks and zirconium-containing solutions should be different in shape. Four selected “blue” spectra n ere accordingly run as shown in Figure 4. The conditions are indicated by cell designations of Table V. The same system of plotting is used as for Figure 1. Comparison of these spectra indicates that the prediction holds. The pH 2.0 spectra are identical, while the pH 1.0 spectra are not. The following Am,, were observed: for both pH 2 . 0 ’ ~710 ~ mp; for the pH 1.0 blank, 705 mp; and for the pH 1.0 zirconium, 730 to 750 mp.
CONCLUSIONS
Perhaps the most important item about this method is the fact that an apparently new type of heteropoly complex, a molybdosulfatozirconate, is formed. Although the method is more sensitive than the molybdozirconophosphate procedure, because of the strict control of sulfate and the number of interferences, it is probably useful only for estimations of zirconium.
W
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500
700
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WAVELENGTH, mfi
Figure 4.
Spectra of reduced species
A.
pH 1.0, blank, 25’ C. B. pH 1.0, Zr, 100’ C. C. pH 2.0, blank, 25’ C. D. pH 2.0, Zr, 25’ C.
The existence of a sulfatomolybdate, the proposed intermediate in the complex formation step. has been reported (8). It is reported to be reducible under extreme conditions. That Zirconium, or for that matter any metal ion, would form an easily reducible complex with this species has not been previously observed. The proposed equilibria take into account the observed experimental facts and a t least qualitatively explain the proliference of interaction terms observed in the analyses of variance. Synergistic effects would be expected if species changes, dependent upon concentration and conditions, were occurring in the system.
LITERATURE CITED
(1) Ainsworth, S., Rabinowitch, E., Science 131, 303 (1960). (2) Guyon, J. c., Mellon, M. G., ANAL. CHEM.34. 856 (1962). (3) Hallopeku, R., -L4nn. Chim. Phys. [7] 19, 92 (1900). (4) Holt, A. S., Jacobs, E. E., Am. J . Botuny 41, 710 (1954). (5) Illingworth, J. W., Keggin, J. F., J . Chem. SOC.1935. 575. (6) Liberti, A., et ai., Ric. Sci. 25, 883 (1955). ( 7 ) Pechard, T., Compt. Rend. 117, 788 (1893). (8) Schriever, K., Toussaint, R., Chern. Ber. 91, 2639 (1958). (9) Shakhova, Z. F., et al., Zh. Aieorg. Khim. 6 , 330 (1961). (10) Souchay, P., Tchakirian, A,, Ann. Chim. [la] 1, 248 (1946). (11) Veitsman, R. M., Zavods. Lab. 26, 927 (1960). (12) Weber, G., LVature 180, 1409 (1957).
RECEIVEDfor review May 14, 1963. Accepted June 26, 1963. Work supported by Eli Lilly and Co., the M. W. Kellogg Co., and the Lubrizol Corp.
Properties and Analytical Applications of the Iron(l1)-2,2’-Bipyrimidine Complex DONALD D. BLY’ with M. G. MELLON Deparfmenf o f Chemisfry, Purdue Universify, Lafayefte, Ind. ,The spectra of several metal ion2,2’-bipyrimidine complexes show that only the iron(ll) and copper(1) complexes have high absorptivities in the visible region. Comparison is made of the iron(ll)-2,2‘-bipyrimidine complex with the analogous 2,2’-bipyridine and 1,lO-phenanthroline complexes. Iron forms more than one complex with 2,2’-bipyrimidine. A Job’s plot indicates that the stoichiometry of the reaction in aqueous solutions is 1 :3 for the complex of maximum color intensity. The data also show that only one pair of the four nitrogens ( 1 , l ’ and 3,3’) enter into coordination with the iron. By following the ultraviolet spectrum with change in pH, the first pK, of 2,2’-bipyrimidine is shown to be 0.6. A stability constant, 83 = K 1 . K 2 . K 3 for the re-
1386
ANALYTICAL CHEMISTRY
+
action 3Bipm Fe+2 S Fe(Bipm)3fZ i s 3.4 X lo’, based on mathematical treatment of Job’s plot data. An absorptimetric method for the determination of iron in the l to 10p.p.m. range employs the 2,2’-bipyrimidine reagent.
