摘要:

Analyst, October, 1979, Vol. 104, pp. 961-972 96 1 Calcium Ion-selective Electrode Measurements in the Presence of Complexing Ligands A. Craggs, G. J. Moody and J. D. R. Thomas Chemistry Department, Ufliversity of Wales Institute of Science and Technology, Cardiff, CFl 3N U Calcium ion-selective electrodes based on calcium bis[di(4-octylphenyl)- phosphate] sensor and dioctyl phenylphosphonate solvent mediator have been used for monitoring free calcium-ion levels (from below lo-' to above 10-3 M) in the presence of citrate, malate, malonate, oxalate, EDTA, NTA, sulphate, orthophosphate, tripolyphosphate and pyrophosphate anion ligand systems under conditions of constant ionic strength maintained by sodium chloride. Log data fall in the range of those previously measured for the various equilibria by alternative methods, thus demonstrating that calcium ion-selective electrodes of the type used here can be employed for free calcium- ion measurements to below the detection limits characteristic of calibrations with serial dilution standards and without disturbing the equilibria of com- plexation.Equilibria existing in the tripolyphosphate and pyrophosphate systems were discerned by application of the MINIQUAD program for computing forma- tion constants and species distribution. Predictions concerning the existence of [CaP30,J3- and [Ca(P,O,,) 2]8- in tripolyphosphate and of [CaP20,I2- and [Ca(P,O,),]G- in pyrophosphate systems are briefly discussed. These calcium ion-selective electrodes are not affected by added phosphate except insofar as free calcium-ion levels are lowered by complexation.Keywords : Calcium ion-selective electrodes ; fovmation constants ; conzplexation equilibria Complexing ligands frequently lower the level of free calcium ions to well below the serial dilution calibration range of calcium ion-selective electrodes. However, calcium ion- selective electrodes can be calibrated to much lower calcium-ion levels with complexing ligands by exploiting this phenomenon where the stability constants for the calcium - ligand complexes are known, such as with EDTA and nitrilotriacetic acid (NTA) ligand~.l-~ Such calibrations suggest that calcium ion-selective electrodes may be used for calcium-ion measurements down to ca. 10-8 M. One way of assessing the utility of calcium ion-selective electrodes for measuring calcium ions in the presence of complexing ligands is to examine the characteristics of the associated equilibria.The procedures are similar to previous applications of ion-selective electrodes for determining stability constants for calcium - carboxylate systems at 0.1 M ionic strength sodium perchlorate6 and of association constants of ADP and ATP complexes at zero ionic The stability data so obtained can then be compared with recorded data obtained by other means and hence used to deduce the suitability of calcium ion-selective electrodes for the measurement of free calcium ions under such conditions. Related to the use of calcium ion-selective electrodes for studying complexing equilibria are determinations of chelating properties of detergent builders.9 These have been based on electrodes using the Orion 92-20-02 calcium liquid ion exchanger, which normally has a relatively short calibration range compared with the newer electrodes based on calcium bis [di (4-octylphenyl)phosphate] sensors.This paper extends the above studies and describes the scope of calcium ion-selective electrodes based on calcium bis [di(4-octylphenyl)phosphate] sensor in the determination of stability constants of calcium with oxalate, citrate, malate, malonate, EDTA, NTA, ortho- phosphate, pyrophosphate, tripolyphosphate and sulphate complexing ligands. Experimental Procedures PVC matrix-membrane calcium ion-selective electrodes were fabricated as described962 CRAGGS, MOODY AND THOMAS CALCIUM ION-SELECTIVE ELECTRODE Analyst, Val. 104 previously.l0Y1' The master membranes were based on 0.36 g of dioctyl phenylphosphonate solvent mediator (Alfred Bader Library of Rare Chemicals, Division of Aldrich Chemical Co.Inc., Milwaukee, Wisc., USA, Cat. No. S57593-3) plus 0.04 g of synthesised12 calcium bis [di(4-octylphenyl)phosphate] sensor in 0.17 g of PVC. A titration vessel was assembled as in Fig. 1 with Radiometer, Type pHM 64, research pH - millivoltmeters being used for measuring pH and e.m.f.s of the calcium ion-selective electrode systems with respect to a Radiometer, Type K401, saturated calomel electrode, respectively. E F D Fig. 1 . Titration vessel assembly: A, millivoltmeter; B, pH meter; C, nitrogen inlet; D, inlet from micro- meter syringe or burette; E, reference electrode ; F, poly(viny1 chloride) matrix membrane calcium ion-selective electrode; and G, combination pH elect rode .Because of possible interference from atmospheric carbon dioxide, all titrations were performed with nitrogen bubbling slowly through the solution conveniently maintained at constant ionic strength ( I ) by sodium chloride because electrodes based on the above calcium ion exchanger are not subject to interference by up to 0.15 M sodium chloride. The perchlorate interference with calcium ion-selective electrodes13 precludes its use for adjust- ment of ionic strength. The calcium ion-selective and reference electrode pairs were calibrated before and after each experiment at the chosen temperature and ionic strength using serially diluted calcium chloride standards.At low calcium-ion levels, the calibration was carried out by spiking from an Agla syringe small aliquots of calcium chloride standards to the sodium chloride solution.14 Calibrations of diluted standards have been shown to extrapolate to lower levels of calcium ions in the presence of complexing ligands, such as EDTA and NTA.5 Electrodes that drifted more than 0.1 mV d-l were discarded. Citrate, Malate and Malonate to 5 x 10-3 M for citrate and malate and 5 x M for malonate and at I = 0.10 M maintained by sodium chloride, were titrated at 25 "C with 0.1-l-cm3 aliquots of 0.10 M sodium citrate, malate or malonate solution containing an equivalent amount of calcium. The pH was adjusted to between 6 and 7 with sodium hydroxide or hydrochloric acid as appropriate.Volumes of 100 cm3 of calcium chloride solutions, usually 5 x toOctober, 197.9 MEASUREMENTS IN THE PRESENCE OF COMPLEXING LIGANDS 963 Oxalate The difficulty of measuring stability constants of systems involving sparingly soluble saljs has been discussed by Bond and Hefter.15 The solubility for calcium oxalate, vix., 2.29 x 10-9 (I = 0) and 1.58 x ( I = 0.10), respectively,16 indicates the possible difficulty with this system, as emphasised by the relatively sparse amount of information available.17 Nevertheless, the calcium oxalate system was included in this study at 25 "C for 1 = 0.05, 0.10 and 0.15 M (maintained with sodium chloride) and the pH adjusted to approximately 9 with sodium hydroxide. M sodium oxalate at I = 0.05 M was titrated with calcium chloride solution, such that the total calcium-ion concentra- tion varied between 2 x loe5 and 5 x M.Slightly different concentrations were used for titrations at I = 0.10 and 0.15 M. Although the reported solubility product was exceeded, no precipitation was observed in the reaction vessel and the final electrode potentials were steady (to about 0.1 mV) over several days. Thus, for example, a solution of Sulphate The calcium sulphate ion pair (Ca, SO,) is fairly weak1' with a logarithmic formation constant in the range 2-2.5. Hence, unless precipitation occurs, the change in free calcium- ion concentration on adding sulphate to a calcium-ion solution will be relatively small. Nevertheless, the electrodes were capable of discerning the small changes in e.m.f.associated with the small calcium-ion activity changes. The procedure described above for citrate, malate and malonate was used, the initial calcium-ion concentrabion being in the range 5 x to 2.5 x M at I = 0.10 M. Maintenance of pH at 8.5 avoided corrections for HSO,- formation. EDTA and NTA These ligands form strong complexes with calcium ions and their presence tests the ability of the calcium-ion electrode reliably to detect low levels of free calcium ions. The procedure used for citrate, malate and malonate at I = 0.10 M was again convenient, the initial calcium- ion concentration being 7.5 x loF5 to loF3 M for EDTA and 7.5 x to lo-* M for NTA. The pH of the EDTA system was maintained at between 5 and 6, this region being chosen to maintain the free calcium-ion level sufficiently high (ca.M) for measurement. Although the predominant ligand species, L, at this pH is H2L2-, complexation occurs according to Ca2+ + H2L2- + CaL2- + 2H-t and can therefore be regarded as being dependent on the L4- concentration. The calcium complex formation constant is significantly lower in the NTA system than with EDTA, thus permitting the use of a sufficiently high pH to ionise the ligand fully, namely 10.5. Measurements at pH < 9 involve protonated species. Tripolyphosphate and Pyrophosphate Although the 1 : 1 complexes of calcium with the condensed phosphates are undoubtedly the most important,l* it is necessary to consider the possible formation of several other species in such systems, for example, protonated complexes, especially for tripolyphosphate and pyrophosphate.The possible formation of other species required the use of available computer programs for computing formation constants and species distribution in equilibrium s y ~ t e m s . l ~ - ~ ~ Of these, MINI QUAD^^ was selected for its capability of dealing with potentiometric titration data.24-26 The modifications required to adapt the program for the present application are discussed below. The pH of the titration test solutions used with tripolyphosphate and pyrophosphate systems is such that the possibility of protonated ligand species had to be taken into account, the relevant equilibria being964 CRAGGS, MOODY AND THOMAS CALCIUM ION-SELECTIVE ELECTRODE Analyst, VOZ. 104 for tripolyphosphate and H2P,072- + 2H+ + P2074- for pyrophosphate.As the MINIQUAD program is sensitive to small changes in the constants for these e q ~ i l i b r i a , ~ ~ the range of published constants necessitated determinations under the conditions of measurements with calcium by straightforward acid - base titration. Thus, typically, sufficient 1 M hydrochloric acid (ca. 0.2 cm3) was added to 100 cm3 of 1 0 - 3 ~ of the appropriate sodium phosphate solution to adjust the pH to about 4, and the ionic strength was adjusted to 0.10 M with sodium chloride. Sodium hydroxide solution (0.1 M) was added from an Agla burette until a pH of 9.5-10 was reached and the monitored pH obtained during the addition used as input to MINIQUAD together with titrant volume. For the calcium ion-selective electrode measurements with both the tripolyphosphate and pyrophosphate systems, 100-cm3 calcium chloride solutions (5 x to 5 x 1 0 - 4 ~ and I = 0.10 M) were titrated at 25 "C with freshly prepared 0.10 M sodium tripolyphosphate solution from an Agla burette to a final tripolyphosphate concentration of 5 x M. Prior to the titration the pH of each solution was adjusted to between 8 and 9. After each addition, the volume added, calcium ion-selective electrode potential and pH were recorded for use as MINIQUAD input data.Orthophosphate Although the most important equilibrium involving calcium and orthophosphate is that yielding a precipitate of Ca,(PO,),, there is one involving a soluble complex, namely CaHPO,, but which is considerably weaker than the complexes of the chain Hence, calculation of the formation constant for this did not require the use of the MINIQUAD program.The experimental procedure for citrate, malate and malonate was used, the initial calcium ion concentration being 10-4 to 2.5 x The titrant was 0.10 M disodium hydrogen phosphate in the appropriate calcium ion-containing solution. The pH was in the range 9.0-9.5. M at I = 0.05, 0.10 and 0.15 M. Calculation Procedures Citmte, malate, malonate, EDTA and NTA Under the experimental conditions employed, the tendency of citrate, malate and malonate to form 1 : 2 complexes or protonated species can be neglected and the stability constant, K , for the reaction between calcium ions and carboxylate ligand, An-, K = [CaA2-n]/[Ca2+] [An-] . . . . .. * ' (1) was calculated from the experimentally measured Ca2+ concentrations with the other quantities being deduced from the initial concentrations of Ca2+ and ligand acid, [H,A]total, and the H+ concentration using the relation6 to calculate [An-].As pH measurements give the activity of the hydrogen ion, the calcula- tions could require conversion of aH to [H+]; hence at the ionic strength studied of 0.10 M an activity coefficient of 0.76 was used. The various dissociation constants taken for the ligand acids are summarised in Table I as pK, data. Equations (1) and (2) were also employed for calculating the formation constants of the 1 : 1 complexes of EDTA and NTA with Ca2+.October, 1979 MEASUREMENTS IN THE PRESENCE OF COMPLEXING LIGANDS TABLE I 965 pK, DATA FOR LIGAND ACIDS Ligand PKI PK, PK, pK, Reference Citrate .. . . . . 