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The chemistry of cadmium in natural water—I a study of cadmium complex formation using the cadmium specific-ion electrode
Water Research Vol. 8, pp. 23 to 30. Pergamon Press 1974. Printed in Great Britain
THE CHEMISTRY OF C A D M I U M IN NATURAL W A T E R - - I A STUDY OF C A D M I U M COMPLEX F O R M A T I O N USING THE C A D M I U M SPECIFIC-ION ELECTRODE J. GARDINER Water Pollution Research Laboratory, Department of the Environment, Stevenage, England (Received 16 June 1973)
Al~tract--The extent of formation of labile complexes of cadmium has been investigated in synthetic solutions and in real samples. A substantial proportion of the total cadmium in river and lake water will usually be present as the free cadmium ion, and this proportion will be larger the lower the pH value and the lower the proportion of sewage emuent present in the water. Humic substances usually account for most of the complexation, followed in importance by carbonate, the complex of which with cadmium has not been reported previously.
cadmium, at the /ag 1-' level, and Matson and coworkers have devised methods for measuring the extent of formation of labile and non-labile complexes (Matson, 1968; Matson, Allen and Rekshan, 1969). For the purposes of this work, however, formation of labile complexes was found not to give sufficiently large changes in signal. The cadmium specific-ion electrode is extremely useful in a study of complex formation since it responds only to free cadmium ion in solution and not to complexes. Consequently, if the total cadmium concentration is known, the proportion complexed can be calculated. Its major disadvantage is that the minimum concentration at which it can be used is about 0.1 mg 1-t, which is considerably higher than the concentrations usually encountered in natural waters, though extrapolation to lower concentrations could be made with reasonable certainty. Measurements were made a few minutes after the addition of cadmium to a particle-free, filtered solution at concentrations below those at which precipitation would occur. Problems arising from adsorption and from slow reactions which may have occurred in some natural media were, therefore, minimized. Because of the comparatively small number of colorimetric reagents for cadmium it was not pbssible to devise an analytical scheme like that for copper (Stiff, 1971b). A considerable amount of data on complexation of cadmium already exists (Sill~n and Marteil, 1964, 1971). These data have been used by Mangel (1971) to predict complexation in sea water and by Hem (1972), and it is clear that a wide range of inorganic and organic ligands are potentially important. Measurements were
INTRODUCTION
The work described is a continuation of research in this Laboratory on complex formation of metal ions in natural waters (Stiff, 1971b). Metal complex formation could have important consequences both with regard to toxicity to aquatic organisms, because the toxicity of complexes can be different from that of the free ion, and also in analysis. Complex formation also affects the rate and extent of the disappearance of the metal ion from solution by precipitation and adsorption. Knowledge of the total metal concentration alone is therefore not always sufficient to enable toxicity and other properties to be predicted. Cadmium was chosen as the next metal for investigation after copper because of the current interest in its distribution throughout the environment and because it is extremely toxic to fish, effects on growth rate being observed even for concentrations between 5 and 10 #g 1- ' (Brown and Shurben, 1973). The toxicity curve for cadmium to rainbow trout, i.e. the curve of median survival time against concentration, shows a plateau region from 20 #g 1- t to 1"0 mg 1- t of cadmium over which the median lethal response time is relatively independent of concentration (Ministry of Technology, 1968). There are several methods available for measuring total cadmium concentration (e.g. atomic absorption spectroscopy and spark source mass spectrometry) but none of these would be suitable for measuring complex formation. Anodic stripping voltammetry is capable of measuring concentrations of some metal ions, including 23
24
J. GARDINER
made only on those systems for which there was no information or for those which, although existing data indicated the possibility of complexation, there was inadequate information at the ionic strength of interest. EXPERIMENTAL Materials I. Cadmi,~m solutions. Dilute stock solutions were freshly made up each day from a more concentrated solution of BDH reagent grade cadmium nitrate. 2. Humic acid solutions. Commercial humic acid (Aldrich Chemical Co.) was partially dissolved in sodium bicarbonate solution, filtered, and extracted by nhexanol. The organic extract was then evaporated down. This procedure was necessary to reduce contamination by metal impurities, notably zinc. 3. Other materials. All other substances were BDH AR grade unless otherwise stated. Doubly distilled water was used. 4. Ri~'er water and sewage samples. Samples of crude sewage, sewage effluent and river water were first allowed to settleifnecessary, then filtered under pressure through GF/A, GF/C, and 0.45/Jm membrane filters successively. One river-water sample contained sufficient cupric ion to contaminate the cadmium specific-ion electrode. A short column of Chelex I00 resin (Bio Rad Laboratories) was prepared by passing down it a concentrated phosphate buffer solution at pH 7, followed by thorough washing with distilled water until the washings failed to give a precipitate with silver nitrate solution. Passage of the filtered river-water sample down this column removed the source of contamination of the specific-ion electrode and left the pH value of the sample practically unaltered. Equipment Measurements were made in a cell containing approximately 100 ml of sample. The cell contained an Orion 94/48A cadmium specific-ion electrode, a silversilver chloride reference electrode, and a glass electrode with internal reference connected to a Radiometer No. 26 pH meter. The potential between the specific-ion and Ag/AgC1 electrodes was measured, using a similar pH meter, to within +0-3 inV. A mixture of nitrogen and carbon dioxide could be bubbled through or blown over the surface of the solution, and there was magnetic stirring in addition. Procedure A known quantity of cadmium nitrate (usually between 0"l and l'0 mg Cd l - I ) was added to the sample. The amount of cadmium already in the sample
was insignificant by comparison with the amount added. When required, potassium nitrate was added to adjust the ionic strength approximately to that of a hard natural water. A period of 15 min was allowed for the electrode system to reach equilibrium, which was probably longer than necessary. Experiments were carried out at room temperature (20-0 + 1.0~C). In most cases the pH value was controlled by varying the composition and flow rate of the mixture of nitrogen and carbon dioxide admitted to the cell. In the case of the hydroxy complex, however, pH values greater than 7 were achieved by slow addition of 0-04 M NaOH from an Agla micrometer syringe. CO_, being excluded from the cell. A reverse titration was then carried out with 0-04 M HCI. In order to determine the extent of complexation it was necessary to know the potential corresponding to no complexation. This could be determined in some cases by lowering the pH value until the potential was independent of pH value, or by measuring, before and after measurements on the sample, the potential developed in a non-complexing solution with the same total cadmium concentration and the same ionic strength. Determination of bicarbonate, chloride and sulphate concentrations Bicarbonate and chloride concentrations in unknowns were measured by the conventional methods, i.e. for bicarbonate, titration with N/50 hydrochloric acid using bromo-phenol blue as indicator, and for chloride, titration with silver nitrate solution, using potassium chromate as indicator. Sulphate was determined turbidimetrically by reaction with barium chloride (American Public Health Association, 1971). Determination of humic concentration An approximate measurement of the concentration of humic material in a filtered river-water or sewageeffluent sample was obtained from the absorbance at 400 nm using an SP.800 spectrophotometer, by comparison with solutions of known composition.