T
HE . w . i L Y i w A i L U ~ I L I T Yof
aromatic: compounds containing the XC-C-K linkage in color forming reactions is very familiar. For example, under proper conditions 2,2‘-bipyridine, 1 10-phenanthroline, and sym-tripyridyl-s-triazine react with certain metals to yield soluble colored complexes ( 2 , 4 ) . Thc nnaljtical u e s of t h e v complexes are extensive. -1new type of compound, containing a double S-C-C--S linkage, 2,2’-
bipyrimidine(I), has been prepared recently ( 1 ) . It was desired to determine whether such a compound would form more stable or selective complexes than the others, and whether or not it mas possible to prepare mixed complexey, or aggregate forms of a single complex by utilizing both the 1,l’ and 3,3’ coordination centers. To find ans\yers to thwe and other question3 about the solubility, sensitivity to pH changes, stability, and utility of these complexes, this study was undertaken. 6
1
4
3
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3’
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1 Present address, E. I. du Pont de Nemours and Co., Nylon Technical Division, Wilniington, Del.
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c :2,2‘-BIPYRIDINE
I
30
20
20
30
500
600
700
WAVELENGTH IN MILLIMICRONS Figure 1 . Spectra of metal-2,2’-bipyrimidine complexes in excess 2,2’-bipyrirridine
PART I. PROPERTIES OF
THE IRON(ll)-2,2’BIPYRIMIDINE COMPLEX
Reagents. Solutions of all inorganic reagents w r e prepared from reagent grade cheInicals and distilled water. The standard iron solutions were prepared b y dissolving the appropriate amount of ACS reagent grade iron wire in a small amount of concentrated hydiochloric acid and diluting t o voluine with distilled mater. This gave predominantly iron (11); but wherevc.r iron(I1) was a reactant, an amount of 10% hydroxylamine hydrochloridl. was added to ensure total reducti in. The standard iron(II1) solution mas prepared by taking a 250.0-ml. aliquot of a 50.0p.p.m. standard iron(I1) solution, adding 5 ml. of 30T0 hydrogen peroxide, boiling for 1 hour with evaporation, quantitatively transferring t o a 250-ml. volumetric flask, ccoling, and diluting to volume. The method for preparing the 2,2’bipyrimidine has tieen reported pre\. iously ( I ) . The coi ipound was purified by crystallization from ethyl acetatepetroleum ether bolutions and vacuunidried to a con-tant 1n.p. of 113” to 115” C. A> the compound diswlve. rapidlj in cold water, no other solvent was u5ed. The buffer was a mixture of equal parts of 0.8M sodium acetate and acetic acid. The reducing; agent was 10% hydroxylamine hydrochloride. Apparatus. d Beckman Model G or Zeromatic meter n a s used for pH measurements. Spectrophotometric measurements 111 1 a r t 1 nere made either oil a General EXectrie recording ~pectrophotometcr tth a spectral band \iidth of 10 nip and matched 1.00-cm.
400
500
600
700
WAVELENGTH IN MILLIMICRONS Figure 2. Comparative spectra of iron(1l) complexes of 2,2’-bipyrimidine, 2,2’-bipyridine, and 1,l O-phenanthroline
glass cells, or on a Cary recording sDectrouhotometer. Model 10-11 &.I. 1;ith matched 1.000-em. quartz cells. In Part 11, a Beckman Model DU spectrophotometer with matched 1.000- i 0.005-cm. quartz cells was also used. General Nature of Complexation Figure 1 with 2,2’-Bipyrimidine.