2.95 4.42 5.74 6 Malate . . . . . . 3.22 4,57 6 Malonate . . .. . . 2.61 5.27 29 EDTA . . . . . . 2.07 2.75 6.24 10.34 16 NTA . . .. . . 1.97 2.57 9.81 16 Orthophosphate . . . . 2.0 6.9 11.7 16 Oxalate Use of high pH (around 9), so as to ignore the protonated ligand species, and a large excess of ligand to metal concentration, in order to equate the free ligand concentration to the total ligand concentration, simplified the calculations for the oxalate system by the following equation : * * (3) K = [CaOx]/ [Ca2+] [OX"] total . . . . Sulphate At the high pH of 8.5 used no corrections were made for the formation of HS04-, or for undissociated sulphate (as NaS0,-). Hence the formation constant data obtained for calcium sulphate ion pair formation ( K = [CaSO,]/ [Ca2+]free [S0,2-]free) were apparent or conditional, valid only for the conditions under which they were obtained.Orthophosphate CaHPO, is the assumed formula of the soluble complex formed in the calcium - ortho- phosphate system,28 with respect to which the mass-balance equation for the total phosphate can be written as + ( [H+]210(P"~ + . . .. - - (4) It can be deduced from equation (4) that at pH 9-9.5 less than 1% of the orthophosphate present is not in the form HP0,2-, apart from that bound to calcium. Hence the approxi- mation can be made and, as [CaHPO,] = [Ca2+]tot,l - [Ca2+]free . . .. - - (6) the formation constant of CaHPO,, was readily calculated. Tripolyphosphate and pyrophosfihate complexes being con~idered.~~~23 while initial estimates of the unknown constants are subsequently refined.MINIQUAD requires input of formation constants and stoicheiometric coefficients of the In this, known formation constants are held constant The free con-966 CRAGGS, MOODY AND THOMAS : CALCIUM ION-SELECTIVE ELECTRODE Analyst, VoZ. 104 centrations are those determined by potentiometric titration data, either by means of the Nernst equation if e.m.f. values are read in, or directly from the decimal antilogarithm as for pH. The unknown free concentrations are calculated for each point by iteration on the relevant mass-balance equations. The output for MINIQUAD consists of computed values of the formation constants with the standard deviation and correlation coefficients, and failure messages are printed if the refinement is unsuccessful.The residuals on the mass-balance equations are then subjected to a statistical analysis for any systematic errors in the input data or the chemical model. Finally, concentrations of all species present may be calculated as functions of a given reactant concentration and printed out in graphical form. Before its implementation in this study, some alterations were made to the original version of MINI QUAD,^^ firstly in respect of the error noted by Leggett30 in subroutine ML, and secondly to allow for the fact that practical ion-selective electrodes do not always obey the ideal Nernst equation. To correct the error in subroutine ML, the section of code22 126 DT(1) = CX(1) - TT(1) to 128 DT(1) = DT(1) + JQR(1, J) *CI( J) was changed to 126 DT(1) = CX(1) - TT(1) DO 1128 J = 1, NK W = HLNB(J) DO 127 I = 1, NMBE CI(J) = EXP(W) DO 128 I = 1, NC 127 W = W + HX(1) * JQR(1, J) 128DT(I) = DT(1) + JQR(1, J) *CI( J) Secondly, in subroutine DINP, the code 116 CX(NCPL) = EXP((EMF(L) - EZERO(L)) * JEL(L)*11.6049/(TEMP + 273.16)) was changed to 116 CX(NCPL) = EXP((EMF(L) - EZERO(L)) *2.303/FN) where FN is the actual slope obtained from the calcium ion-selective electrode calibration.Solution pH values, measured with a glass electrode, were input directly into the program, so that the computed formation constants were "mixed" in that the activity of the hydrogen ion has been used in the calculations while all other species were in terms of concentration.TABLE I1 AT 25 "C (I = 0.10 M) PROTON ASSOCIATION CONSTANTS OF PYRO- AND TRIPOLYPHOSPHATE ANIONS Computed values (this work) No. of Literature values" titration I L 3 Ligand points pl (s.d.)* p2 (s.d.)* Log p1 Log Pa Log Log pp Medium Pyrophosphate . . . . 74 1.12 X 10' 6.95 X 1013 8.05 13.84 9.11 15.47 (CH,),NCl (0.12 x 10') (1.35 x lo1,) 8.3 14.3 NaNO, 9.00 15.23 (CH,),NCl 8.93 16.05 (CH8),NCl Tripolyphosphate . . 90 6.99 X lo' 2.51 X 10'' 7.84 13.40 8.81 14.64 (CH,),NCI (0.31 x 107) (0.20 x 1018) 8.65 14.40 (CH,),NBr 7.9 13.5 NaNO, 8.06 13.49 KCI -~ - 8.74 14.76 (CH,),NCl 8.82 14.75 (CH,),NNO, 8.81 14.60 (CH,),NCl * Standard deviations in parentheses.October, 1979 MEASUREMENTS IN THE PRESENCE OF COMPLEXING LIGANDS 967 As mentioned above, the monitored pH obtained during the titration of solutions in hydrochloric acid of sodium tripolyphosphate or sodium pyrophosphatc with sodium hydroxide was input into MINIQUAD in order to determine thz proton associ- ation constants of the parent acid anions.Also used were values for pK, (taken as 14) and initial estimates of the overall proton association constants of the tripolyphosphoric and pyrophosphoric acids, taken as lo8 and lo1* for p1 and p2, respectively, for both of the acids. The refined values of /3, and /I2, shown in Table 11, are compared with literature values in related systems. The slightly lower p1 and p2 in media containing sodium and potassium ions compared with those in tetraalkylammonium ions as supporting electrolyte are indicative of weak complexing of the alkali metal ions by these phosphate ligands.Such complexing cannot be incorporated in the calculations so that the computed data are conditional and apply only to solutions at I = 0.10 M. The program used for calculating the p data was also required to plot the percentage of each species, relative to the total ligand concentration at each point of the titration curve. This is illustrated in Fig. 2 for the pyrophosphate ligand titration with 53 points, at the end of which the free P,O,4- ligand is present. Proton association constants. Percentage 0 50 100 t t + + t + t + + + + t + t t + + t t t t t t t + t + + + + t + t + + t + + + + + + t + t + + + + + t + + + t + + + 1 + A B 2 + A B 3 t A 6 4 + A B 5 + A B 6 + A B 7 + A B 8 t A 6 9 t A B 10 t A B 1 1 t B 12 t A €3 13 t A B 14 t A B 15 + A B 16 t 1 7 + 18 t B A 19 t 8 A 20 + 8 A 21 + B A A B AB 22 + 8 23 + .r 24 + 0 25 + a 2 6 t c 27 t 0 28 t '3 29 t m 3 0 t B b 3 1 t B F 32 +B 33 +B B B B B B El B B A A A A A A A A A A A A 34 t B A 35 + A 36 + A 37 t A 38 t A 39 t A 40 + A 41 t A 42 t A 43 + A 44 + A 45 t A 46 t A 47 + A 48 t A 49 t A 50 + A 51 + A 52 + A 53 + A Fig.2. Percentages of species present for each stage of the titration of an acidic solution of sodium pyrophosphate with sodium hydroxide solution. A, Species LH; and B, species LH,, where L = pyrophosphate (P,0,4-). Formation constants of calcium complexes. Complex equilibria between calcium and the tripolyphosphate anion can involve various c o m p l e x e ~ ~ ~ ~ ~ ~ ~ ~ ~ including CaP,OIo3-, CaHP3OIo2-, CaOHP3012- and CaNaP30,,2-.The evaluation of formation constants in these systems can be hampered33 by limited solubility. For example, for the formation of Ca,(P,01,)2 the solubility product is 1.6 x968 CRAGGS, MOODY AND THOMAS : CALCIUM ION-SELECTIVE ELECTRODE Analyst, VoZ. 104 10-37 at 25 "C and zero ionic strength.34 Nevertheless, although in the present titrations the solubility product was probably exceeded no visible precipitation occurred during the titrations. For the MINIQUAD program, species such as CaNaP3O1,2-, NaHP30,,3- and NaP3OlO4- were not included in the computation and hence the calculated formation con- stants are conditional and valid for I = 0.10 M sodium chloride. Calculations corresponding to the titration points (runs 1-4 in Table 111) indicated that over the concentration range studied the calcium tripolyphosphate system was best repre- sented by the formation of two species, namely CaP3OIo3- and Ca(P3010)28-.Data selected from appropriate titration points where these species would be predominant were then selected to obtain refined values for their formation constants (runs 5 and 6 in Table 111). Of the refined values, the output of run 5 (Table 111) was taken for ML and of run 6 (where the ligand was in excess) for ML,. TABLE I11 MINIQUAD OUTPUT FOR THE TRIPOLYPHOSPHATE - CALCIUM SYSTEM No. of Computer titration run points Model* Log P Standard deviation Sum of squares R factor 1 136 ML 5.12 0.02 0.157 x 10-6 0.112 2 136 ML + (M),L 5.13 (ML); -ve [(M),I-] 0.03 (ML) 0.289 x 0.103 7 136 ML + M(L)z 5.35 (ML); 10.08 [MIL),] 0.24 (ML); 0.83 [M(L),] 0.215 x 0.282 4 136 ML + MLH 7 0 (ML). 12 0 (MLH) 1.29 (ML); 89.7 (MLH) 0.961 x 0.597 5 50 ML + M(L), 5:05 (ML!; 9:86 [M(L),] 0.02 (ML); 0.64 [M(L),] 0.291 x 0.069 6 26 ML + M(L), 9.41 [M(L),]; -ve (ML) 0.07 [M(L),I 0.438 x lo-* 0.084 * M = calcium; L = tripolyphosphate; H = hydrogen.Similar considerations apply to the pyrophosphate system although the formation con- stants are lower. A 1 : 1 calcium - pyrophosphate complex is well established18j28J1J5 for calcium, whereas with other metals higher complexes, e.g., CU(P,O,),~-, have been observed.36 Protonated species, such as CaHP,O,-, have also been reported.31 Assumed sodium-containing species, such as NaP,0,3-, CaNaP,O,- and NaHP,O,%, were excluded from computations so that the formation constants are again conditional to 0.10 M ionic strength maintained by sodium chloride.Of the data in Table IV the most appro- priate model for the pyrophosphate - calcium system points to CaP,0,2- and Ca(P,O,),6- (run a), as negative values were obtained for log p of complexes in the succeeding runs. TABLE IV MINIQUAD OUTPUT FOR THE PYROPHOSPHATE - CALCIUM SYSTEM No. of Computer titration run points Model* Log P Standard deviation Sum of squares R factor 1 91 ML 4.37 0.01 0.117 x 0.035 2 91 ML + M(L)z 4.33 (ML); 7.21 [M(L),] 0.01 (ML); 0.08 [M(L),] 0.994 x 0.032 7 91 ML + MLH 4.37 (ML); -ve (MLH) 0.01 (ML) 0.117 x 0.035 4 91 ML + (M) L 4.37 (ML); -ve [(M),L)] 0.01 (ML) 0.117 x lo-' 0.035 5 82 ML + M(i), + 4.30 (ML); 7.01 [M(L),]; 0.006 (ML); 0.08 [M(L),] 0.155 x 0.013 M(Lh -ve [M(L),I * M = calcium; L = pyrophosphate; H = hydrogen. Results Table V summarises the stability constant (log p) data obtained by the above procedures and the range of calcium-ion concentrations measured by the calcium ion-selective electrodes.Table VI gives a comparison of log p data obtained in this work with literature data obtained by other methods. Discussion The data in Table V demonstrate that calcium ion-selective electrodes based on calcium bis [di (4-octylphenyl)phosphate] sensor can be used for free calcium-ion measurements in the presence of the various complexing anion ligands over a very wide concentration range, extending from below The fact that to above 1 0 - 3 ~ free calcium-ion concentration.October, 1979 MEASUREMENTS IN THE PRESENCE OF COMPLEXING LIGANDS TABLE V STABILITY CONSTANT (LOG p) DATA FOR VARIOUS CALCIUM - ANION LIGAND SYSTEMS Cal+ + Caa+ + Caa+ + CaP+ + Caa+ + Ca4+ + Caa+ + Cat+ + Caa+ + Ca*+ + Gas+ + Ca'+ + Equilibrium at 25 "C citratea- + Ca citrate malatea- + Ca malate malonates- $ Ca malonate oxalatea- $ Ca oxalate EDTAP-? + [CaEDTA]'- NTA*- + [CaNTAI- Soda- f Ca, SO4 HP0,'- + CaHPO, P,OlOs- + [CaP,0,,18- 2(P3010)*- + [Ca(P,Olo)zls- P,0,4- $ [CaP,O,]*- 2(P,0J4- + [Ca(P,O,),Ia- Ionic strength/ Titrations No.of M Range of [Caa+]rree/~ (n) runs 0.10 4.12 x to 7.10 X 33 3 0.10 3.12 x to 4.80 x lo-' 42 4 0.10 4.12 x 10-4 to 9.71 x 10-4 14 2 20 7 1 0.15 - (5 to 7.5) x 10-6 24 1 0.10 9.37 x 10-8 to 2.87 x 10-4 24 4 0.10 3.27 x 10-8 to 6.56 x lo-' 20 3 0.10 4.63 x to 2.47 x 28 4 6.13 x 10-6 to 9.84 x lo-' 14 1 5.12 x 10-6 to 2.19 x 25 3 8.56 x 10-6 to 9.92 x 11 1 - (1 to 6) x 75 6* 50 l { { } 0.10 -10-8 to -10-4 } 0.10 -10-6 to -10-4 91 969 Log P ( U " - l ) 3.42 (0.03) 2.00 (0.08) 1.52 (0.02) 2.70 (0.03) 2.54 (0.09) 1.99 (0.04) 10.93 (0.01) 6.31 (0.01) 1.39 (0.005) 2.14 (0.08) 1.87 (0.06) 1.57 (0.07) 5.05$ (0.02) 9.415 (0.07) 7.217 (0.08) 4.337 (0.01) * Different electrodes used for each run.t H,EDTAs- present (see text). $ Computer model 5 of Table 111. 