RESULTS AND DISCUSSION Behaviour of the specific-ion electrode Below pH values at which formation of the hydroxy complex was significant (see below) down to pH 5 and probably below, the potential was found to be independent of the pH value. Readings were very reproducible provided that the system was allowed to reach equilibrium, but rapid changes of pH value did cause a temporary departure from the equilibrium value.
The chemistry of cadmium in natural water--I
25
Table 1. Potential between the cadmium specific-ion electrode and Ag/AgCI reference electrode in a 1.0 mg 1-1 cadmium solution (8.9 x I0 -~ M) in 4 x I0-* M NaHCO 3 at pH 7.5 at various ionic strengths Ionic strength (KN03) (raM) Negative potential (mV)
2'0 178.1
4-0 6.0 8.0 10'0 12-0 14-0 18"0 20"0 180-0 181.4 182.5 183.4 184.1 185.0 186.0 186-8
Theoretically the potential, E, follows the relation E = A + Blogacd_,+ =
A + B Iogfcd.,- + BlogECdZ÷],
(1)
where A and B are constants and acd:" and fed:- are the activity and activity coefficient of the cadmium ion respectively. Plots o r e against log I-Cd-"+], with [Cd z +] from 1.0 x 10 -4 to 2.0 x 10 -6 M in 2 x 10 -3 M KNO3 were straight lines. The slope was 30 (-I- 1) mV for each 10-fold increase in concentration, which is not inconsistent with the theoretical value of 29-6 mV for B in equation (1). The form of equation adopted, therefore,
0"5 mg 1- ~, with the total cadmium concentration less than 1-0 mg 1- t the value orB decreased below 29.6 mV. Consequently, whenever possible, [Cd]r was kept at 1.0 mg 1-t or above. Contaminationof the electrode by cupric ion caused a change in the characteristics of the electrode, even after the source of the contamination had been removed. It was then necessary to lightly clean the crystal with fine emery cloth. Let [Cd]r be the total cadmium concentration and E0 the potential when there is no complexation. From equation (2) the potential change on complexation, Eo - E, is given by
was
E = A ' + BlogECd2+],
where B = 29'6 mV and the value of A' varies with the ionic strength of the solution (see Table 1), and also from day to day at constant ionic strength over a range of about 8 mV. The response of the electrode to log [Cd 2+] was linear down to 0-5 mg 1- ~ but the value of B in equation (2) then began to decrease (see Fig. 1) until below 1.0 #g I- t the signal was independent of cadmium concentration. The specifications of the electrode state that it is sensitive to [Cd z+] down to 10 ngl--t provided that the total cadmium concentration, [Cd]r, is greater than 10 ,ug I-t. It was found, however, that when complex formation occurred so as to lower [Cd-" +] below about
19(
Eo - E
(2) but
:
B
log
[Cd]~
[Cd:']'
[Cd]r = [Cd 2 +] + l e d complex],
(4)
where [Cd complex] is the concentration of cadmium complexed, assuming no losses by precipitation or adsorption. Using equations (3) and (4) the ratio [Cd complex]/[Cd-" +] can be calculated from E o - E.
Complexation by individual ligands Hydroxide. A comparatively low cadmium concentration of0"l mg 1-t was used so that a pH value of 10 could be reached without precipitation of cadmium hydroxide. All potential ligands except hydroxide were absent. The concentration of nitr.ate (2 x 10 -3 M) was
+ txpmu,4cm'~ ~04~Ts
Z
m 21(
~
Z~
+ .........
-4
), I"
24C
2SC
(3)
l t % ~ O ~ N T ~ T , O N . (,~n) ~0'
0~
Fig. 1. Variation in potential with concentration of the cadmium specific-ion electrode.
:6
J. GARDINER
too small to give an)complexation (Sill6n and Martell, 1964). On addition of sodium hydroxide the potential remained constant up to pH 7.4. but above this value it began to d:crease.