shows t h e visible absorption spectra, taken on the General Electric instrument, of 100 p.p.111. of zinc(II), nickel(II), cobalt(II), chroinium(III), 5 1J.p.m. of copper(I), and 4 p.p.m. of iron(I1) in the presence of a large excess of 2,2’-bipyrimidine a t pH 3 in aqueoua solutions. It is shown below t h a t the reagent is i n its unprotonated form a t this pH. It can be seen t h a t only the iron(I1) and copper(1) complexes have high absorptivities. The coniplexes all absorb strongly in the ultraviolet region, hut the spectra there are obscured by the absorption of the reagent itself. Of the other 30 metal ions studied (li>tecl in thc Interfering Ion. study of Part II), only palladium( II), d y e r (I) , and zirconiuni(I1) formed complexei a t the 100-l1.p.m. levcl with 2,2’-bipyriniidine \t hich could be easily been with the eye. Because of the high sensitivity, stability, and rapid formation of the iron complex, it was chosen for further study to ascertain the nature of the complexing ability of 2,2’-bipyriniidine. Comparison of Iron Complex to Analogous 2,2’-Bipyridine and 1 , l O Phenanthroline Complexes. Figure 2 compare> the >pcctrum of the iron
(II)-2,2’-bipyrimidine complex t o thc spectra of the analogous 2,2’-bipyridine and 1,lO-phenanthrolinc coniplexes a t low iron concentration. Although the spectra of the latter complexes ha\ e been reported elsewhere ( I O , I I ) , the variables of allowable pH’s, buffers, time limits, temperatures, and reducing agents mere made t o orerlap in this case, so that all the bpectra nere generated under as optimum and siniilar conditions as possible. -4, shorn in Figure 2, the s1)ectra are \cry much alike. Howevcr, the iron(II)-2,2’-bipyrimidine system 1 aried from the others in two respects. To obtain maximum color development for this amount of iron, it was necessary to add a 250-fold molar excess of 2,2’-bipyrimidine t o iron(I1). An excess greater than 250-fold produced no further increascl in color intensity. In this respect, 4-fold molar excesses of 2,2’-bipyridine and l,l0-phenanthroline w r e sufficient. Secondly, to obtain similar absorbance values, it was necessary to make the iron(I1) concentration in the 2,2’-bipyrimidnie yolution t n ice that in the other t n o solution.. These differences imply that the iron(II)-2.2’-bipyrimidine coinplex is less stable and has a loaer molar abJorptivity than thrx 2,2’-bipyridine or 1,lO-phenanthrolinc complexcs. The lower stabilitx is 1 erihed in the calculation of BI in thiz paper. The loaer molar absorptivity i-, veil in the value of 5,350 for the 2,2’-bipyrimidine coniplex coinpared to 8,550 and 10,900 for the respective 2,2’-bipyridiiiu arid 1 , l O VOL. 35, NO. IO, SEPTEMBER 1963
1387
I
0.0
0.1
I
0.2
0.3
-M F e / i 450
WAVELENGTH,
500
IN
550
1388
ANALYTICAL CHEMISTRY
1 0.6
0.7
8
Fe+M Bipm
rnp
Figure 3. Plot of absorbance vs. wavelength, on log scale, for indicated ratios of 2,2’-bipyrimidine to iron
phcnanthroline complexes. These a, values were calculated from the maximum A value in Figure 2 for each respective curve. Nature of Iron-2,2‘-Bipyrimidine Reaction. Depending on the ratio of iron(I1) t o 2,2’-bipyrimidine i n solution, more than one complex formed. This was determined in the following way. Known volumes of standard iron(I1) and 2,2’-bipyrimidine solutions were added t o 50-nil. beakers along with 3 ml. of fresh 10% hydroxylamine hydrochloridc reductant; the pH was adjusted to 3 with dilute hydrochloric acid and/or ammonium hydroxide and the solutions subsequently were diluted t o 25.0 ml. The amounts of standard iron and 2,2’-bipyrimidine solutions were taken to make the final concentration of the sum of reagents 3.00 X lO+M, and the ratio of the 2,2’-bipyrimidine to iron(I1) vary from 3 : l to 1:2. The absorption spectra were taken on the Cary instrument. Plots of A (on log scale) us. wavelength are shown in Figure 3, where the log A values were calculated from the experimental curves a t every 10 mp in the range shown. These plots are not superimposable, which indicates that more than one complex forms. None of these complexes could be isolated, however, and their contributions t o the respective curves could not be satisfactorily subtracted out. Stoichiometry of Reaction. The stoichiometry of the reaction which forms the complex of maximum color intensity as a function of 1111 w:~s determined by the method of continuous variations proposed by Job
0.5
Figure 4. Variation of absorbance with mole fraction of iron as function of pH
I 400
1 0.4
(6). It involved plotting 11, at 490 inp. us. the ratio mole Fe(II)/[mole Fe(I1)
+
mole Bipm] at various pH’s, where Bipm stands for 2,2’-bipyrimidine. The data are represented in Figure 4. Fractions of standard iron and 2,2’bipyrimidine solutions were taken so that the total volume used was 10.0 ml. and the final concentration of the sum of reagents in 25.0 ml. was 4.00 X l O - 3 M . After combining the reagents, 3 ml. of 10% hydroxylaniine hydrochloride was added t o ensure total reduction. The pH was adjusted with dilute hydrochloric acid and/or ammonium hydroxide, and the solutions Tv-ere diluted t o volume. hnalysis of Figure 4 shows that the stoichiometry is unaffected by p H in the range 1.5 to 6.0 and that the ratio of iron to 2,2‘bipyrimidine for the complex of maxi-