5 Computer model 6 of Table 111. 7 Computer model 2 of Table IV. the log fl data fall into the range of those measured previously by various alternative methods is indicative of the useful scope of the calcium ion-selective electrode in monitoring free calcium-ion levels without disturbing the prevalent equilibria.This is of special importance as many of the anions included in this study occur in analytical samples. The resistance of this electrode system to sodium and other cation interference5 is also helpful in such appli- cations. The formulae of the complexes contributing to the equilibria summarised in Table V are generally well established. Nevertheless, the tripolyphosphate and pyrophosphate systems call for further comment on this point and also regarding the effect of supporting electrolyte. The formation constant for CaP,OlO3- (log /I1 = 5.20), although equal to the value obtained by Ellison and MartelP2 (log = 5.20 for I = 0.1 M maintained by potassium chloride) is, as expected, lower than the value obtained (log Bl = 6.41) when the 0.1 M ionic strength is maintained by tetramethylammonium bromide17 because of the contribution to binding by alkali metal ions as mentioned above.The existence of Ca(P,010),8- has been inferred previously37 and log /3, for this species is quoted in Table V. Although the high negative charge suggests its existence is unlikely, it is interesting and relevant that species such as Mn(P2O7)20- have been reported.,* Never- theless, the results obtained from a program such as MINIQUAD by indicating the ML, species cannot be regarded as unequivocal, for they depend on the permutations of chemical models and experimental data. The constraints imposed by working at such low concentrations must also be considered but, on the available evidence, it can be deduced that the most probable complexes formed in the calcium - tripolyphosphate system are those named in Table V.For the pyrophosphate system, the computer program also led to a two-complex model, namely CaP,O,,- and Ca(P,0,),6- (Tables IV and V). Again log 18, for the 1 : 1 complex is lower than that for the 0.1 M ionic strength system maintained by tetramethylammonium bromide. No comment can be made on the 1 : 2 complex, for which no other data have been published, although similar higher complexes have been observed for other metals,31 e.g., CU(P,O,),~-. Protonated complexes, such as CaHP,O,-, which had a negative log ,8 value for the model of run 3 (Table IV), have also been reported.,l The ionic strength studies with the oxalate and orthophosphate systems yieldd the following equations for variations of log ,8 with ionic strength, I : log p = 3.12 - 7.1 I (oxalate) .. . . .. - (8)TABLE VI STABILITY CONSTANTS FOR CALCIUM - ANION LIGAND SYSTEMS DERIVED BY SEVERAL METHODS AT 25 "C FOR EQUILIBRIA SUMMARISED IN TABLE V Complex Ca citrate Ca malate , Ca malonate Ca oxalate [ CaEDTA1'- Ionic strength 0.1 0 0 0.15 0.1 0.16 0 0 0.1 0.2 0.1 0.04 0.1 0.2 0.1 0 1.0 0.05 0.1 0.15 0.1 0.1 0.1 0.1 0.1 0.1 Medium NaCIO, NaCl NaCl NaCl NaClO, KCI NaCl NaCIO, KCl NaCl NaCIO, NaCl - - - - - - KNOa KCI NaCIO, KNO, KCI NaCl Log P Method 3.67 CaP+ electrode 4.68 Ion exchange 4.90 Solubility 3.17 Amalgam electrode 3.42 Ca*+ electrode 2.06 Ion exchange 2.66 Conductivity 2.24 Kinetic 2.66 Caz+ electrode 1.80 H, electrode 2.00 Ca*+ electrode } Glass electrode 1.46 H, electrode 1.52 Caz+ electrode 3.00 Conductivitv (18 OC\ 1.66 Distribution ' ' 1 } Caz+ electrode 1 1.99 11 Electrophoresis (20 "C) 10.59 Glass electrode (30 "C) 10.7 Amalgam electrode 10.42 Glass electrode (25.3 "C) 10.57 Hydrogen electrode 10.93 Ca*+ electrode Reference Complex 6 [CaNTAI- ..39 40 41 This work Ca,SO, .. 42 43 44 45 This work 6 CaHPO, .. 48 49 This work 17 17 17 17 17 This work [CaP,0,I2- . . Ionic strength 0.1 0.1 0.1 0 0 0 0.1 0.2 0.05 0.1 0.15 0.1 0.1 0.1 0.1 0.i Medium Log p Method Reference KNO, 6.57 Calorimetry 17 KCI 6.46 Glass electrode (20 "C) 17 NaCl 6.31 Ca2+ electrode This work - 2.0 Solubility (20" C) 17 - 2.31 Various thermodynamic data 17 - 2.27 to 2.48 Freezing point (0 "C) 17 NaCl 1.39 Cae+ electrode This work (C,H,),NBr 1.70 Glass electrode 50 NaCl { ptii } Caz+ electrode This work KCI 5.20 Glass electrode 32 (CH,),NBr 6.41 Nephelometry 17 NaCl 5.05 Caz+ electrode This work (CH,),NBr 5.3 9 Nephelome try 17 NaCl 4.33 CaL+ electrode This workOctober, 1979 MEASUREMENTS IN THE PRESENCE OF COMPLEXING LIGANDS 971 and The respective correlation coefficients are -0.953 and -0.999.log /3 = 2.43 - 5.7 I (orthophosphate) . . . . * - (9) Effect of Phosphate on Calcium Ion-selective Electrode Response Finally, it is necessary to allay previous fears that phosphate can interfere apparently inexplicably with the response of calcium ion-selective electrodes. Thus, an Orion 92-20 calcium ion-selective electrode based on a calcium bis(dialky1phosphate) sensor led to the expected fall in apparent free calcium-ion level on addition of sodium tripolyphosphate, but showed an increase when the ratio of about 1: 1 sodium tripolyphosphate to calcium was exceeded.38 Such a minimum was not observed with the calcium ion-selective electrodes used in this study and the anomaly can be attributed to the susceptibility of the Orion electrode to sodium-ion interference.That the actual decrease in apparent free calcium-ion level on addition of phosphate to a calcium-containing solution is due to complexation of calcium by phosphate may be con- firmed by adding phosphate to a solution of calcium ions buffered to a fixed level with a complexing ligand such as citrate. Thus, a calcium ion-selective electrode reference electrode pair immersed in a solution containing 2 x loF2 M sodium citrate and M calcium chloride maintained at ionic strength 0.1 M with sodium chloride and adjusted to a pH of about 10 with additions of sodium hydroxide and/or hydrochloric acid forms such a buffered system with a free calcium-ion concentration of ca. 3 x Such a cell gave a steady e.m.f.(-31.9 to -33.1 mV calcium ion-selective electrode potential with respect to the saturated calomel electrode) and additions of concentrated trisodium orthophosphate (system pH = 9.92), tetrasodium pyrophosphate (system pH = 9.97) and pentasodium triphosphate (system pH = 10.07) until the cell solution contained M phosphate did not alter the initial e.m.f. readings, thus confirming absence of the type of interference previously observed38 for an earlier calcium ion-selective electrode model.The authors thank the Science Research Council for a studentship (to A.C.) under the CASE scheme in conjunction with Unilever Research, Port Sunlight Laboratory. They also thank Dr. B. J. Birch for very helpful discussions. M. 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 16. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. References RbZiCka, J., Hansen, E. H., and Tjell, J . C., Analytica Chim. Acta, 1973, 67, 155. Brown, H. M., Pemberton, J . P., and Owen, J . D., Analytica Chim. Acta, 1976, 85, 261. Ammann, D., Giiggi, M., Pretsch, E., and Simon, W., Analyt. Lett., 1975, 8, 709. Birch, B., Craggs, A., Moody, G. J . , and Thomas, J . D. R., in Pungor, E., Editor, “Ion-Selective Craggs, A., Moody, G.J., and Thomas, J. D. R., Analyst, 1979, 104, 412. Rechnitz, G. A., and Hseu, T. M., Analyt. Chem., 1969, 41, 111. Rechnitz, G. A., Science, 1975, 190, 234. Fogt, E. J., and Rechnitz, G. A., Arch. Biochem. Biophys., 1974, 165, 604. Blay, J . A., and Ryland, J . H., Analyt. Lett., 1971, 4, 653. Moody, G. J., Oke, R. B., and Thomas, J . D. R., Analyst, 1970, 95, 910. Craggs, A., Moody, G. J., and Thomas, J . D. R., J . Chem. Educ., 1974, 51, 541. Craggs, A., Delduca, P. G., Keil, L., Key, B. J., Moody, G. J., and Thomas, J. D. R., J . Inorg. Nucl. “Instruction Manual for Calcium Ion-Electrode Model 92-20,” Orion Research Inc., Cambridge, Keil, L., Moody, G. J., and Thomas, J . D. R., Analyst, 1977, 102, 274. Bond, A. M., and Hefter, G., Inorg.Chem., 1970, 8, 1021. Ringbom, A., “Complexation in Analytical Chemistry,” Wiley-Interscience, New York, 1963. Martell, A. E., and SillCn, L. G., “Stability Constants of Metal-Ion Complexes,” Spec. Publ. Nos. 17 Overbeek, J . T. G., and Wolhoff, J. A., R e d . Trav. Chim. Pays-Bas Belge, 1959, 78, 759. Rossotti, F. J. C., Rossotti, H. S., and Wherell, R. C., J . Inorg. Nucl. Chem., 1971, 33, 2051. Ingri, N., and SillCn, L. G., A r k . Kemi, 1964, 23, 47. Sayee, I. G., Talanta, 1968, 15, 1397. Gans, P., Sabatini, A., and Vacca, A., Talanta, 1974, 21, 53. Gans, P., Sabatini, A., and Vacca, A., Inorg. Chim. Acta, 1976, 18, 237. Graham, R. D., and Williams, D. R., J . Chem. Soc. Dalton Trans., 1974, 1123. Brookes, G., and Petit, L. D., Chem. Commun., 1974, 813. Israeli, M., Loing, D. K., and Petit, L. D., J . Chem. SOC. Dalton Trans., 1974, 2194. Electrodes,” Akademiai Kiad6, Budapest, 1978, p. 335. Chem., 1978, 40, 1483. Mass., 1966. and 25, Chemical Society, London, 1964 and 1971.972 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. CRAGGS, MOODY AND THOMAS May, P. M., personal COM~UniCatiOn. Collis, C. F., and Van Wazer, J. R., Chem. Rev., 1958, 58, 1011. Powell, J . E., Farell, J . L., Neillie, W. F. S., and Russell, R., J . Inorg. NucZ. Chem., 1974, 20, 2223. Leggett, D. J., Talanta, 1977, 24, 535. Lambert, S. M., and Watters, J . I., J . Am. Chem. SOC., 1959, 81, 3201. Ellison, H., and Martell, A. E., J . Inorg. Nucl. Chem., 1964, 26, 1555. Quimby, 0. T., J . Phys. Chem., 1954, 58, 603. Birch, B. J., personal communication. Vasil’ev, V. P., and Yatsimirskii, K. B., Zh. Fiz. Khim., 1956, 30, 28. Reynolds, C., and Rogers, L., J . Am. Chem. SOC., 1949, 71, 2081. Blay, J . A., and Ryland, J . H., Analyt. Lett., 1971, 4, 653. Clarke, D. E., personal communication. Davies, C. W., and Hoyle, B. E., J . Chem. SOC., 1955, 1038. Bates, R. G., and Pinching, G. D., J . A m . Chem. SOC., 1949, 71, 1274. Joseph, N. R., J . Biol. Chem., 1946, 164, 529. Schubert, J . , and Lindenbaum, A., J . A m . Chem. SOC., 1952, 74, 3529. Topp, N. E., and Davies, C. W., J . Chem. SOC., 1940, 87. Bell, R. P., and Waind, G. M., J . Chem. SOC., 1951, 2357. Cannan, R. K., and Kibrick, A., J . A m . Chem. SOC., 1938, 60, 2314. Stock, D. I., and Davies, C. W., J . Chem. SOC., 1949, 1371. Campi, E., Annuli Chim., 1963, 53, 96. Money, R. W., and Davies, C. W., Trans. Faraday SOC., 1932, 20, 609. Hasegawa, Y., Maki, K., and Sekine, T., Bull. Chem. SOC. Japan, 1967, 40, 1845. Alberty, R. A., and Smith, R. M., J . A m . Chem. SOC., 1956, 78, 2376. Received ApriZ 20th, 1979 Accepted May 24th. 1979