known to occur at high carbonate concentrations (f13 = 1.7 × 10 6 M -3, Lake and Goodings, 1958). It is known that soluble complexes of C a : " and Cu-' " exist (Garrels and Thomson, 1962; Stiff. 1971a) and their £d(H:O)~" ~ CdOHtH.O)~ + H*. (5) stability constants are similar to those of oxalate complexes with the same metal ions (Sill~n and Martell, Let 19641. The stability constant of the cadmium oxalate complex is given variously as 10s5. 103.9 and 10"t° M - t FCdOH(HzO);] ~tH "',- [ C d O H ' ] {H~} by Sill~n and Martell (1964), and the stability constant K°H = [Cd(H:O)~*] = [Cd z +] ' of a carbonate complex should be in this region. (6) Experiments were carried out by measuring the variawhere Cd(H.,O)~ ÷ and CdOH(H,O)~ are abbreviated tion in the extent of complex formation with pH value to Cd" ~ and CdOH + respectively. in a pH range in which the extent of formation of CdOH + was low. in the presence of various concentrations of sodium bicarbonate. The results, shown in Log [CdZ+] = l o g K o . + pH, (7) Fig. 3, indicate 1:1 complex formation involving one proton per molecule and are consistent with the where pH = - l o g [ H + ~,. The readings were stable, and on adding acid the following reaction: potential at a given pH value was shown to be indepenCd '~ + HCO~ ~ C d C O 3 + H + (8) dent of whether the pH was rising or falling, i.e. the change was completely reversible. A plot of log [CdOH +]/[Cd: +] against pH value was a straight line where CdCO~ is a soluble complex. From a knowledge (Fig. 2), as predicted by equation (7), but the slope was of the acid association constant of the carbonate ion 0-84 instead of 1.0 as expected. This was probably caused [Hco{] by the approach to the detection limit of the electrode. K . = [ C O ~ - ] (H +} (9) From Fig. 2, log K o , was calculated to be -9.06 (_+0.08) at 20°C, ~ = 2 x l0 -3 M (KNO3). Carbonate. A l : I carbonate complex of cadmium has ( K , --- 2-0 × 10 '° M - t , Sill~n and Martell. 1964)it is not been reported previously, although Cd(CO3) ~- is possible to calculate values for Kco~ the stability constant of the carbonate complex,
[CdOH÷2
0
//
[CdCOa] Kco,=(Cd,.] [COF]'
(10)
0.4
0.2
a
9
0
I
tJ.
-o.2
L
pH vAt.U(
-0"4 U
S-0., / ) ) -I'0
- 1-2
- 1.4
-I-6
Fig. 2. Graph of log[CdOH*]/[Cd-'*] against pH value. Total cadmium concentration. 0"10 mg 1- ~.
as respectively 10't°3, 10'~'14 and 104.0-' M - t at 2"5, 5"0 and 10.0 mM H C O j . Of these values the last is the most accurate. The solubility product of CdCO 3 is well documented. Previous work in this Laboratory (Stiff and Gardiner, 1973) has shown that in a clear solution the system is slow to reach equilibrium. In the experiments described here, with [Cd]r at 0-9 mg 1-*, and at the higher bicarbonate concentrations, the solubility product was frequently exceeded. Precipitation occurred only in some of the preliminary experiments and reversible changes then followed, cloudiness of the solution eventually becoming evident. The results of Fig. 3 could not have been due to partial precipitation because changes in potential with pH value were reversible and reproducible, unlike the effects observed with precipitation. Furthermore, the observed extent of complexation was, within experimental error, independent of [Cd]r, whereas the extent of precipitation would be dependent on [Cd]r.
27
The chemistry of cadmium in natural water--!
+÷
= 0.Zl-
,/
0
-0-~
?-
.
'
'
~"
//s//s /
, •
/
-0"4
INtos"],,s,,u
/
" [ H C O ! " ] I tO a M
-0-6
s
/s
//
7, II
/.
+
[.cos-l.=.s../
i/I
.,'
,
//~'HEOm[TICAL
/
t.co:l-o
Fig. 3. Effect of pH value on log(led complex]/[CdZ+]) in bicarbonate media. Total cadmium concentration, 1"0 x 10-s M (0.9 mg I-t). There was therefore strong evidence for a soluble carbonate complex with log Kco~ = 4.02 (+0.04) at 20'C, # = 10 x 10-3 M, in good agreement with the prediction based on comparison with the oxalate complex. Chloride and sulphate. Values given by Sill~n and Martell (1964) for the stability constants of the 1:1 complexes of chloride and sulphate with cadmium indicate that these complexes can be expected in natural waters. The extent of complexation was consequently measured in solutions containing various concentrations of sodium chloride and sodium sulphate respectivelyat an ionic strength typical of a hard natural water, and with a total cadmium concentration of i~) nag 1- t.
[CdCV]
slopes of which gave, respectively, Ko = 48 (+8) M - t and Kso, = 220 (+ 10)M- t both at # = 10 x 10 -3 M, and 20°C. The extent of complex formation was, of course, independent of pH value in both cases. Humic material. Because "hurnic acid" is a mixture and its mol. wt and precise acid-base properties are unknown, it is not possible to establish stability constants for complexation of metals in the usual way, nor is it certain that a mixture of humic material obtained from one source will behave in the same way as that from another. Solutions were made up containing 86, 20 and 6-6 mg 1-1 of humic acid purified as described, representative of a high, moderate and low level of coloration respectively. Each solution contained 3 x 10 -4 M bicarbonate, which was sufficient to allow the pH value to be varied over the range of interest by varying the flow of the N,/CO., mixture to the cell, but not sufficient to give appreciable carbonate complex formation. The fraction complexed was found to be slightly dependent on pH value at all three concentrations of humic material, and the ratio of complexed to uncomplexed cadmium ion did not increase linearly with humic acid concentration (see Table 2).
(II)
Ko = [Cd2 +] [ C l - ] ' [CdS04]
(12)
tqo, = [Cd2+] [SO=_]. In both cases plots of [Cd complex]/[Cd' +] against the concentration of ligand were straight lines, the:
[Cd complex]
['Cdz+]
oc
[ligand] °6'~
{H+} °'2° •
The low p H dependence could indicatethat the ligand is largely unprotonated in the p H range studied or, more probably, that the fraction of humic material responsible for most of the complexation is a strong chelating ligand relativelyunaffected by pH changes. The ligand dependence shows that some of the ligand molecules complex more than one cadmium ion, which is consistent with a polymeric ligand structure and perhaps with polymerization of the complex. Results were independent of cadmium concentration over the range [Cd]r 0-I-I'0 mg I-', afterallowing for changes caused by the approach to the detection limit, and assuming the average tool. wt of the humic material to be about 700 (Schnitzer, 1971), the ligand was in excess in all cases. A convenient formula based on equation (13) for calculating the degree of complexation by humic
Table 2. Dependence of the degree of complexation of cadmium by humic substances on pH value pH value
6.0
6"5
7.0
7.5
8-0
8"5
2.50 1.00 0.50
3-20 • 1.26 0-62
[Cd complex]/[C-d 2+] Humic concentrations (rag I-t)
86 20 6"6
1"00 0.40 0-20
(13)
1-26 0-50 0-25
1-59 0.62 0.32
2-00 0.77 0"40
28
J. GARDINER
Table 3. Composition of samples used in measuring extent of complexation of cadmium
Sample •
No.