N
N NN
Figure 5. Hypothetical model of nonexistent polymeric iron(11)-2,2 ’-bipyrimidine complex
mum color intensity is 1:3. It also indicates that the extent of complex formation is not very dependent on pH in the range 3 to 6. The 1:3 ratio of iron(I1) to 2,2’-bipyrimidine is analogous to the 1:3 ratios found for the octahedral iron(I1) complexes of 2,2‘bipyridine and 1,lO-phenanthroline (2). 2,2’-Bipyrimidine Coordinates with Fe(I1) at One Site Only. Since i n
iron(I1) solutions containing excess 2,2’-bipyrimidine t h e respective ratio in the complex is 1:3, it can be shown that the 2,2’-bipyrimidine must be coordinating on one side only, that is, at either the 1,l’ or 3,3‘ poiitions. The argumrnt involvrs proking that hypothetical agglomerate complexes (where the 2,2’-bipyrimidine is coordinating a t both the 1,l’ and 3,3’ positions) do not fit the experimental data. Figure 5 represents a hypothetical, symmetrical complex of arbitrary size. The dots are octahedral Fe(I1) ions bonded to one another by bridges of 2,2’-bipyrimidines (shown schematically) , while the lines represent the direction of bonding. I n the true 3dimensional model the angles would be 120°, but the angles have been distorted in this planar view to prevent lines from hiding one another as they nould a t 120’. For this model, it can be shown that total number of Fe(1I) ions must be given by the formula 3(2%)- 2,regardless of the size of the symmetrical complex. The number of internal irons (all except the external shell of irons) is given by 3(2”-’) - 2, while the number of irons in the external shell is given by 3(2n-1). The number n is any integer >= 0, and repremits thr iiuiiher of repetltirig distances from the center t o the outer edge of thc complex. The t o t d iiuiiibclr of
2,2’-bipyrimidines is calculated from the internal and exteinal irons as follows. For any value of n 1. Each interrral irvn bonds witli 3 X 1/2 Bipins (2,2’-bipyrimidine is abbreviated Bipm to prevent confusion with the numbers), so that the number of such Ripma is 3,/2 X 3(2“ -l) - 3//z X 2, or 9/2(2n-1) - 3. 2. Each external iron bonds with 5/2 Bipms, since the reagent is present in excess, so that the number of such Bipms is X or 15/2(2n-1). Thus, the total lumber of irons is (i(2r1-1)- 2, and the total number of Bipms is 12(2n-l) - 3. It can be been t h a t for large n the limit ratio of irons to Bipms is 1:2; in fact, even for as low as n = 3, the ratio is 22:45. Likewise, it can be shown that if a dimer is assumed, where only one 2,2‘bipyrimidine mole1:ule coordinates a t both the 1,l’ and 3,3’ sites and the rest a t one site only, t k e ratio limit is 2 : 5, or 1: 2.5. But the experimental value in solutions of excess 2,2’-bipyrimidine is 1:3. It can be sc’en from Figure 4, where a 1 : 3 ratio of reagents corresponds to a 0.250-mole fraction for iron and a 1:2.5 ratio of reagents corresponds to a 0.286-mole fraction, that the difference between a 1 : 3 ratio and a 1:2.5 ratio is certainly greater than experimental error. Thus, in solutions of excess 2,2’-bipyrimidine, the ratio cannot be 1:2.5. The large agglomerate model and the dimer model, therefore, must not represent the actual case. Other models, such as unsymmetrical ones, or ones where not all of the coordination sites of t k e iron are taken up, serve only to lower the limit ratio of iron to 2,2’-hipyrimidine, rather than to increase it towards the observed 1:3 ratio. The only model left LS that of the simple species where three 2,2’-bipyrimidines are complexed to one iron, and each 2,2’-bipyrimidine coordinates a t one nitrogen pair. Only this model has a theoretical limit ratio of 1:3. Thus, the structure of the complex is analogous to
that of the 2,2’-bipyridinc and ],lophenanthroline complexes. The structure and stoichiometry of the copper(1)-2,2’-bipyrimidine complex were not investigated, but they should be interesting. When the two reagents are present in equal molar quantities, or there is a slight excess of copper(1), a very insoluble precipitate forms, even in strong acid, which contains copper and 2,2’-bipyrimidine. This might imply a polymeric or aggregated species. Stability Constants and pK, Values. It was desired to determine the stability constants for the iron-2,2’bipyrimidine complexes, or a t least for the one of maximum color intensity. But, since none of these could be separated from solution, and since they obviously are in equilibrium, a value B3 was determined. Assuming that three complexes form of the type shown in (I), then Ba = K1.K~.K3, and describes the equilibrium shown by (2). 3Bipm Fe+2F?