ISSN:0003-2654    

DOI:10.1039/AN9790400961

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Analyst, Octobei, 1979, Vol. 104, pp. 973-976 973 Interferences of a Barium Ion-selective Electrode Used for the Potentiometric Titration of Sulphate Dilys L. Jones, G. J. Moody and J. D. R. Thomas and Magdolna Hangos Chemistry Department, University of Wales Institute of’ Science and Technology, Cardifl, CF1 3N U Institute for Geneva1 and Analytical Chemistry, Technical University, 1502 Budapest X I , Hungary Certain cations and anions interfere with the potentiometric titration of sulphate with barium chloride as titrant and a barium ion-selective indicator electrode. Thus, cations, for example potassium, which can be sensed by the barium ion-selec tive electrode, give distorted titration curves, while pH values below 1.5 lead to electrode breakdown. Cations, such as calcium, which interact with sulphate give low sulphate recovery, while anions that inter- act with the barium titrant lead to apparently high sulphate.Treatment of samples with cation-exchange resins in the sodium form can remove cation interferences, but acidification to pH 2 with hydrochloric acid prevents interference from anions, such as phosphate, carbonate - hydrogen carbonate and organic anions. Keywords : Barium ion-selective electrode ; sulphate titration ; interferences A sensor based on the tetraphenylborate(II1) (TPB) salt of a barium complex with a nonyl- phenoxypoly(ethy1eneoxy)ethanol (NP) (Antarox CO-880.Ba.2TPB), forms the basis of good barium ion-selective electrodes,ls2 and the use of 2-nitrophenyl phenyl ether as solvent mediator with the sensor in a PVC matrix membrane gives electrodes with long lifetimes.2 Among the possible applications of the PVC matrix-membrane electrodes is the potentio- metric titration of sulphate with barium ions, such a finish having been described for determining sulphur in organic compounds by the oxygen-flask method2 and for determining sulphate in the presence of chr~mium(VI).~ The electrode has also been tested for deter- mining sulphate by analate ~ubtraction.~ A liquid-membrane barium ion-selective electrode with the same sensor but with 4-nitroethylbenzene as solvent mediator has been used for the potentiometric titration of sulphate in sea water and certain natural waters5 These and other analytical systems in which sulphate has to be determined can frequently be complex; hence, the present study was undertaken to assess the effects of various possible interferents on potentiometric titrations using the PVC matrix-membrane barium ion- selective electrode as the indicator electrode.Experimental Ion-selective Electrode and Potentiometric Titration Assembly Barium ion-selective indicator electrodes, fabricated as previously described,2 were used in the titration system in conjunction with a saturated low-leak (approximately 0.01 cm3 h-l) calomel reference electrode (Corning No. 476107). The e.m.f .s of the potentiometric titra- tion cell were recorded with a high-impedance millivoltmeter (Corning, Model 112) reading to 0.1 mV. Chemicals The chemicals used were of analytical-reagent grade. Antarox CO-880, used as the starting material for making the barium ion sensor, was a gift from GAF (Great Britain) Ltd., Manchester (PVC matrix-membrane electrodes selective to barium ions and based on the same sensor are available from EDT Research, London.) Titration System sulphates. anions or in hydrochloric acid.Barium chloride solution (0.1 M) was used to titrate 25 cm3 of 0.01 M solutions of various The sulphates were dissolved in water, in aqueous solutions of sodium salts of974 JONES et al. : INTERFERENCES OF A BARIUM ION-SELECTIVE Analyst, Vd. 104 Results and Discussion The parameters studied included low pH and the effect of various cations and anions, the titration system having previously been shown to give a theoretical equivalence point for 0.01 M sodium sulphate (Fig. 1, curve A).Effect of Acid Titration of 0.01 M sulphuric acid gave a theoretical equivalence point and the titration curve (Fig. 1 , curve B) was similar to that for sodium sulphate (Fig. 1 , curve A). The corresponding titration of sodium sulphate dissolved in 0.1 M hydrochloric acid showed a considerably decreased step at the equivalence point (Fig. 1 , curve C) but titrations in 0.1 M hydrochloric acid had a deleterious effect on the electrode, the slope to barium ion response dropping to about 19 mV decade-l. These observations in acidified solutions are consistent with the effective limit of about pH 1.5 in the range of barium ion-selective electrodes for 0.001 M barium chloride.2 Effect of Cations according to the cation present. Potassium and ammonium The break in the titration curve for potassium and ammonium sulphate was considerably reduced because of the potassium and ammonium ion sensitivity of the electrode prior to the equivalence point (Fig. 2, curves B and C).A later titration of potassium sulphate showed some hysteresis, although the end-point was properly detected (Fig. 2, curve D). A subsequent titration of sodium sulphate (Fig. 2, curve E) showed a conventional titration curve, but with a shortened break at the end-point, demonstrating the deterioration of barium ion-selective electrodes on prolonged contact with potassium and ammonium ions. Such deteriorations did not occur with successive titrations of sodium sulphate. In order to control and recognise any hysteresis by ions other than those present in titra- tions of sodium sulphate with barium chloride, the remaining titrations of this study were undertaken with fresh barium ion-selective electrodes that had been checked in titrations of sodium sulphate.Titrations of 0.01 M sulphates other than sodium sulphate gave various interesting effects -20 -40 -60 2 4 Volume of 0.1 M BaCI, solution/cm3 Fig. 1. Potentiometric titration of 25 cm3 of 0.01 M sulphate solution: A, sodium sulphate ; B, sulphuric acid ; and C, sodium sulphate plus 0.1 M hydrochloric acid. r 20 0 -20 -40 -60 2 4 Volume of 0.1 M BaCI, solution/cm3 Fig. 2. A series of potentiometric titrations of 25 cm3 of 0.01 M sulphate solutions in sequence A to E: A, sodium sulphate solution; B, potassium sulphate solution ; C , ammonium sulphate solution; D, potassium sulphate solution; and E, sodium sulphate solution.October, 1979 ELECTRODE USED FOR POTEN~OMETRIC TITRATION OF SULPHATE 975 In an assessment of the effect of increasing amounts of potassium chloride, added to the sodium sulphate, these gave a reproducible hysteresis when 0.01 M sodium sulphate was present in a background of 0.01 M potassium chloride.This was characterised by a steady fall in e.m.f. from the commencement of the titration to a minimum (cf, Fig. 2, curve D) before the increase at the equivalence point. The titration curve became conventional in shape on raising the potassium chloride background to 0.1 M, except that the shorter inflection step at the end-point elongated considerably from the vertical. A further increase in back- ground potassium chloride to 1 M resulted in an almost horizontal titration curve of e.m.f.versus volume of barium chloride solution at about 5-10 mV veyszfis S.C.E. Despite the above characteristics, the end-points could be properly detected at background potassium chloride levels of 0.1 M and less, and the mean sulphate recovery was 100.1 (standard deviation, Other cations added to sodium sulphate. = 0.90)0/$. Table I summarises the effects of various cations in metal sulphates and also of metal salts TABLE I POTENTIOMETRIC TITRATION CURVE DATA AND SULPHATE RECOVERIES I N THE PRESENCE OF VARIOUS CATIONS Break at inflection* to 0.01 M Na,SO,) Composition of solution (normalised relative SO,2- recovery (from titration curve), yo 100.0 (s.d., = 0.66) 100 100 ... . .. .. 0.81 100 .. .. .. .. 0.83 100 (with 0.01 M SO,,-) :t Na,SO, . . .. .. .. .. .. .. .. .. . . .. .. .. .. . . 0.66 H2S04 3%, .. (NH4)2S04 * * A12(S04) S * * .. .. .. . . 1.07 100 FeSO, .. .. .. .. .. 1.15 100 MgSO4 Na,SO, + lo-* M Fe(NO,), .. .. 1.03 100 .. . . .. .. 1.03 100 .. . . 0.65 100 Na,SO, + 0.1 M HC1' Na2S04 + <$ x lo-, M Fe(NO,), . . .. 1 .oo 98.8 Na,SO, + 1.7 x M Fe(NO,), . . 0.73 97.6 Na,SOI + 4 x M Fe(NO,), . . .. 0.70 94.8 Na,SO, + 2.9 x lo-, M CaCl, . . . . 0.69 93.0 . . Na,SO, + 1.5 x M CaC1, . . .. 0.78 93.9 Na,SO, + 6 x M CaCl, . . . . 0.63 93.9 * This is AE corresponding to 20% of the titration each side of the end-point. t For 11 electrodes, the mean inflection was 69.6 (s.d., i3n-l = 3.85) mV. Good near-vertical inflections were obtained for all of the titration curves.Breaks smaller than those of the reference sodium sulphate may be attributed to a greater sensitivity of the electrode for the metal ion than for sodium ions. The converse applies to larger breaks, for example, iron(I1) and magnesium. These trends may be deduced from the general pattern of selectivity coefficient data, although it must be realised that k:;:= data are not directly comparable for different valence states of B ions 9.5 x 4.4 x 6.0 x 2.2 x 3.7 x 6.8 x and 6.6 x for B = Na+, K+, NH4+, Ca2+, Mg2+, Fe2+, Fe3+ and A13+, respectively]. Prolonged contact of the electrode with interfering ions led to shortened voltage swings, as illustrated by curve E (Fig. 2) compared with curve A (Fig. 2). Low sulphate recoveries in the presence of calcium and iron(II1) salts arise from the expected interactions of the metal ions with sulphate, which can occur with any cation yielding a low-solubility sulphate.The iron interferences here could be removed by passing the test solution through a column of Zerolit 225 (Na+ form) ion-exchange resins, but because of the low solubility of calcium sulphate, the resin was only partially successful for eliminating calcium-ion interference. Thus, for example, the sulphate recovery from 0.01 M sodium sulphate containing 2.9 x M calcium chloride rose only to 97% on ion-exchange resin treatment. Prasad3 overcame chromium(V1) interference by reducing with ascorbic acid ; this also had the effect of complexing chromium(II1). (for 0.1 M of B) = 3.0 x976 JONES, MOODY, THOMAS AND HANGOS Effect of Anions Apart from some shortening of the titration curve inflections, backgrounds of 0.1 M nitrate or chloride are without effect, but, in addition, anions likely to give insoluble barium salts were examined, with the results summarised in Table 11.TABLE I1 POTENTIOMETRIC TITRATION CURVE DATA AND SULPHATE RECOVERIES I N THE PRESENCE OF VARIOUS ANIONS ADDED AS SODIUM SALTS Break a t inflection* Concentration of added sodium salt (normalised to SO,2- recovery (from (with 0.01 M Na,SO,) 0.01 M Na2S0,) titration curve), yo 0.1 M c1- . . . . .. .. . . .. 0.96 100 0.1 M NO3- . . . . .. .. .. .. 0.86 100 6 x lo-, M HPO,- . . . . .. .. . . 0.87 104 1.2 x 10-3 M HPO,- . . . . . . .. .. 0.67 110 2.4 x 10-3 M HPO,- . . . . . . .. . . 0.74 116 2.9 x 10-3 M HPO,- . . . . ,. .. .. 0.81 121 8 X lov3 M HC03- . . .. .. . . .. 0.81 105 2 x M HC03- . . . . . . .. .. 0.70 114 3.5 x M HC03- . . . . . . . . . . 0.64 124 2.9 x 10-3 M HCO,- . . . . . . . . . . 0.75 115 1 x M salicylate- . . .. .. .. 0.79 103 2.9 x M salicylate- . . .. . . . . 0.79 104 2.9 x M oxalate2- . . . . . . . . 1 106 2.9 x 1 0 - 3 ~ citrate3- . . . . . . . . 0.78 113 2.9 x 10-3 M citrate3- adjusted to pH 2 with HC1 1 100 * See footnote to Table I. The general effect is a shortening of the step at the inflection corresponding to the end- point and an over-consumption of barium chloride titrant for those anions likely to interact with barium. The over-consumption of titrant increases with increase in the level of inter- acting anion.The titration curves were of conventional shape (cf., Fig. 1) in all instances except for citrate, where there was considerable elongation, but prior adjustment to pH 2 with hydro- chloric acid resulted in conventionally shaped titration curves and 100 yo sulphate recovery (Table 11). Indeed, the interferences of all of the organic and inorganic anions could be removed by previously adjusting the pH to 2 with hydrochloric acid, when theoretical recoveries of sulphate were obtained. Conclusion Various cations and anions interfere in the potentiometric titration of sulphate with barium chloride using a barium ion-selective indicator electrode. Although the cations would require removal, acidification to pH 2 with hydrochloric acid controls many anion interferences. The authors are grateful for support under the Programme of Cultural, Education and Scientific Exchanges of the British Council between Britain and Hungary, permitting a short stay (by M.H.) at UWIST for commencing the study, and to Professor E. Pungor and Dr. K. T6th for their interest. The Science Research Council is also thanked for a studentship (to D.L. J.) under the CASE scheme in conjunction with Unilever Research Laboratory, Port Sunlight, whose representative, Dr. B. Birch, is thanked for helpful hscussions. References 1. 2. 3. 4. 5. Levins, R. J., Analyt. Ckem., 1971, 43, 1045. Jaber, A. M. Y., Moody, G. J., and Thomas, J . D. R., Analyst, 1976, 101, 179. Prasad, R., Analyst, 1979, 104, 164. Moody, G. J., and Thomas, J. D. R., Lab. Pract., 1979, 28, 125. Ouzounian, G., and Michard, G., Analytica Chim. Acta, 1978, 96, 405. Received April loth, 1979 Accepted May 24th, 1979