Type ofsample Fresh ground-water* Filtered settled sewage'l" Filtered poor quality percolating filter efltuent Filtered good quality percolating filter ettluent Filtered river water Filtered river water
7
Filtered river water
8
Filtered river water
9
Filtered river water
10
Filtered river water
11
Filtered river water
Point of collection
Absorbance Approx in40-mm ionic cell at strength 400nm (mM)
pH value
[HCO~] (mM)
[CI-] (mM)
[SO~-] (mM)
WPRL WPRL
7.2 7-7
5" 1
0.39 2'8
0'030 0-34
0
8.4
0-22
17
WPRL Filter 30
8-0
8'2
3-0
0"31
0.22
!7
WPRL Filter 1
8'0
5-8
3'1
0"33
0.16
14
Stevenage Brook WPRL R. Ivel, Stotfold, Herts Pix Brook, Arlesey, Herts R. Hiz, Arlesey, Herts R. Lee, Luton++ R. Lee, Luton§ R. Trent, Swarkeston¶
7-7
4-4
1.8
0"59
0-07
l1
8"1
5.2
0-94
0.20
0.01
9
8"1
5"5
2-8
0-44
0.07
12
8.1
5.2
1'5
0.36
0.03
Il
8"0
4. I
3.7
0-40
0"02
12
7-9
4.8
3-6
0.68
0:05
14
7.7
2.6
1'3
0-35
0"06
8
8
* Used in fish toxicity tests. ? +2 drops I00 vol H20 2. :~Upstream of Luton sewage works outflow. § Downstream of Luton sewage works outflow. ¶ Four miles south of Derby. material from the absorbance of the solution is log
[Cd complex] [Cd 2+] 0"64 log A4oo + 0.20 pH - 1.14, 04)
where A4oo is the absorbance of the filtered solution at 400 nm in a 40-ram cell, assuming ferric ion and other interfering substances are absent or negligible.
Complexation in samples of river water and sewage effluent. The source, type and composition of the sewage and river water samples used are given in Table 3. From the concentrations of the ligands present, using the measured stability constants and equation (14), it was possible to calculate the proportion of cadmium present in complex forms and as the free ion (see Table 4). The results of a direct determination of the extent of complexafion, obtained by adding 0.1-10mgl - t of cadmium to the sample and measuring the proportion remaining as free ion, are also given in Table 4 for
comparison. The fact that there was, for most of the samples, good agreement between the calculated and observed values indicates that there was no appreciable formation of complexes other than those mentioned in Table 4. The extent of complexation, in those samples in which it was studied, was independent of the total cadmium concentration in the range stated. The extent of complexation increased with pH value and with the proportion of sewage effluent in the water (cf. especially Samples 9 and 10). Most of the complexation in sewage and sewage effluent (Samples 2--4) resulted from the presence of humic material, with the carbonate complex usually as the next most prominenL In the unpolluted water (Sample 1) approximately 90 per cent of the cadmium was present as the free ion. CONCLUSIONS
The work described is concerned only with the chemistry of cadmium in solution, where the dissolved
The chemistry of cadmium in natural water--l
29
Table 4. Extent of complexation of cadmium in borehole water, settled sewage, sewage effluents and river water samples Calculated proportion (~o)as Sample No. 1
2 3 4 5 6 7 8 9 10 11
CdOH +
CdCO3
CdCI ÷
CdSO4
Cd humic complex
1.4 l'8 3.2 3.6 2.8 6-5 4.9 5-7 4.8 3.6 2.6
3-9 9.0 t5 12 6.1 21 16 18 12 9.7 3.9
1.8 5"3 5"2 6.2 4.6 2.6 6"0 3"8 10 9-2 3"5
0-6 3.1 2"5 3.0 7.2 2.6 4-3 4-1 5-1 7.7 7.2
0 39 38 37 24 9-3 24 16 12 20 24
(mV)
Observed proportion as Cd-' ~ (O.g)
1.5 12.3 14-4 15.9 6.0 7.8 13.8 8.6 5-5 8.1 " 7.9
89 38 32 29 63 54 35 51 65 53 54
Eo-E Cd-' + 92 41 35 38 55 59 44 52 56 51 58
Total added cadmium concentration was 1'0 mg 1-t except for Samples 2-4 (0.3 mg l -t) and Sample 5 (2 mg I -t)
state is defined by the ability to pass through a 0.45/~m filter. The question of precipitation and adsorption will be dealt with in Part II of this Study. The stability constants measured compare well with published values obtained under slightly different conditions (Sill6n and Marteil, 1964), in those cases in which comparisons are possible. The results for complexation by humic substances are in good agreement with those found by Schnitzer (1971) for zinc and manganese. Equal concentrations of the extracted humic acid and natural humic material in sewage effluent seemed to complex the metal to a similar extent. The humic complexes were labile in the sense that they were substantially formed in l or 2 min, and probably less, but incorporation into a polymerized complex to give non-labile complexes of the type studied by Matson (I 968) may occur over a longer period. Whether the interaction is regarded as complex formation or adsorption will depend on the size of the humic molecule. Of the more common potential ligands not considered, phosphate and cyanide are likely to give complexation of cadmium only if their concentrations are exceptionally high (Sill6n and Martell, 1964) e.g. 0-3 mg 1- ~ of free cyanide at pH 8.0. Carboxylic and hydroxycarboxylic acids would also need to be at abnormally hi'gh concentrations, e.g. citrate at about 2 mg 1- t. Amino acids could give slight complexation, but for most of them, giving complexes of stability similar to those of glycine, the concentration would have to be higher than that commonly encountered in sewage-effluent. In the sewage-effluent samples of Table 4. little or no contribution to the proportion complexed was detectable from this source. No evidence has so far been produced for naturally occurring alkyl complexes
of cadmium like those of mercury, and in any case alkylating organisms would not have passed the filter in these experiments. It is necessary to consider whether, at cadmium concentrations in the/~g 1-~ region, the extent of complexation would be similar to that found at the mg 1- t concentrations in this work. Contrary to a widespread misconception, the ratio of complexed to uncomplexed cadmium depends only on the stability constant and the concentration of ligand, not on the total cadmium concentration, provided the ligand remains in excess. Consequently the extent of complexation by such ligands as carbonate and chloride will be the same at the lower cadmium concentration. In unpolluted water with a very low concentration of humic material (as in Sample 1, Tables 3 and 4), the ratio of free to complexed cadmium is certain to remain high at total cadmium concentrations of about I pg 1-~. In polluted water and otherwise unpolluted water containing a high concentration of humic material, this may not be the case. As humic acid is a mixture, and different fractions may have different complexing properties, it is possible that very low concentrations of metal ion may be more highly complexed at a given humic concentration than at higher metal concentration when less active sites may be employed. Highly complexing ligands at concentrations in the #g 1-t region may also become important. Their effect would obviously be negligible if, as in the above experiments, the metal was in large excess. For a ligand to be effective at a concentration of, say, 10-8 M however, the stability constant of its complex with cadmium would have to be about 108 M -~ or above. This excludes all naturally occurring ligands except any free cysteine which may be present and perhaps certain highly active humic acid fractions.