+
Ki
+ Fe(Bipm)+2F? Kr Bipm + Fe(Bipm)2+2e Ka
2Bipm
Fe(Bipm)l + 2 3Bipm
+ Fe S Fe(Biprn)S’+
(1) (2)
Similar work has been done with iron(I1) complexes of 2,2’-bipyridine (2, 5, 7 ) and 1,lO-phenanthroline (9, 8, 9).
The equilibrium postulated in (2) assumes that acid does not enter into the reaction. It was necessary to prove this before proceeding with the determination of B3. A single pK, of 0.6 for 2,2’-bipyrimidine was found by following the change in absorbance of a 5.00 X 1O-bM solution with pH a t 265 mp (see Figure 6). For comparison, the first pK. of pyrimidine is 1.3 and of 2chloropyrimidine is 0.8 (S), while those for 2,2’-bipyridine and for 1,lO-phenan-
o
throline are 4.3 and 5.2. rcsprrtively (2J8 ) . For determining pK,, the pH of the solutioiia \\‘as adjusted with dilute itinmonium hydroxide and/or dilute hydrochloric acid. The pH was read after the photometric readings, which were taken on the Cary instrument. I n the case of the strongly acidic solutions, between pH -1 and 0, known amounts of constant-boiling hydrochloric acid, 20.263y0 and d = 1.1019, mere added by pipet, and the pH was calculated. Corrections for activity were not made. Also, it had previously been determined that the interaction of 2,2’-bipyrimidine with chloride was negligible. Thus, a t pII 3 the first PIC, of 0.6 shows that one need not consider acid in equilibrium (2). The calculation of B3 was made with data from two Job’s plots. These were obtained by plotting absorbance us. mole fraction of iron(II), a t 490 mp, a t pH 3.0, and a t constant chloride concentration of 0.2M. The chloride arises from the addition of 3 ml. of 10% hydroxylamine hydrochloride and enough sodium chloride to make 0.2X. This concentration of chloride produced no visible change in the free iron concentration in equilibrium (2). Known amounts of a standard iron solution were taken by pipet, 3 ml. of fresh 10% hydroxylamine hydrochloride was added, and then the required amount of sodium chloride, followed by the proper amount of standard 2,2’-bipyrimidine solution. The pH was adjusted to 3.0, the solutions were diluted to 25.0 ml., and the photometric readings taken. The data are shown in Figure 7 . The upper curve represents a molar sum of reagents of 3.50 X lop3, while the lower curve represents a molar sum of reagents of 3.00 x 10-3. The calculation follows. An absorbance value (Figure 7 ) was chosen common to both curves, arbitrarily a t A = 0.25 ==I 0.01. At this point the
’
6
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05
04
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0
03
m 02 01
00 -1
00651
2
3
4
5
6
7
8
DH Figure 6. Variatior of absorbance with p H for 5.00 X 10-5M 2,2’-bipyrimilJine at 265 mp
a
-M Fe/M
9
Fe +M Bipm
Figure 7. Variation of absorbance with mole fraction of iron a t two concentration levels for calculation
of 63 VOL. 35, NO. 10, SEPTEMBER 1963
1389
same amount of complex will have formed in both cases, but the concentration of reagents is different between the two curves. S o w let
physical meaning. This would mean that the amount of complex formed would be greater than the limiting reagent, iron. However, the value q = 9.7 X does have phpical meaning. It indicates that about 2.8% of the iron is in the complex form, or forms which are giving rise to the measured absorbance. The choosing of other A values gives varying al, a2, and q values, but B3 remains constant. Using q = (9.7 i 0.1) X to solve for B3, the value i q B3 = (3.4 + 0.5) X IO’, or log B3 = 7.5 i 0.07. When ahorbance values of 0.200, 0.300, and 0.325 were chosen, the calculations for BS all yielded values agreeing with this value within the estimated error. The values of log B3 for the iron(I1) complexes of 2,2’-bipyridine and of 1,lO-phenanthroline have been reported t o be approximately 16 and 21, respectively ( 2 ) . Obviously, the 2,2‘bipyrimidine complex is much weaker than the other two. Thiq result was expected, as discussed above in the analysis of Figure 2.