ISSN:0003-2654    

DOI:10.1039/AN9790400973

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Analyst, October, 1979 SHORT PAPERS 977 Detection of L-Cysteine, Methionine, Thiourea, Allylt hiourea and a- Nap hthyl thiourea in Sub-milligram Amounts Using Acidified Potassium Chromate Solution M. Nasim Beg, Fasih A. Siddiqui, M. Mohtashim Beg, R. Shyam and M. Arshad Department of Chemistry, Aligarh Muslim University, A ligarh-202001, India Keywords T h i o compound detection ; acidified potassium chromate colour reagent A number of methods for the identification of thio compounds are available.'-7 Feig17 employed the catalytic acceleration of the iodine - azide reaction for the small-scale identifi- cation of thioketones and thiols. Kiroaki et a1.8 detected cysteine, related thiols and disulphide compounds on thin-layer chromatograms using N-(9-acridinyl)maleimide as a fluorescent detection reagent in a buffer of pH 8.8.Wawrzyczekg used osmium(VII1) oxide (OsO,) for the detection of cysteine, methionine and taurine in the presence of an acetate buffer. Garcia-Blanco and Pascual-Leonelo used vanadic acid for the identification of sulphur conta,ining compounds such as cysteine, cystine and methionine. These methods are tedious and require a large amount of the sample for the analysis. We describe here a simple method for the detection of L-cysteine, methionine, thiourea, allylthiourea and a-naphthylthiourea using acidified potassium chromate solution as a colour reagent. The amount of the substance required for the detection is of the sub-milligram order. Experimental Reagents All reagents were of analytical-reagent grade (supplied by BDH Chemicals and E.Merck) and doubly distilled water was used throughout. The solutions of methionine, thiourea] allylthiourea and cc-naphthylthiourea were prepared directly] by dissolving the substances in water, whereas L-cysteine was first dissolved in a small volume of 4 N hydrochloric acid and this solution was then diluted with water. The acidified potassium chromate solution was prepared by adding 5 ml of 12 N hydrochloric acid to 3 ml of freshly prepared 1% potassium chromate solution. Procedure The acidified potassium chromate solution gave a green complex with methionine, thiourea, allylthiourea, a-naphthylthiourea and L-cysteine solutions. The intensity of the colour slowly increased with time and became stable after about 15 min. The green complex has a Amax.of 420 nm, irrespective of which thio compound is taken. A systematic study of the influence of the concentration of potassium chromate solution on the detection was carried out by taking 3 ml of 0.5, 1 , 2, 3, 4 and 5% solutions of acidified potassium chromate. To each solution 1 ml of a 100 pg ml-l solution of thio compound was added. It was seen that the solution containing 1% potassium chromate solution gave the colour with maximum intensity. Similarly, the effect of pH on the reaction was also studied by taking 5 ml of 4, 6, 8, 10 and 12 N hyd.rochloric acid. To each of the acid solutions 3 ml of the 1% potassium chromate solution and 1 ml of the 100 pgml-l solution of a thio compound were added. It was observed that the intensity of the colour increased with the increase in the acid concentra-978 SHORT PAPERS Analyst, Vol.104 tion. A Beckman DU spectrometer was used to record the absorbance of the green solutions, measured against a reagent blank. A graph of absorbance veisus wavelength was used to obtain the value for A,,,, Results and Discussion Atoms with unshared electrons, such as oxygen, nitrogen and sulphur, can act as ligands for various metallic ions. These atoms are abundant in proteins and thioureas, which, therefore, are effective chelating agents for various metallic ions. Sulphur, particularly in the form of thiol groups, is highly reactive biochemically and fulfils a number of vital functions in living matte+ ; for example, L-cysteine, HSCH,CH(NH,)COOH, forms com- plexes with metal ions12 and has three possible co-ordination sites at the nitrogen, oxygen and sulphur centres. The configurations adopted in these species are largely dependent on the reactive metal ion.It is largely because of the kinetic inertness that so many complex species of this ion have been isolated as solids and that they persist for relatively long periods of time in solution, even under conditions where they are thermodynamically unstable. I t was thought worthwhile, therefore, to study the complexes formed by chromium(II1) with a number of compounds such as L-cysteine, methionine, thiourea, allylthiourea and ct-naphthylthiourea with particular reference to their detection. The complexes formed by L-cysteine, methionine, thiourea, allylthiourea and a-naphthyl- thiourea with acidified potassium chromate solution are green, and the intensity of the colour increases slowly and becomes steady after about 15min.The intensity of the colour was greatly affected by the concentrations of both the potassium chromate solution and the hydrochloric acid. The green complex has a A,,,. of 420 nm, irrespective of the thio compound taken. The detection limits of the substances were established by adding 1% potassium chromate solution (in 12 N hydrochloric acid) to solutions containing different concentrations of the thio compounds and observing the intensity of the colour after about 15min. The limits of detection thus obtained are given in Table I. Chromium(II1) is known to have a large number of kinetically inert complexes. TABLE I LIMITS OF DETECTION OF THE THIO COMPOUNDS Limit of detkctionl Sample Pf3 L-Cysteine .. . . .. .. 100 Methionine .. . . .. 150 Allylthiourea . . .. .. 100 Thiourea . . . . . . .. 80 or-Naphthylthiourea . . .. 200 The method was also tested in the detection of a large number of other compounds such as glycine, alanine, cystine, 4-aminobenzoic acid, carbohydrates, chloroacetic acid and sulphosalicylic acid. However, these compounds gave no coloured complex with the potassium chromate solution. In the light of these investigations it is concluded that L-cysteine, methionine, thiourea, allylthiourea and a-naphthylthiourea can be detected iv sub-milligram amounts using acidified potassium chromate solution as a reagent. The reaction may be frequently employed in the detection of these compounds by the spot-test technique.References 1. 2. Feigl, F., and Anger, V., in “Spot Tests in Organic Analysis,” Seventh Edition, Elsevier, Amsterdam, Feigl, F.. and Anger, V., in “Spot Tests in Inorganic Analysis,” Sixth Edition, Elsevier, Amsterdam, 1966. 1972.October, 1979 SHORT PAPERS 979 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. Yufera, E. P., Parareda, J . S., and Aberda, J., Revta Agroquim. Technol. Aliment., 1965, 5 (2), 211. Cunha, 0. D. L. R., and Cunha, A. P. D., BoZ. Esc. Farm., Univ. Coimbra, Ed. Cient., 1958, 18, 130. Calvo, J. M., and Ruiz, A. S., Bull. SOC. Chim. BioZ., 1957, 39, 1557. Bayfield, R. F., and Cole, E. R., J . Chromat., 1969, 40, 470. Feigl, F., in "Spot Tests in Organic Analysis," Sixth Edition, Elsevier, Amsterdam, 1960. Hiroaki, T., Yasunari, N., and Katura, T., Agric. Biol. Chem.. 1976, 40, 2493. Wawrzyczek, W., 2. Analyt. Chem., 1962, 185, 446. Garcia-Blanco, J., and Pascual-Leone, A. M., Revta Esp. Fisid., 1955, 11, 149. Eldjarn, L., Scand. J . Clin. Lab. Invest., SuppZ., 1965, 86, 7. McAnliffe, C. A., and Murray, S. G., Inorg. Chim. Acta Rev. 1972, 6, 103. Received November 1 lth, 1978 Accepted April 6th, 1979