30
J. GAROISeR
Synthetic chelating agents like NTA and EDTA. the second of which is at present found in sewage effluent. could be important in this context; but other factors arise, such as competition between metal ions for the available iigand, which it is hoped will be dealt with in a later paper. Only at E D T A concentrations of 20/~g 1or above is complexation likely to be of importance. (The concentration in sewage effluent is expected to be up to 100/zg 1- ~.) It can consequently be predicted that although the proportion of free cadmium ion may be reduced at the /~g 1- t level it is likely to form a substantial fraction of the total cadmium concentration in polluted fresh water, and a major fraction in unpolluted fresh water with little humic material present. The behaviour of cadmium in this respect is therefore very different from that of copper (Stiff, 1971b), which is invariably highly complexed in polluted and unpolluted fresh water. Because a substantial proportion of the total dissolved cadmium is expected to be present as the free ion, the time of survival of fish in most natural waters containing cadmium.can be assessed with reasonable accuracy from a knowledge of the total cadmium concentration, without any reference to the degree of complexation, unless the toxicity of cadmium complexes present is very much greater than the toxicity of the free ion. Most experimental work on the toxicity of cadmium to fish has been performed in relatively pure fresh ground water in which most of the cadmium must have been present as the free ion. This would help to explain why, unlike copper..the toxicity of cadmium is relatively independent of hardness (Brown and Shurben, 1973).
Acknowledgement--Crown copyright. Reproduced by permission of the Controller, H.M. Stationery Office.
R EFERENCES
American Public Health Association 11971) Standard Methods for the Examination of Water and Waste-water. 13th edn, pp. 334--335. Brown V. M. and Shurben D. G. (1973) Unpublished work at Water Pollution Research Laboratory. Garrels R. M. and Thompson M. E. (1962) A chemical model for sea water at 25"C and one atmosphere total pressure. Am. J. Sci. 260, 57-66. Hem J. D. (1972) Chemistry and occurrence of cadmium and zinc in surface water and ground-water. War. Resour. Res. 8, 661-679. Lake P. E. and Goodings J. M. (1958) The nature of the. cadmium ions in hydroxide arid carbonate solutions. Can. J. Chem. 36, 1089-1096. Mangel M. S. (1971) A treatment of complex ions in sea water. Mar. Geol. 11, M24-M26. Matson W. R. (1968)Trace metals, equilibrium and kinetics of trace metal complexes in natural media. Ph.D. Thesis, Massachusetts Institute of Technology. Matson W. R., Allen H. E. and Rekshan P. (1969) Trace metal complexes in the Great Lakes. Presented before the Division of Water, Air and Waste Chemistry. American Chemical Society, Minneapolis, Minnesota. Ministry of Technology (1968) Water Pollution Research 1967, p, 60. H.M. Stationery Office, London. Schnitzer M. (1971) Organic Compounds in Aquatic Environmerits (Edited by Faust S. D. and Hunter J. V.). Dekker, New York, 297-315. Sill6n L. G. and Martell A. E. (1964) Stability Constants of Metal-lon Complexes. Special publication No. 17. The Chemical Society, London. Sill~n L. G. and Martell A. E. (1971) Supplement No. 1 to Stability constants of metal-ion complexes. Special publication No. 25. The Chemical Society, London. Stiff M. J. (197 la) C0pper-bicarbonate equilibria in solutions of bicarbonate ion at concentrations similar to those found in natural water. Water Research 5, 171-1"/'6. Stiff M. J. (1971b) The chemical states of copper in polluted fresh water and a scheme of analysis to differentiate them, Water Research 5, 585-599. Stiff M. J. and Gardiner D. K. (1973) Unpublished work at Water Pollution Research Laboratory.