CZ = total concentration = 3.50 X 10-3JI in case I1
C1
=
total concentration
=
3.00 X 10-3M in case I
Then in case I, for the mole fraction values corresponding to the chosen A (Figure 7 ) , let a, =
the amount of Fe initially presenti.e., the mole fraction x Ci = 0.10 x 3.0 x 10-3 0.01) x 10-~AE
=
(0.30 i=
q = the amount of complex formed al - p = the amount of Fe left over
C1
- a1
a3 =
=
after complex formation the amount of 2,2’-bipyrimidine initially present, which also is its mole fraction x
Ci = (2.70 f 0.01) X 10-3~ the amount of Fe initially presenti.e., the mole fraction x C2 = 0.057.5 X 3.50 X = (0.201 =!= 0.01) X 10-3N and in case 11, a?
Then, (3)
PART II. APPLICATION OF THE IRON[II)2,2’-BIPYRIMIDINE COMPLEX TO THE DETERMINATION OF IRON
An analogous series is set up involving C2 and az. To solve for q, advantage is taken of the fact that d was chosen the same in both cases I and 11, so that the same amount of complex has formed in both cases. Then, (4) holds. (a,- a ) (Cl
As shown in Figure 2, 4 p.p.m. of Fe(I1) in the presence of greater than 250-fold excess of 2,2’-bipyrimidine gives a n absorbance of 0.34 a t 490 mp. This high sensitivity for the complex suggested that 2,2’-bipyrimidine could be used for the spectrophotometric determination of iron. Sensitive spectrophotometric methods for iron are already known, using for example, 2,2’-bipyridine or 1,lO-phenanthroline (10, 11). However, experience with the Fe(II)-2,2’-bipyrimidine system showed that this method would have several advantages in addition to high sensitivity. The reagent dissolves rapidly in cold water. The colored complex
- - 3q)3 =
(a2
a1
- q ) (CZ - a2 - 3913
(4)
Upon putting in the values for the C’s and a’s and graphically solving the cubic equation, three real roots are obtained, two identical a t q = 0.99 X lo-% and one a t q = 9.7 X loF5. Since q is defined as the amount of complex formed, and since there was only 0.35 X 10-3Jf of iron initially present, the value p = 0.99 X can have no Table I.