ISSN:0003-2654    

DOI:10.1039/AN9790400977

出版商:RSC    

年代:1979

数据来源: RSC

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October, 1979 SHORT PAPERS 979 Implementation of a Sensitive Method for Determining Mercury in Surface Waters and Sediments by Cold-vapour Atomic-absorption Spectrophotometry R. L. Lutze State Pollution Control Commission, Central Square Building, 323 Castlereagh Street, Sydney 2001, N . S . W . , Australia Keywords : Mercury determination ; atomic-absorption spectrophotometry ; cold-vapour method comparisons ; environmental samples The determination of mercury at low concentrations in surface waters and sediments has been an important task of the State Pollution Control Commission’s water laboratory for several years. An earlier method partitioned the reduced mercury between a fixed volume of gas and the liquid phase, while agitating vigorously with a magnetic stirrer. The equilibrated gas phase was then transferred into the absorption cell by a gas purge.The method gave only weak, transient absorptions and was time consuming and imprecise. Improved partitioning and aeration methods were implemented and compared , utilising conventional single-beam spectrophotometers. The adopted method was used routinely for over a year on unpolluted surface freshwaters and also on polluted sediments. Experimental and Results Reagents and Standards All glassware (pipettes, calibrated flasks, etc.) was soaked for at least 1 d in 20% nitric acid, followed by 5% nitric acid containing O.Olyo potassium dichromate solution, made up with doubly glass-distilled water.l The latter solution was used as the soak mixture for all glassware until they were required. The reagents used were of analytical-reagent grade and included tin(I1) chloride, potassium dichromate , magnesium perchlorate (granular) , nitric acid , hydrochloric acid, sulphuric acid, mercury(I1) chloride and phenylmercury(I1) acetate.Tin(I1) chloride was freshly prepared as a 20% suspension in 1 M sulphuric acid, and stored in a dark glass bottle under nitrogen gas. A stock solution containing 1.000 g 1-1 of mercury(I1) was prepared from mercury(I1) chloride in an aqueous solution of 5% nitric acid and 0.01 % potassium dichromate. Standard mercury solutions containing 100, 10 and 1 pg 1-1 of mercury(I1) made by dilution of aliquots of the stock solution, when made up to 5% in nitric acid and 0.01% in potassium dichromate, were stable for several months.Standard solutions in the concentration range 0.2-1.0 pg 1-1 prepared in the same way were stable for at least several days. The high oxidation potential of the acidic dichromate solution and its apparent surface coating properties2 together prevented detectable mercury losses by hydro- lysis or reduction.980 SHORT PAPERS Analyst, Vol. 104 Instrumentation The determinations were carried out on Varian Techtron, Model AA5 and Model 1250, atomic-absorption spectrophotometers, using the Varian Techtron BC-6 background corrector unit in all instances. Absorbance tracings were recorded by a Varian Techtron A25 high-speed chart recorder. Maximum signal stability was obtained by utilising overnight instrument warm-up before the determinations. Development and Description of Cold-vapour Methods In view of the high relative atomic mass and very low vapour pressure of elemental mercury at room temperature, it was thought that an aeration accompanying the reduction might enhance the sensitivity.A coarse fritted glass bubbler was used to mix and aerate the solution with dry nitrogen. A sample volume of 50 ml in a 125-ml reduction bottle was used, together with 2 ml of the 20y0 tin(I1) chloride suspension in sulphuric acid. The carrier gas flow-rate was reduced to 600 ml min-1 from the previously used 1.5 1 mi@. The reduction bottle was connected to a flow-through absorption cell via a moisture trap. Enhanced sensitivity and precision were obtained from this apparatus (see Table I). However, at lower flow-rates (which reduce the dilution of the mercury vapour) the aeration efficiency declined, and the absorbance signal became broader and erratic.Also, the moisture trap quickly became damp, and introduced some memory effects. In order to counter these shortcomings, a novel and improved partitioning method was devised, which would eliminate gas-purge dilution of the mercury vapour. The tin(I1) chloride reduction of the sample was performed in a glass syringe, where the air volume was twice that of the sample. Mechanical or manual agitation for 2 min transferred elemental mercury from the liquid into the vapour space. The vapour was then injected through a drying tube into the silica window enclosed absorption cell. A large, persistent absorbance was obtained, enabling a precise measurement to be made.The sensitivity obtained was so superior to the previous methods that the sample volume was reduced to 10 ml (see Table I). However, prompt purging of the cell was required in order to prevent mercury condensation and memory effects. Also, the charging, shaking and syringe-cleaning cycles were slow and liable to contamination. TABLE I COMPARISON OF COLD-VAPOUR GENERATION OF MERCURY METHODS Operating conditions : hollow-cathode lamp current 3.0 mA ; wavelength 253.6 nm ; spectral band pass 0.5 nm; damping “B”; recorder 10 mV; chart speed 2.5 cm min-’, background correction used; volume of tin(I1) chloride solution 2 ml. Mercury Mass of mercury Solution concentration/ Scale Method type injected/ng volume/ml pg I-’ expansion Gas-purged partitioning .. . . 0 50 0.00 7 x 10 50 0.20 7 x 20 50 0.40 7 x 50 50 1.00 7 x Bottle aeration . . . . . . 0 50 0.00 7 x 10 50 0.20 7 x 50 50 1.00 7 x Syringe-injected partitioning . 0 10 0.00 6 x 2 10 0.20 6 x 10 10 1.00 6 x 20 10 2.00 6 x Dual-bubbler . . . . . . 0 10 0.00 10 x 2 10 0.20 10 x 10 10 1 .oo 10 x Flow-rate/ ml min-’ 1500 1500 1500 1500 600 600 600 480 480 480 Mean absorbance (6 readings) 0.013 0.053 0.097 0.217 0.030 0.096 0.352 0.020 0.056 0.198 0.385 0.018 0.101 0.412 Relative standard deviation, % 23 20 13 8 15 10 8 12 7 5 4 10 6 4 An apparatus was desired that combined high speed of reduction with very low blank values and that still had sensitivity and precision comparable to those of the syringe- partitioning method. At that time it was thought that a re-designed aeration apparatus would be suitable if it eliminated the shortcomings noted earlier.October, 1979 SHORT PAPERS 981 The mercury vapour generation technique finally adopted was an aeration apparatus, termed a dual-bubbler system, as described by Simpson and Nick1ess.l A medium-porosity frit forms the bottom of each reduction tube and the tubes are enlarged n3ar their top to reduce foaming and foam carry-over. A long, narrow drying tube, using cotton-wool rather than glass-wool plugs, provided efficient drying for long periods with no obszrvable memory effects.On aeration of standard mercury solutions, absorbances were consistently equal to those obtained using the syringe-partitioning method, and were prolonged enough for precise recording either by a chart recorder or a peak-read facility. Further, the reductions were performed far more rapidly.Operational Advantages An additional advantage of the design includes pre-aeration of the tin(I1) chloride solution, removing any residual mercury from both. the solution and the glass bubbler prior to sample addition; this eliminates the need for any special “low in mercury’’ tin(I1) chloride reagent. Also, spent samples or standards are bottom-emptied under sealed conditions. The self- cleaning and self-emptying action contribute to the very low blank values obtained, as the apparatus is sealed off from the laboratory environment most of the time. Further, the cycle of sample reduction and aeration is rapid, and the dual-bubbler design permits two samples to be determined consecutively in less than 2 min.There is one other advantage to the bubbler reduction design not found in other methods examined, particularly the partitioning methods. The single reduction - aeration step removes all detectable mercury present, rather than some (perhaps) constant fraction of the mercury. This allows consecutive sample or standard injections to be carried out on a suitable volume of tin(I1) chloride solution, without replacing the solution. More importantly, it offers the analytical flexibility of standard calibration in the sample matrix, by injecting a known mass of mercury of suitably small volume into the mercury-spent sample and repeating the aeration. This is because the volume change is negligible, and hence the sensitivity is unaltered, and the tin(I1) chloride solution is present in large excess.This facility was found to be very helpful when analysing fish digests and some oily sedi- ment digests, as appropriate synthetic standard matrices would have been difficult to prepare, and the method of standard additions to a significant number of samples prior to digestion would have been too tedious. In all of the techniques of mercury vapour generation examined, a flow-through absorption cell was used, with subsequent scrubbing through an acidic potassium permanganate solution. It is thought that this precaution reduced laboratory contamination by elemental mercury and contributed to the very low blank values obtained. Calibraaion, Blanks and Detection Limit The calibration graphs have a high degree of linearity from 0 to 400ng of mercury; calibrations gave a correlation coefficient of unity.I t was found that the calibrations were stable from day to day during constant use, provided that optimum conditions were main- tained and the tin(I1) chloride solution was fresh. Three injections of the mercury standard were sufficient for calibration, as the range was within &2y0 of the mean peak height. For mercury masses between 10 and 100ng, the relative standard deviation was 2%. The blank value for 5% nitric acid and O.Olyo potassium &chromate solution was 25 ng 1-l. The detection limit, without scale expansion and using background correction, was 1 ng, or 0.1 p g 1-1 for a 10-ml aliquot. Using approximately 10 x scale expansion, the detection limit was 0.4 ng, or 0.04 pg 1-1 for a 10-ml aliquot.However, at a lamp current of 6 mA without background correction, the detection limit at high scale expansion (about 15 x ) was 0.1 ng or 0.01 pg 1-1 for a 10-ml aliquot. Water Analysis The Simpson and Nickless apparatus2 has been used routinely for the determination of mercury in samples from freshwater rivers in New South Wales. The sample bottles were982 SHORTPAPERS Analyst, Vol. 104 soaked repeatedly in the acidic dichromate rinse before sampling, and the river water was preserved with nitric acid and dichromate. The river water was not filtered, because the non-filtrable residues were less than 50 mg 1-l. Mercury concentrations in these river waters ranged from 100 to 1200 ng 1-1, and sample oxidation by the potassium permanganate - sulphuric acid method did not increase the amount detected, probably indicating the absence of organic mercury compounds. As these river waters were sampled from rural catchment areas, the concentrations are a good indication of the natural background levels to be found in unpolluted surface freshwaters.Sediment Analysis In the first, samples were weighed, before drying, into PTFE pressure digestion vessels, their moisture content being determined separately. Aqua regia was added, and the samples were left overnight at room temperature. The resulting suspensions were filtered, and rinsed, through Whatman 541 ashless filter-papers, using a doubly-distilled water wash containing O.Olyo of potassium dichromate, giving a total volume of 50 ml.The second method examined was similar to that ycommended by the US Environmental Protection Agency, and used samples that had been oven-dried at 60 "C. Random portions of the dried sediment were roughly powdered with the spoon-shaped end of a spatula, and weighed directly into 50-ml B-grade calibrated flasks fitted with ground-glass stoppers. Aqua regia was added and the flasks were warmed at about 50 "C for 3 h. The contents were diluted to volume with 0.01% potassium dichromate solution, the solution was mixed and the particulate matter was left to settle. The supernate was analysed for mercury using sample aliquots of 0.25-1.0 ml on to 1 ml of tin(I1) chloride solution. In this way, peaks corresponding to 20-100 ng of mercury were obtained, and foaming problems avoided.Mercury concentrations in a group of oily sediments from a polluted urban canal were found to be the same by both methods, and were in the range 0.3-2.0pgg-l (dry mass). The recovery of spiked additions of both mercury(I1) chloride and phenylmercury(I1) acetate was 100~o. The relative standard deviation obtained by weighing the sample wet was 13- 20%, and by weighing the sample dry less than half this range, demonstrating a considerable gain in precision and convenience for the second method. Two sediment digestion methods were used and compared. This was followed by oven heating at 95 "C for 1 h. Conclusions The dual-bubbler apparatus is the most rapid and sensitive cold-vapour mercury detection system so far evaluated in this laboratory. Its precision and low blank values have been verified by application to routine mercury analysis of surface waters and sediments, and its use is currently being extended to biological material, such as fish tissue. Further improve- ments in the detection limit are limited by the stability and intensity of mercury line emission sources and the stability and sensitivity of the detector and its peripheral electronics. A much lower detection limit is possible with the existing instrumentation by recourse to a pre-concentration step, followed by reduction to elemental mercury vapour in a minimum volume of carrier gas. However, these refinements reduce the speed of analysis and do not appear to be warranted for the State Pollution Control Commission's work at present. The earlier methods used in this laboratory and mentioned in the introduction were developed by Ms. K. Wyatt. References 1. 2. Simpson, W. R., and Nickless, G., Analyst, 1977, 102, 86. Feldman, C., Analyt. Chem., 1974, 46, 99. Received February 8th, 1979 Accepted March 9th, 1979