THE CHEMISTRY OF C A D M I U M IN NATURAL W A T E R - - I A STUDY OF C A D M I U M COMPLEX F O R M A T I O N USING THE C A D M I U M SPECIFIC-ION ELECTRODE J. GARDINER Water Pollution Research Laboratory, Department of the Environment, Stevenage, England (Received 16 June 1973)
Al~tract--The extent of formation of labile complexes of cadmium has been investigated in synthetic solutions and in real samples. A substantial proportion of the total cadmium in river and lake water will usually be present as the free cadmium ion, and this proportion will be larger the lower the pH value and the lower the proportion of sewage emuent present in the water. Humic substances usually account for most of the complexation, followed in importance by carbonate, the complex of which with cadmium has not been reported previously.
cadmium, at the /ag 1-' level, and Matson and coworkers have devised methods for measuring the extent of formation of labile and non-labile complexes (Matson, 1968; Matson, Allen and Rekshan, 1969). For the purposes of this work, however, formation of labile complexes was found not to give sufficiently large changes in signal. The cadmium specific-ion electrode is extremely useful in a study of complex formation since it responds only to free cadmium ion in solution and not to complexes. Consequently, if the total cadmium concentration is known, the proportion complexed can be calculated. Its major disadvantage is that the minimum concentration at which it can be used is about 0.1 mg 1-t, which is considerably higher than the concentrations usually encountered in natural waters, though extrapolation to lower concentrations could be made with reasonable certainty. Measurements were made a few minutes after the addition of cadmium to a particle-free, filtered solution at concentrations below those at which precipitation would occur. Problems arising from adsorption and from slow reactions which may have occurred in some natural media were, therefore, minimized. Because of the comparatively small number of colorimetric reagents for cadmium it was not pbssible to devise an analytical scheme like that for copper (Stiff, 1971b). A considerable amount of data on complexation of cadmium already exists (Sill~n and Marteil, 1964, 1971). These data have been used by Mangel (1971) to predict complexation in sea water and by Hem (1972), and it is clear that a wide range of inorganic and organic ligands are potentially important. Measurements were
INTRODUCTION
The work described is a continuation of research in this Laboratory on complex formation of metal ions in natural waters (Stiff, 1971b). Metal complex formation could have important consequences both with regard to toxicity to aquatic organisms, because the toxicity of complexes can be different from that of the free ion, and also in analysis. Complex formation also affects the rate and extent of the disappearance of the metal ion from solution by precipitation and adsorption. Knowledge of the total metal concentration alone is therefore not always sufficient to enable toxicity and other properties to be predicted. Cadmium was chosen as the next metal for investigation after copper because of the current interest in its distribution throughout the environment and because it is extremely toxic to fish, effects on growth rate being observed even for concentrations between 5 and 10 #g 1- ' (Brown and Shurben, 1973). The toxicity curve for cadmium to rainbow trout, i.e. the curve of median survival time against concentration, shows a plateau region from 20 #g 1- t to 1"0 mg 1- t of cadmium over which the median lethal response time is relatively independent of concentration (Ministry of Technology, 1968). There are several methods available for measuring total cadmium concentration (e.g. atomic absorption spectroscopy and spark source mass spectrometry) but none of these would be suitable for measuring complex formation. Anodic stripping voltammetry is capable of measuring concentrations of some metal ions, including 23
24
J. GARDINER
made only on those systems for which there was no information or for those which, although existing data indicated the possibility of complexation, there was inadequate information at the ionic strength of interest. EXPERIMENTAL Materials I. Cadmi,~m solutions. Dilute stock solutions were freshly made up each day from a more concentrated solution of BDH reagent grade cadmium nitrate. 2. Humic acid solutions. Commercial humic acid (Aldrich Chemical Co.) was partially dissolved in sodium bicarbonate solution, filtered, and extracted by nhexanol. The organic extract was then evaporated down. This procedure was necessary to reduce contamination by metal impurities, notably zinc. 3. Other materials. All other substances were BDH AR grade unless otherwise stated. Doubly distilled water was used. 4. Ri~'er water and sewage samples. Samples of crude sewage, sewage effluent and river water were first allowed to settleifnecessary, then filtered under pressure through GF/A, GF/C, and 0.45/Jm membrane filters successively. One river-water sample contained sufficient cupric ion to contaminate the cadmium specific-ion electrode. A short column of Chelex I00 resin (Bio Rad Laboratories) was prepared by passing down it a concentrated phosphate buffer solution at pH 7, followed by thorough washing with distilled water until the washings failed to give a precipitate with silver nitrate solution. Passage of the filtered river-water sample down this column removed the source of contamination of the specific-ion electrode and left the pH value of the sample practically unaltered. Equipment Measurements were made in a cell containing approximately 100 ml of sample. The cell contained an Orion 94/48A cadmium specific-ion electrode, a silversilver chloride reference electrode, and a glass electrode with internal reference connected to a Radiometer No. 26 pH meter. The potential between the specific-ion and Ag/AgC1 electrodes was measured, using a similar pH meter, to within +0-3 inV. A mixture of nitrogen and carbon dioxide could be bubbled through or blown over the surface of the solution, and there was magnetic stirring in addition. Procedure A known quantity of cadmium nitrate (usually between 0"l and l'0 mg Cd l - I ) was added to the sample. The amount of cadmium already in the sample
was insignificant by comparison with the amount added. When required, potassium nitrate was added to adjust the ionic strength approximately to that of a hard natural water. A period of 15 min was allowed for the electrode system to reach equilibrium, which was probably longer than necessary. Experiments were carried out at room temperature (20-0 + 1.0~C). In most cases the pH value was controlled by varying the composition and flow rate of the mixture of nitrogen and carbon dioxide admitted to the cell. In the case of the hydroxy complex, however, pH values greater than 7 were achieved by slow addition of 0-04 M NaOH from an Agla micrometer syringe. CO_, being excluded from the cell. A reverse titration was then carried out with 0-04 M HCI. In order to determine the extent of complexation it was necessary to know the potential corresponding to no complexation. This could be determined in some cases by lowering the pH value until the potential was independent of pH value, or by measuring, before and after measurements on the sample, the potential developed in a non-complexing solution with the same total cadmium concentration and the same ionic strength. Determination of bicarbonate, chloride and sulphate concentrations Bicarbonate and chloride concentrations in unknowns were measured by the conventional methods, i.e. for bicarbonate, titration with N/50 hydrochloric acid using bromo-phenol blue as indicator, and for chloride, titration with silver nitrate solution, using potassium chromate as indicator. Sulphate was determined turbidimetrically by reaction with barium chloride (American Public Health Association, 1971). Determination of humic concentration An approximate measurement of the concentration of humic material in a filtered river-water or sewageeffluent sample was obtained from the absorbance at 400 nm using an SP.800 spectrophotometer, by comparison with solutions of known composition.