Effect of 2,2’-Bipyrimidine Concentration and pH on Color Development
Excess Bipm ,If ratio
6.0
10 50 100 150 200 2.50 300 350 10 50 100 150 200 250 300 350
1390
Absorbance at, pH 4.6
3.0
0.031 0.219 0.219 0,270 0.296 0,322 0.330 0.331
4 p.p.ni. Fe 0.028 0.107 0.216 0.277 0.316 0.325 0,327 0.328
0.026 0.114 0.222 0.284 0.322 0.335 0.345 0.348
0.015 0.087 0.194 0.258 0.290 0.309 0.328 0.333
0.073 0.411 0.604 0.650 0.688 0.688 0,690 0.688
8 p.p.m. Fe 0.068 0.394 0,569 0.643 0.688 0.695 0.700 0.701
0 . OGG 0.400 0.602 0.668 0.690 0.690 0.692 0,693
0,042 0.410 0.594 0.648 0.679 0.688 0.688 0.688
ANALYTICAL CHEMISTRY
5.2
forms immediately a t room temperature, and the absorbance does not change for months. The extent of complex formation is insensitive to pH in the range 2 to 6. And finally, several interferences arc different from those in the other cases. The greatest disadvantage is that the reagent is not yet commercially available. Apparatus and Reagents. T h e same apparatus was used as in Part I, and the reagent solutions were prepared in the same way. However, only one standard stock iron solution and one standard stock 2,2’-bipyrimidinc solution m s used throughout the work i n P a r t 11. The iron solution (50.00 p.p.m. Fe) was prepared by dissolving 0.0500 gram of ACS reagent grade iron wire in a small amount of concentrated hydrochloric acid and diluting t o 1000.0 ml. The 2,2’bipyrimidine solution (2.0% w./v.) mas prepared by dissolving 5.0 grams of the purified compound in distilled water, filtering, and diluting t o 250.0 ml. Twenty-five-milliliter flasks were used throughout this study to keep to a minimum the total amount of 2,2‘bipyrimidine used. EXPERIMENTAL WORK
Effect of Reagent Concentration a s Function of pH. As discussed previously, a molar ratio of reagent t o iron greater t h a n 250-fold is necessary for maximum color development a t low concentrations of iron. This conclusion was confirmed by a systematic study of the excess of reagent required, as a function of pH. for 4 and 8 p.p.ni. of iron. The procedure follows. ,211 required amounts of iron, 2,2‘bipyrimidine, and hydroxylamine hydrochloride were added from burets to 50-ml. beakers. The solutions were diluted to about 15 ml. and the p H n-as adjusted with dilute hydrochloric acid and/or ammonium hydroxide. After transferring the solutions to 25-ml. volumetric flasks and diluting t o volume, the absorbances were read a t 490 mp z‘s. a distilled water blank. Table I gives the data. The data show that, in all cases except two, the absorbances between the 250and 350-fold excesses of reagent agree within the =t3 u limits (discussed under Interfering Ions) set on the system. The fact that A does not increase with reagent concentration beyond a certain value means that a calibration curve for iron can be constructed. The data in Table I also show that the excess necewary for maximum color development is not dependent on pH in the range 3.0 to 6.0. Effect of pH. T o find the pH range of mavimum sensitivity for the complex, it was necessary t o vary the pH in the solutions from 0.5 t o 6.5.
Table 11.
Effect of p H on Absorbance
True pH
A
0.72
0.228
i.oS
1.70 2.24 2.70 3.22 3.54 3.93 4.42 .5.00 -5.40 5.92 6.50
0.289 0.301 0.303 0.303 0.306 0.311 0.311 0.315 0.316 0.316 0.314 0.314
Below p H 0.5 the color faded rapidly, while above 6.5 the iron precipitated from solution. T h e procedure was t o pipet 2.0 ml. of stock iron solution, 3.5 ml. of stock !2,2’-bipyrimidine solution, and 3.0 ml. of 10% hydroxylamine hydrochloride into 50-ml. beakers and adjust the pH with dilute hydrorhloric acid and/or ammonium hydroxide. The solutions were transferred to 25-ml. volumetric flasks, diluted to volume, and the absorbances read a t 490 mp. Aftei- the photometric readings, the pH was again determined for each solution. These are recorded in Table 11. The absorbance vaIues were corrected for the reagent blank of 0.024. r s i n g the value of 0.306 a t pH 3.22 as a reference, all of the absorbances between pH 1.70 and 6.5 agree within A 3 u of the 0.306 value. There is a slight trend, however, towards higher absorbances as the p d increases. I n this range the slope of ,he line AA /ApH is 0.004, or an O.llyochange per p H unit. Thiq effect i? probably due to varying hydration of the iron rather than to any effect on the reagent. Effect of Time anti Temperature. I n the presence of sufficient reducing agent, excess 2,2’-bipyrimidine, and p H range 2 to 6, the color forms immediately and the absorbance does not change for months. T h e color is bleachc d considerably a t the boiling temperature, but returns to the initial value again upon cooling t o room temperature. The abqorbance is unaffected bv a change of &loo t o 25’ C. Reducing Agents. The color reaction is with ferrous iron only, b u t ferric or total iron ma!’ be determined by using a suitable reducing agent. -\scorbic acid, sodium sulfite, sulfurous acid, hydrazine sulfate, and hydroxylamine salts wwe all effective. Hydrovylamine hydrochloride was selected for this work as i t was easy t o work with and introduced only chloride ion into the solutions. It