ISSN:0003-2654    

DOI:10.1039/AN9790400979

出版商:RSC    

年代:1979

数据来源: RSC

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October, 1979 SHORT PAPERS Spectrophotometric Determination of Trace 983 Amounts of Boron in Solutions Containing Large Amounts of Nitrate H. J. Rosenfeld and A. R. Selmer-Olsen Agricultural University of Norway, Chemical Research Laboratory, N - 1432 AS-NLH, Norway Keywords : Boron determination ; nitrate interference ; hydrazine hydrate ; carminic acid ; spectrop hotometry The best known reagents for the determination of microgram amounts of boron are 1,l’- dianthrimide, curcumin and carminic acid. The method described by Hatcher and Wilcoxl and its modifications by Callicoat and Wolszon2 are among the most common methods using carminic acid. The determination of boron in nutrient solutions used in horticulture and sewages is of interest. Such water samples may contain large amounts of nitrate, which seriously interfere with the spectrophotometric determination of boron in the 1 ,l’-dianthrimide, curcumin and carminic acid methods.These problems have been solved in the curcumin method by Goldman et al.3 and by Goulden and Kakar4 in the 1,l’-dianthrimide method. In the carminic acid procedure Hatcher and Wilcoxl added hydrochloric acid to prevent the interference of nitrate. Ross and White5 found that in the presence of large amounts of nitrate this proved to be unsatisfactory. Instead they introduced a step in order to destroy up to 3 M nitrate in solution using formic and sulphuric acids, under reflux conditions, before the formation of the carminic acid - boron complex. LionneP determined boron by an automated method in sewage, sewage effluents and river waters.By adding phenol the interference of up to 40 mg 1-1 of nitrate-nitrogen was reduced, but at higher concentrations an interfering effect occurred. The carminic acid procedure has been used for several years in our laboratory for deter- mining trace amounts of boron. In this investigation efforts were made to remove the interfering effects of nitrate by adding hydrazine hydrate to the samples. Experiment a1 Apparatus 50-mm light paths were used. S$ectrojdzotometer. Polyfiropylene beakers. A Hitachi, Model 100-20, spectrophotometer and cells with 10- and Reagents Szdphuric acid. Carminic acid. Carminic acid reagent. Hydrazine hydrate (NH2NH2.H20). Boric acid standard solution. Pro analisi grade (Merck) . Pro analisi grade (Merck).Prepared by dissolving 125 mg of carminic acid in concentrated sulphuric acid and diluting to 1 1 with concentrated sulphuric acid. Analytical-reagent grade, >99% purity (Fluka) . Prepared by dissolving 0.571 6 g of analytical-reagent grade boric acid (Riedel de Haen) in 1 1 of de-ionised water; 1 ml of solution is equivalent to 0.1 mg of boron. Procedure A 2-ml aliquot of the water sample is transferred into a boron-free polypropylene beaker, 2 drops (0.1 ml) of hydrazine hydrate are then added and well mixed; 10 mi of concentrated sulphuric acid are added, followed by 10 ml of carminic acid reagent. The solutions are well mixed and set aside, atmospheric moisture being excluded. Maximum colour develop- ment is reached after 90 min, according to Callicoat and Wolszon.2 This was confirmed in our laboratory.The difference in absorbance shown by increasing standing time from984 SHORT PAPERS Analyst, VoZ. 104 60 to 90 min was less than 2%. The absorbance of the solution is measured at 610 nm in 10- or 50-mm cells. Beer’s law is obeyed in the range 0-40 x g. Results and Discussion When nitrate was added to samples of the carminic acid - boron complex a strong change in the colour of the carminic acid took place after a few minutes. Depending on the nitrate concentration the colour became purple, blue or completely disappeared. This change of colour is seen in the change in the absorption graphs shown in Fig. 1. A nitrate concentra- tion of 250 mg 1-1 gave a clear blue colour, while at 2000 mg 1-1 of nitrate the solution was almost entirely decoloured.By adding different amounts of nitrate to solutions containing 10 mg 1-1 of boron the absorbance at 610 nm first decreased, then increased strongly for up to 350 mg 1-1 of nitrate. Further increases in the amount of nitrate resulted in decreased absorbances (Table I). Wavelengthhm Fig. 1. Spectral absorption curves for a carminic acid - boron complex measured against water. Nitrate contents: A = 0, B = 400, C = 1000 and D = 2 000 mg 1-1. Broken line: carminic acid. I t appears that the addition of hydrazine completely eliminates the interfering effect of nitrate, even at levels of 2000 mg l-l, if the equivalent amount of hydrazine is added. To secure an effective elimination of the interference hydrazine should be added in excess.In practice, 2 drops of hydrazine (concentrated) remove the interference of nitrate for up to 6000 mg 1-l. The reaction between sulphuric acid and hydrazine hydrate is violent, and the use of hydrazine should therefore be kept to a minimum. Up to 4 drops per 2 ml of sample solution gave satisfactory results, but losses on account of spattering occurred. TABLE I INFLUENCE OF DIFFERENT NITRATE CONCENTRATIONS ON THE ABSORBANCE OF THE CARMINIC ACID - BORON COMPLEX AT 610 nm, WITH AND WITHOUT HYDRAZLNE ADDITION Each sample contains 10 mg 1-l of boron. Nitrate concentration/mg 1-1 r * A 0 100 400 1000 2000 Absorbance without hydrazine . . . . 0.165 0.080 0.472 0.271 0.059 Absorbance with hydrazine . . . . 0.167 0.165 0.168 0.167 0.166 Recovery tests using model samples show that boron was recovered within a limit of 3%, Recovery tests in sewage effluents with nitrate rates as high as 1000Omgl-l (Table 11).October, 1979 SHORT PAPERS 985 TABLE I1 RECOVERY OF BORON AT VARIOUS LEVELS O F BORON AKD NITRATE I N MODEL SAMPLES Nitrate Boron/mg 1-' Average A added/ I ~-\ recovery, mgl-1 Added Found Added Found Added Found YO 0 2.00 2.00 4.00 4.00 10.0 10.0 100 2000 2.00 1.93 4.00 3.90 10.0 10.0 98 4 000 2.00 1.95 4.00 3.92 10.0 9.78 98 loo00 2.00 1.93 4.00 3.87 10.0 9.75 97 and nutrient solutions gave similar results (Table 111).The highest deviation seemed to occur at the highest levels of nitrate concentration. The nitrate content of nutrient solutions generally does not exceed 200 mgl-1 and sewage effluents seldom contain more than 200 mg 1-1 of nitrate. Concentrated nutrient solutions contain as much as 5000-10000mg1-1 of nitrate. The addition of hydrazine hydrate makes it possible to carry out a determination of boron with an error of *3%. The recovery was &3%. TABLE I11 RECOVERY OF BORON IN NUTRIENT SOLUTION AND SEWAGE EFFLUENT Sample Superba nutrient solution Sewage effluent . . . . Original Added concentrations/ concentrations/ mg 1-' mg 1-1 7- B NO, B NO, 0.50 100 2.00 0 4.00 1000 10.0 4000 0.45 145 2.00 0 4.00 1000 10.0 4000 B concentration Recovery, found/mg I-' Y O 2.45 98 4.41 98 10.18 97 2.40 98 4.50 101 10.76 103 References 1. 2. 3. 4. 5. 6. Hatcher, J . T., and Wilcox, L. V., Analyt. Chem., 1950, 22, 567. Callicoat, D. L., and Wolszon, J . D., Analyt. Chem., 1959, 31, 1434. Goldman, E., Taornina, S., and Castillo, M., J . Am. Wat. Whs Ass., 1975, 67, 14. Goulden, P. D., and Kakar, Y . P., Wat. Res., 1976, 10, 491. Ross, W. J., and White, J . C., Talanta, 1960, 3, 311. Lionnel, L. J., Analyst, 1970, 95, 194. Received March 12th, 1979 Accepted May 9th, 1979

ISSN:0003-2654    

DOI:10.1039/AN9790400983

出版商:RSC    

年代:1979

数据来源: RSC

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October, 1979 SHORT PAPERS 985 Spectrophotometric Determination of Nitrite Using 4,5- Di hyd roxycou m a ri n Motoshi Nakamura and Akira Murata Faculty of Engisceering, Shizuoka University, 3-5-1, Johoku, Hamamatsu-shi, Shizuoka-ken, 432, Japan Keywords : Nitrite determination ; spectrophotovnetry ; 4,5-dihydvoxycouuuzarin The hydroxyl derivatives of coumarin have been studied as analytical reagents for deter- mining metal ions,l but hardly any studies of the reactions with inorganic anions have been reported, except for 4-methylumbelliferone applied to the determination of itr rate.^,^986 SHORT PAPERS Analyst, Vol. 104 Anschiitz4 reported that 4-hydroxycoumarin was readily converted into oximidobenzo- tetronic acid by the action of nitrous acid. We found that 4,5-dihydroxycoumarin (DHC) also reacted with nitrite in acidic solution and that the reaction occurred considerably quicker than the reaction of 4-hydroxycoumarin with nitrite.The resulting yellow product was unstable in the acidic solution. However, it was found that in an organic solvent the same reaction readily took place on shaking the acidic solution containing the nitrite with the organic solvent containing dissolved DHC; the reaction product in the organic solvent was extremely stable. This reaction was applied to the simple and selective spectrophoto- metric determination of nitrite. Experimental Apparatus A Hitachi, Model 124, recording spectrophotometer was used for measuring the absorption spectra and a Hitachi, Model 139, spectrophotometer was used for the determination of nitrite.Reagents All chemicals used were of analytical-reagent grade. 4,5-Dihydroxycoumarin solution. DHC, which was prepared by the method described by Desai and Sethna,5 was dissolved in benzene containing 4% V/V of methanol, to give a 5 x M solution. Nitrite standard solution. A 1 x M solution of nitrite was prepared by dissolving sodium nitrite, which had been dried at 110 "C for 4 h, in water. Solutions of lower con- centrations were prepared by dilution of this standard solution. Procedure tube and add 5 ml of 5 N hydrochloric acid. and shake vigorously for 30s. sodium sulphate. Transfer 20 ml of a solution containing 0.14-3.50 p.p.m. of nitrite-nitrogen into a test- Then add 10 ml of the DHC in benzene solution Separate the benzene phase and dry it over anhydrous Measure the absorbance at 410 nm.Results and Discussion Absorption spectra for the product and for DHC in benzene are shown in Fig. 1. Maxi- Benzene is suitable as the solvent because it gives mum absorbance was shown at 410 nm. a low reagent blank, although chloroform or isoamyl acetate, etc., could be used. 0.6 0 400 450 ' 500 Wavelength/nm Fig. 1. Absorption spectra. A, Absorption spectrum of nitrite - DHC product ueYsu.s reagent blank, nitrite-nitrogen = 1.4 p.p.m.; and B, absorption spectrum of reagent blank veYsus water. The effect of the acidity of the aqueous phase on the absorbance was examined (Fig. 2). With hydrochloric acid a constant absorbance was obtained in the range 0.25-2.0 N, butOctober, 1979 SHORT PAPERS 987 "'6 0 1 2 3 4 5 ( a 1 I 1 I I - 0 2 4 6 8 10 12 14 (b) Acidity/N Fig.2. Effect of acidity. The acidity of the aqueous phase containing 1.4 p.p.m. of nitrite-nitrogen was adjusted by addition of hydrochloric acid [line A, axis of abscissa, ( a ) ] ; perchloric acid [B, ( a ) ] ; and sulphuric acid (C, ( b ) ] . with sulphuric and perchloric acid the absorbances varied with the acidity and the absorbance maxima were obtained a t 5 and 1.5 N, respectively. The effect of the reagent concentration was also examined. The maximum and constant absorbance was obtained above 1 x The addition of methanol to the benzene was necessary in order to prepare a DHC solution of the recommended concentration (5 x 10-3 M) , because the solubility of DHC in a solvent such as benzene is low.Variation of the amount of methanol added did not influence the absorbance in the concentration range 4-20y0 V / V . The maximum absorbance was reached after shaking for 5 s and was constant up to 60 s; shaking for longer than 60 s slightly reduced the absorbance. M of DHC in benzene. Fig. 3 shows the effect of the shaking time for the reaction on the absorbance. I , 1 60 120 180 Timeh Fig. 3. Effect of shaking time for reaction. Nitrite- nitrogen = 1.4 p.p.m.; DHC = 6 X w 3 M . At 410 nm the nitrite - DHC system follows Beer's law over the concentration range 0.14-3.50 p.p.m. of nitrite-nitrogen. The absorbance remained unchanged for at least 24 h. A statistical study of seven samples, each containing 1.4 p.p.m. of nitrite-nitrogen, which gave a mean absorbance of 0.400, was carried out over a period of 2 weeks and gave a relative standard deviation of 1.0%.The effects of various interferences were investigated on a sample containing 1.4 p.p.m. of nitrite-nitrogen. Most of the diverse ions examined were tolerated when present in large amounts. However, iodide, sulphite, chromium(V1) and iron(I1) interfered at the 10 p.p.m. level. The reaction product that is utilised as the analytical species is thought to be a hydroxyl derivative of oximidobenzotetronic acid, although it has not been identified. The results are shown in Table I.988 SHORT PAPERS TABLE I EFFECT OF DIVERSE IONS Ion Nitrate, acetate, chloride, phosphate, carbonate, bromide, perchlorate, ammonium, iron(III), magnesium, tin(IV), cobalt, lead, zinc and calcium . . .. .. . . Sulphate and aluminium . . . . . . .. .. Fluoride . . . . . . . . .. . . .. Iodate, nickel, cadmium and chromium(II1) . . .. Copper . . . . . . . . . . . . .. .. Iodide, sulphite, chromium(V1) and iron(I1) . . * . Tolerance limitp p.p.m.* 50 000 2 5000 12 600 5 000 2 500 < 10 * Amount of ion causing an error of less than 5% in the determination of 1.4 p.p.m. of nitrite-nitrogen. References 1 . 2. 3. 4. 5. Katayl, M., and Singh, H. B., Talanta, 1968, 15, 347. Skujins, J . J., Analyt. Chem., 1964, 36, 240. Keil, R., 2. Analyt. Chem., 1974, 271, 359. Anschiitz, R., Justus Liebigs Annln Chem., 1909, 367, 169. Desai, N. J., and Sethna, S., J . Org. Chem., 1957, 22, 388,. Received April 9th, 1979 Accepted May 1&h, 1979