RESULTS AND DISCUSSION Behaviour of the specific-ion electrode Below pH values at which formation of the hydroxy complex was significant (see below) down to pH 5 and probably below, the potential was found to be independent of the pH value. Readings were very reproducible provided that the system was allowed to reach equilibrium, but rapid changes of pH value did cause a temporary departure from the equilibrium value.
The chemistry of cadmium in natural water--I
25
Table 1. Potential between the cadmium specific-ion electrode and Ag/AgCI reference electrode in a 1.0 mg 1-1 cadmium solution (8.9 x I0 -~ M) in 4 x I0-* M NaHCO 3 at pH 7.5 at various ionic strengths Ionic strength (KN03) (raM) Negative potential (mV)
2'0 178.1
4-0 6.0 8.0 10'0 12-0 14-0 18"0 20"0 180-0 181.4 182.5 183.4 184.1 185.0 186.0 186-8
Theoretically the potential, E, follows the relation E = A + Blogacd_,+ =
A + B Iogfcd.,- + BlogECdZ÷],
(1)
where A and B are constants and acd:" and fed:- are the activity and activity coefficient of the cadmium ion respectively. Plots o r e against log I-Cd-"+], with [Cd z +] from 1.0 x 10 -4 to 2.0 x 10 -6 M in 2 x 10 -3 M KNO3 were straight lines. The slope was 30 (-I- 1) mV for each 10-fold increase in concentration, which is not inconsistent with the theoretical value of 29-6 mV for B in equation (1). The form of equation adopted, therefore,
0"5 mg 1- ~, with the total cadmium concentration less than 1-0 mg 1- t the value orB decreased below 29.6 mV. Consequently, whenever possible, [Cd]r was kept at 1.0 mg 1-t or above. Contaminationof the electrode by cupric ion caused a change in the characteristics of the electrode, even after the source of the contamination had been removed. It was then necessary to lightly clean the crystal with fine emery cloth. Let [Cd]r be the total cadmium concentration and E0 the potential when there is no complexation. From equation (2) the potential change on complexation, Eo - E, is given by
was
E = A ' + BlogECd2+],
where B = 29'6 mV and the value of A' varies with the ionic strength of the solution (see Table 1), and also from day to day at constant ionic strength over a range of about 8 mV. The response of the electrode to log [Cd 2+] was linear down to 0-5 mg 1- ~ but the value of B in equation (2) then began to decrease (see Fig. 1) until below 1.0 #g I- t the signal was independent of cadmium concentration. The specifications of the electrode state that it is sensitive to [Cd z+] down to 10 ngl--t provided that the total cadmium concentration, [Cd]r, is greater than 10 ,ug I-t. It was found, however, that when complex formation occurred so as to lower [Cd-" +] below about
19(
Eo - E
(2) but
:
B
log
[Cd]~
[Cd:']'
[Cd]r = [Cd 2 +] + l e d complex],
(4)
where [Cd complex] is the concentration of cadmium complexed, assuming no losses by precipitation or adsorption. Using equations (3) and (4) the ratio [Cd complex]/[Cd-" +] can be calculated from E o - E.
Complexation by individual ligands Hydroxide. A comparatively low cadmium concentration of0"l mg 1-t was used so that a pH value of 10 could be reached without precipitation of cadmium hydroxide. All potential ligands except hydroxide were absent. The concentration of nitr.ate (2 x 10 -3 M) was
+ txpmu,4cm'~ ~04~Ts
Z
m 21(
~
Z~
+ .........
-4
), I"
24C
2SC
(3)
l t % ~ O ~ N T ~ T , O N . (,~n) ~0'
0~
Fig. 1. Variation in potential with concentration of the cadmium specific-ion electrode.
:6
J. GARDINER
too small to give an)complexation (Sill6n and Martell, 1964). On addition of sodium hydroxide the potential remained constant up to pH 7.4. but above this value it began to d:crease.
known to occur at high carbonate concentrations (f13 = 1.7 × 10 6 M -3, Lake and Goodings, 1958). It is known that soluble complexes of C a : " and Cu-' " exist (Garrels and Thomson, 1962; Stiff. 1971a) and their £d(H:O)~" ~ CdOHtH.O)~ + H*. (5) stability constants are similar to those of oxalate complexes with the same metal ions (Sill~n and Martell, Let 19641. The stability constant of the cadmium oxalate complex is given variously as 10s5. 103.9 and 10"t° M - t FCdOH(HzO);] ~tH "',- [ C d O H ' ] {H~} by Sill~n and Martell (1964), and the stability constant K°H = [Cd(H:O)~*] = [Cd z +] ' of a carbonate complex should be in this region. (6) Experiments were carried out by measuring the variawhere Cd(H.,O)~ ÷ and CdOH(H,O)~ are abbreviated tion in the extent of complex formation with pH value to Cd" ~ and CdOH + respectively. in a pH range in which the extent of formation of CdOH + was low. in the presence of various concentrations of sodium bicarbonate. The results, shown in Log [CdZ+] = l o g K o . + pH, (7) Fig. 3, indicate 1:1 complex formation involving one proton per molecule and are consistent with the where pH = - l o g [ H + ~,. The readings were stable, and on adding acid the following reaction: potential at a given pH value was shown to be indepenCd '~ + HCO~ ~ C d C O 3 + H + (8) dent of whether the pH was rising or falling, i.e. the change was completely reversible. A plot of log [CdOH +]/[Cd: +] against pH value was a straight line where CdCO~ is a soluble complex. From a knowledge (Fig. 2), as predicted by equation (7), but the slope was of the acid association constant of the carbonate ion 0-84 instead of 1.0 as expected. This was probably caused [Hco{] by the approach to the detection limit of the electrode. K . = [ C O ~ - ] (H +} (9) From Fig. 2, log K o , was calculated to be -9.06 (_+0.08) at 20°C, ~ = 2 x l0 -3 M (KNO3). Carbonate. A l : I carbonate complex of cadmium has ( K , --- 2-0 × 10 '° M - t , Sill~n and Martell. 1964)it is not been reported previously, although Cd(CO3) ~- is possible to calculate values for Kco~ the stability constant of the carbonate complex,
[CdOH÷2
0
//
[CdCOa] Kco,=(Cd,.] [COF]'
(10)
0.4
0.2
a
9
0
I
tJ.