required slightly greater than 0.5 ml. of 10% hydroxylamine hydrochloride to reduce completely 4 p.p.m. of iron(II1) in 25 ml. of solution. The normal amount used in this stiidy \\’as 3 ml. Effect of Buffer. Because the effect of interfering ions might vary with p H , i t was necessary t o buffer the solutions. As all buffers studied reduced somewhat the absolute absorbances, i t was deemed best t o work only with one buffer. A sodium acetateacetic acid buffer interfered only slightly and the least of any buffer studied. It was prepared 0 . S N in sodium acetate and acetic acid and 3 t o 4 ml. was the usual quantity used. Effect of Iron Concentration. The buffered system followed Beer’s law for 1 to 10 p.p.m. of iron in the presence of enough 2,2’-bipyrimidine t o be 250-fold molar excess for the 10 p.p.m. A typical calibration curve was prepared as follows. Appropriate amounts of standard iron solution were pipetted into 25-nil. volumetric flasks. Then were added 4 ml. of sodium acetate-acetic acid buffer, 3 ml. of 10% hydroxylamine hydrochloride reductant, and 8 ml. of 2.0570 2,2’-bipyriniidine solution. The pH was adjusted with dilute hydrochloric acid and/or ammonium 0.5. After dilution hydroxide to 4.5 to volume, the solutions were mixed, and the absorbances read a t 490 mp. Corrections mere made for reagent blanks, even though they were small. Even though the reagent mas in exceis by 2500 molar equivalents for 1 pap.m. of iron and by 250 molar equivalents for 10 p.p.m., the system does follow Beer’s law. But one must keep in mind that a t all times the minimum excess required is 250-fold. Interfering Ions. T o determine whether various diverse ions would
*
Table 111.
Ion
Added as
AsOdW3
interfere, confidence limits were established for the calibration procedure described. On 5 consecutive days one tleterniination was made per day for 4 p.11.m. of iron following the adopted proceduie. For these 5 determinations thc mean absorbance minus the 111~aiiblank was 0.319, the range was 0.010, and the standard deviation was A0.00332. The 3 u limit was taken as 0.010. Then, if the absorbance of any sample containing 4 p.p.m. of iron and a diverse ion had an absorbance of 0.010 grentpr or less than 0.319, this ion waq wid to interfere. The tolerable concentration was calculated on the assumption that Beer’s law was valid for the interfering ion. Blanks were always determined. The diverse ions rrere added to give a final concentration of 400 p.p.m. If an ion interfered seriously a t this concentration, it was again studied a t a lower concentration, and the calculation made. 111 solutions were prepared exactly as in the ralibration procedure, except that the diverse ion was added to the 25-ml. volumetric flasks just after the iron solution. Metal ions were added as the chloride, nitrate, or sulfate, n-hile the anions were added as sodium, potassium, or ammonium salts. The following ions did not interfere : Al+3, B+3, Ba+2, Bi+3, Br-, Br03-, Ca+2, C1-, C103-, Clod-, Cof2, Kf, H ~ TI-, ~ ,IO4-, Mg+2, hI00~-~, S a + , Ydf3, KO2-, SO3-, SCS-, SOa-’, S203-2, Sr+2, tartrate, Th+4, U02+z,and TrOa-. Precipitates , were formed by Ag+, Bi+3,H F + ~Pd+2, SeOdFZ,Snt2, S203-z, Ti+3, and JT+6. I n the cases of Bi+3, Hg+2,8203-2,and Ti+3, precipitates formed which could by filtered off and the absorbance was read without residual interference. The ions listed in Table I11 can be tolerated in the concentrations indicated.
Effects of Interfering Ions
Concentrations nrlded (p.p.m.) 100 50
Cz04-2
Citrate
400
Cd +2 CY -
c u +’
400 50 10
Cr + 3
100 400
F-
Ni +2
Pb + 2 Pd + 2 Pt + 2
Sb +3 Si03-2 Zn + 2 ZrO + 2
XiCL Pb(iYOs)? PdC12 in IlCl KHgPOd PtCI2 !n HC1 SbC4 in HC1 NazSiO3
ZnSOd
ZrO( NOs)n
100 400
50 50 400 400
100 100 50
>Tax. perniissihle Concn. (p.p.m.) 18 50 180 166 8