ISSN:0003-2654    

DOI:10.1039/AN9790400985

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Analyst, October, 1979 Book Reviews 989 ENZYME LABELLED IMMUNOASSAY OF HORMONES AND DRUGS. Edited by S. B. PAL. Proceedings of the International Symposium on Enzyme Labelled Irnrnunoassay of H orrnones and Drugs, Ulm, West Germany, J u l y 1 0 and 11, 1978. Pp. xxvi + 475. Berlin and New York: Walter de Gruyter. 1978. DM130. This timely book is probably the best publication to date on the development and application of enzyme immunoassays. It consists of a collection of papers delivered at an International Meeting on Enzyme Immunoassay held a t Ulm, West Germany, in 1978. The value of the publication is that it covers both the basic and applied aspects of the subject, so this work should be of interest to both the researcher and the user. A few papers are devoted to new homogeneous assays, e.g., for oestriol and to established assays of this type, e.g., for thyroxine and anti-epileptic drugs.However, the majority of the contribu- tions deal with heterogeneous enzyme immunoassays. All aspects of these are dealt with in depth including technical procedure and reagent preparation. The use of various separation methods, such as double antibody, and solid-phase techniques are explained and the importance of the solid phase used is stressed. Different groups describe tests in tubes, beads and microplates made of polystyrene, nylon and other materials. The enzyme label and its conjugation to the immunological reactants is one of the most important topics in this book. Enzymes mentioned include glucose oxidase, peroxidase, /3-galactosidase, penicillinase and malate dehydrogenase and labelling methods mentioned include periodate and maleimide methods as well as the use of glutaraldehyde and bisfunctional reagents.Much attention is also devoted to the measurement of the enzyme activity by chromogenic and fluorogenic substrates as well as by chemiluminescence and thermometry. The laboratory and clinical use of enzyme immunoassays for both steroid and protein hormones and for drugs is comprehensively reviewed. The contents of this book indicate that a vast amount of research and development is going into enzyme immunoassays and suggests that they might soon take their proper place in the routine laboratory. The reviewer heartily recommends this publication. A. VOLLER MANUAL ON WATER. Fourth Edition.Edited by C. E. HAMILTON. ASTM Special Tecknical Philadelphia, Pa: American Society for Testing and Publication 44214. Materials. 1978. Price $29.50. Pp. vi + 472. To quote from the introduction “This manual is intended as a brief reference source of information on water. It will not replace an adequate library on the subject, but it does provide basic informa- tion for routine use and cites references to the technical literature, thus serving as a point of depar- ture for more specific and detailed studies.” Given these aims, the book is a useful source of information though much of the text is concerned with examples and procedures from the USA and, in particular, ASTM publications. For example, in the brief discussion of sources of water, all numerical data are for the TJSA. Similarly, the extensive discussions of sampling and analytical procedures are centred around standard ASTM methods.Nevertheless, the book does provide a useful overview of the field for readers from other countries, and would be a worthwhile addition to the reference shelves of libraries for those requiring a brief account of the main aspects of the many topics included. A list of the chapter titles is useful in indicating the topics covered : “Water Sources and Supply” (15 pp.) ; “Uses of Industrial Water” (10 pp.) ; “Production and Preservation of Ultrapure Water” (10 pp.); “Effects of Composition on Industrial Use” (8 pp.); “Treatment of Process Water and Waste Water” (38 pp.) ; “Technology of Industrial Water Re-use’’ (13 pp.) ; “Self-purification and Other Natural Quality Recovery Mechanisms” (10 pp.) ; “Thermal Loading of Water Supplies” (11 pp.) ; “Sampling and Flow Measurement of Water” (27 pp.) ; “Water Quality Monitoring” (32 pp.) ; “Analysis of Water and Waste Water” (77 pp.) ; “Sensory Examination of Water” (22pp.) ;990 BOOK REVIEWS Analyst, Vol. 104 “Sampling and Identification of Deposits in Steam and Water Systems” (31 pp.) ; “Chemical Analy- sis of Deposits” (12 pp.) ; “Sediment in Streams and Other Water Bodies” (8 pp.) ; “Radioactive Nuclides in Water” (25 pp.) ; “Nuclear Water Technology” (35 pp.) ; and “Practices for Measure- ment of Radioactivity” (56 pp.). There is also an Appendix giving Reference Tables and curves. Note, though, that biological sampling and examination procedures are not included.A. L. WILSON MONITORING TOXIC SUBSTANCES. Based on a symposium sponsored by the ACS Division of Industrial and Engineering Chemistry a t the 174th Meeting of the American Chemical Society, Chicago, Illinois, August 31, 1977. Edited by DENNIS SCHUETZLE. ACS Sympo- s i u m Series 94. Pp. xii + 290. Washington D.C.: American Chemical Society. 1979. Price $26.50. This book is based on presentations a t a symposium sponsored by the ACS Division of Industrial and Engineering Chemistry a t Chicago in August 1977. Although the symposium was held about two years ago there has been a commendable effort to update the references to include work published since the symposium. In addition, several papers of relevance, but not presented a t the symposium, have been added.The papers bring together various disciplines associated with toxic substances, in particular biological screening and chemical examination. The first chapter is a discussion of the develop- ment and use of the “Ames Test” by the originator, a most appropriate introduction to a book of this nature. Other techniques for testing compounds for mutagenicity and carcinogenicity are described in the second chapter. There are several contributions concerned with the analysis of organic pollutants in waste water and drinking water. The importance of appropriate sampling is stressed and analytical procedures using gas and liquid chromatography, the former also combined with mass spectrometry, are presented. The need for a satisfactory combined high-performance liquid chromatograph and mass spectrometer is clear.Two chapters are devoted to the detection of trace metals in air and the two usual methods, namely atomic-absorption and inductively coupled plasma - atomic-emission spectroscopy are both discussed. The various procedures for surface analysis as applied to atmospheric particu- lates are covered. Less common procedures for gas analysis include Fourier transform - infrared spectroscopy and optoacoustic spectroscopy. Using Fourier transform - infrared spectroscopy, mixtures of both inorganic and low relative molecular mass organic materials have been quantitatively analysed, with an excellent limit of detection. The use of ion chromatography (basically high-performance liquid chromatography) is important in the context of pollution as many toxic materials are ionic.The principles of ion chromato- graphy are described and an extensive list of ions for which this technique is applicable is given. Examples are given of the analysis of water and air for these pollutants. The final chapter is devoted to the NIH/EPA Chemical Information System, with emphasis on its value to analytical chemists. The book is well presented with clear typescript and figures, and its value is considerably en- hanced by the inclusion of a single index covering all chapters. T. A. GOUGH AIR POLLUTION REFERENCE MEASUREMENT METHODS AND SYSTEMS. Proceedings of the Inter- Edited by T. SCHNEIDER, H. W. Pp. viii + 168. national Workshop, Bilthoven, December 12-16, 1977. DE KONING AND L. J.BRASSER. Amsterdam, Oxford and New York: Elsevier. 1978. Price $35.55; Dfl180. Studies in Environmental Science 2. This work is the proceedings of the international workshop held in Bilthoven in 1977, organised jointly by The National Institute of Public Health, Bilthoven, and the World Health Organization. Of the 44 participants only 5 are from outside Europe and 25 are from The Netherlands. The work is prefaced by The Prince of The Netherlands who chaired some of the forum. A general introduction is followed by 17 papers and a report summarising the proceedings. Several papers are devoted to the philosophy of reference methods. The purpose of reference methods is discussed, as are the principles on which such methods are based, including the use of reference materials and the evaluation of methods.Several speakers pay attention to the calibra-October, 1979 BOOK REVIEWS 991 tion of. methods and inter-laboratory comparisons. Applications of reference methods are con- sidered for the measurement of dust, the use of portable units and suitable systems for use in the developing countries. One paper gives an account of the economics of a calibration unit in The Netherlands. Thus, accounts are presented of work in USA, Japan, India and Germany. On the larger scale there are reviews of the European Monitoring and Evaluation Programme (EMEP), “Euregios,” i.e., areas boarded by several other countries, the Pan American Air Monitoring Network (REDPANAIRE) and the Global Environmental Monitoring System (GEMS). There is also a contribution on standard- isation in meteorology.I t is recognised that some countries already have expensive monitoring systems whilst manual methods are more appropriate to poorer nations. The first is an excellent discussion of the generation of standard atmospheres of gas mixtures using permea- tion tubes. A further paper introduces a new technique of analysing aerosols for mass size distribu- tion with an apparatus fitted with piezo-electric quartz sensors. The final paper compares the various methods of measuring particulate matter. The book is reproduced from typescript (mostly double spaced) and is well illustrated, but the papers contain few references. Although nominally 168 pages, 19 pages are taken up with the list of participants or are blank. The main merits of this work are : a sound presentation of the principles of reference methods ; several excellent technical papers ; and an interesting account of air monitoring in several countries and larger scale projects.In the whole, the work lacks depth and contains quite a lot of duplication and one feels it is an expensive buy. A number of speakers gave reviews of national activity in air-pollution measurement. Three papers are more specifically technical in nature and these are the most useful. Thus, I consider the work very expensive. S. CRISP SULFUR IN THE ENVIRONMENT. PART I : THE ATOMSPHERIC CYCLE. PART I1 : ECOLOGICAL IMPACTS. Environmental Science and Technology Series. Part I, New York, Chichester, Edited by JEROME 0. NRIAGU. Pp. xiv + 464 + 8-page errata pamphlet; Part 11, Pp.xii + 482. Brisbane and Toronto: John Wiley, 1978. Price: Part I i 2 3 ; Part I1 L24. These volumes cover a wide range of aspects of sulphur pollution. In assessing such a compre- hensive work my approach was to consider a three-fold question: what value are the volumes to the analytical chemist, the general chemist and the non-chemist ? This is an excellent chapter, one of the best in the volumes. The analyses of sulphur dioxide, hydrogen sulphide and organic sulphur compounds by continuous monitors and manual procedures are reviewed. Detail is avoided but it is difficult to prevent the treatment becoming a catalogue. The discussion of the analysis of particulate sulphur compounds is fascinating and topics covered include determination of oxidation state, speciation, size distribution and problems of sampling on filters.Other chapters of interest to the analytical chemist include the first on “Production and Uses of Sulphur,” a useful review of the possible origins of sulphur compounds. Three chapters discuss various aspects of factory emission, dispersal, transport and deposition of sulphur compounds and there is some duplication of treatment. The main criticism of these chapters is the failure to introduce specialist (e.g., meteorological) terms, and the first even uses undefined symbols and has unlabelled diagrams. There is an interesting contrast of styles between the following two chapters. The former is far too long and mathematically detailed, tensor calculus is assumed, whilst the latter is readable although containing substantial mathematics.The chapter on acid precipitation contains much information and concentrates on ecological effects but with a somewhat confused discussion of pH. Chapters of more general chemical interest are: Part I, “The Global Sulphur Cycle,” “Organo- sulphur Emissions from Industrial Sources” and “Atmospheric Chemistry of Sulphur-containing Pollutions” ; and in Part I1 “Deteriorative Effects of Sulphur Pollution on Materials,” “Chemistry of Pollutant Sulphur in Natural Waters” and “The Acid Mine Drainage.” The last is far too long and detailed. Especially interesting are those chapters on materials deterioration and natural waters although the latter would be difficult to follow without some knowledge of chemical thermo- dynamics and reaction kinetics.Other subjects cover cost/benefit analysis of emission control, effects of sulphur compounds on animals, human health, plants, aquatic ecosystems, soil and microbes. A large variation of Only one chapter in this work is specifically devoted to chemical analysis.992 BOOK REVIEWS Analyst, Vol. 104 standards were encountered. A number of authors made no attempt to explain specialised terms, something essential for the success of a multi-disciplinary work. The chapters concerning plants were especially poor comprising tedious detail ; however, readable accounts were presented on human health effects and soil pollution. Generally, the chapters are well referenced (1977) and illustrated and contain much data but a number of editorial errors were found.The analytical chemist will find little more on method- ology than in other recent works in this field. The main strength is the presentation of a broad account of sulphur pollution, unfortunately marred by specialists unwilling (or perhaps unable) to introduce their subjects to other scientists outside their own expertise. S. CRISP MODERN METHODS FOR TRACE ELEMENT ANALYSIS. By MAURICE PINTA. Pp. xii + 492. Ann Arbor, Mich. : Ann Arbor Science Publishers. 1978. Price klS.60; $32.45. This book starts with an introductory section in which the author explains his choice of methods, which, in the opinion of this reviewer, is not unexceptionable. Many would regret the absence of molecular-absorption spectrometry and electrical methods on grounds that are at least debatable.There follow chapters in which the chosen methods are discussed. Fluorescence is given a reasonable treatment and a number of useable methods are given. How- ever, the chapter is by no means self-contained and i t would be desirable to refer to the original text of the numerous references before applying the technique. After an adequate treatment of the theoretical principles, there follows an account of methods and apparatus where this slightness of treatment is in evidence. The section on sources is marred by historical inaccuracy and wrong emphasis, dates without references and a t least one incorrect reference. It is possible to speculate on the absence of furnace methods in this chapter on emission, as they are included in a later chapter on atomic absorption. I t contains a worth- while section on interferences. There are many references and actual applications. Atomic-fluorescence and X-ray fluorescence spectrometry chapters have the same lack of detail as that on emission; the chapter on activation analysis is rather better. I t is difficult to decide a t whom this book is aimed. A student would find it valuable as it contains concise, if limited, accounts of important techniques. He may find the price prohibitive. A non-specialist analytical chemist, whose knowledge of trace methods needed revision, would find the book of value but would need to refer to the original specialist texts if he wished to apply the techniques. A scientist, whose discipline was other than analytical chemistry, but who wished to acquire a knowledge of trace methods, would also find the book useful. An expert in the field of trace analysis would find the treatment too superficial. This superficiality is even more evident in the chapt'ers on emission. The chapter on atomic absorption is well written and complete in itself. The section on non-flame (furnace) methods is useful. S. GREENFIELD

ISSN:0003-2654    

DOI:10.1039/AN9790400989

出版商:RSC    

年代:1979

数据来源: RSC