-o.2
L
pH vAt.U(
-0"4 U
S-0., / ) ) -I'0
- 1-2
- 1.4
-I-6
Fig. 2. Graph of log[CdOH*]/[Cd-'*] against pH value. Total cadmium concentration. 0"10 mg 1- ~.
as respectively 10't°3, 10'~'14 and 104.0-' M - t at 2"5, 5"0 and 10.0 mM H C O j . Of these values the last is the most accurate. The solubility product of CdCO 3 is well documented. Previous work in this Laboratory (Stiff and Gardiner, 1973) has shown that in a clear solution the system is slow to reach equilibrium. In the experiments described here, with [Cd]r at 0-9 mg 1-*, and at the higher bicarbonate concentrations, the solubility product was frequently exceeded. Precipitation occurred only in some of the preliminary experiments and reversible changes then followed, cloudiness of the solution eventually becoming evident. The results of Fig. 3 could not have been due to partial precipitation because changes in potential with pH value were reversible and reproducible, unlike the effects observed with precipitation. Furthermore, the observed extent of complexation was, within experimental error, independent of [Cd]r, whereas the extent of precipitation would be dependent on [Cd]r.
27
The chemistry of cadmium in natural water--!
+÷
= 0.Zl-
,/
0
-0-~
?-
.
'
'
~"
//s//s /
, •
/
-0"4
INtos"],,s,,u
/
" [ H C O ! " ] I tO a M
-0-6
s
/s
//
7, II
/.
+
[.cos-l.=.s../
i/I
.,'
,
//~'HEOm[TICAL
/
t.co:l-o
Fig. 3. Effect of pH value on log(led complex]/[CdZ+]) in bicarbonate media. Total cadmium concentration, 1"0 x 10-s M (0.9 mg I-t). There was therefore strong evidence for a soluble carbonate complex with log Kco~ = 4.02 (+0.04) at 20'C, # = 10 x 10-3 M, in good agreement with the prediction based on comparison with the oxalate complex. Chloride and sulphate. Values given by Sill~n and Martell (1964) for the stability constants of the 1:1 complexes of chloride and sulphate with cadmium indicate that these complexes can be expected in natural waters. The extent of complexation was consequently measured in solutions containing various concentrations of sodium chloride and sodium sulphate respectivelyat an ionic strength typical of a hard natural water, and with a total cadmium concentration of i~) nag 1- t.
[CdCV]
slopes of which gave, respectively, Ko = 48 (+8) M - t and Kso, = 220 (+ 10)M- t both at # = 10 x 10 -3 M, and 20°C. The extent of complex formation was, of course, independent of pH value in both cases. Humic material. Because "hurnic acid" is a mixture and its mol. wt and precise acid-base properties are unknown, it is not possible to establish stability constants for complexation of metals in the usual way, nor is it certain that a mixture of humic material obtained from one source will behave in the same way as that from another. Solutions were made up containing 86, 20 and 6-6 mg 1-1 of humic acid purified as described, representative of a high, moderate and low level of coloration respectively. Each solution contained 3 x 10 -4 M bicarbonate, which was sufficient to allow the pH value to be varied over the range of interest by varying the flow of the N,/CO., mixture to the cell, but not sufficient to give appreciable carbonate complex formation. The fraction complexed was found to be slightly dependent on pH value at all three concentrations of humic material, and the ratio of complexed to uncomplexed cadmium ion did not increase linearly with humic acid concentration (see Table 2).
(II)
Ko = [Cd2 +] [ C l - ] ' [CdS04]
(12)
tqo, = [Cd2+] [SO=_]. In both cases plots of [Cd complex]/[Cd' +] against the concentration of ligand were straight lines, the:
[Cd complex]
['Cdz+]
oc
[ligand] °6'~
{H+} °'2° •
The low p H dependence could indicatethat the ligand is largely unprotonated in the p H range studied or, more probably, that the fraction of humic material responsible for most of the complexation is a strong chelating ligand relativelyunaffected by pH changes. The ligand dependence shows that some of the ligand molecules complex more than one cadmium ion, which is consistent with a polymeric ligand structure and perhaps with polymerization of the complex. Results were independent of cadmium concentration over the range [Cd]r 0-I-I'0 mg I-', afterallowing for changes caused by the approach to the detection limit, and assuming the average tool. wt of the humic material to be about 700 (Schnitzer, 1971), the ligand was in excess in all cases. A convenient formula based on equation (13) for calculating the degree of complexation by humic
Table 2. Dependence of the degree of complexation of cadmium by humic substances on pH value pH value
6.0
6"5
7.0
7.5
8-0
8"5
2.50 1.00 0.50
3-20 • 1.26 0-62
[Cd complex]/[C-d 2+] Humic concentrations (rag I-t)
86 20 6"6
1"00 0.40 0-20
(13)
1-26 0-50 0-25
1-59 0.62 0.32
2-00 0.77 0"40
28
J. GARDINER
Table 3. Composition of samples used in measuring extent of complexation of cadmium
Sample •
No.
Type ofsample Fresh ground-water* Filtered settled sewage'l" Filtered poor quality percolating filter efltuent Filtered good quality percolating filter ettluent Filtered river water Filtered river water
7
Filtered river water
8
Filtered river water
9
Filtered river water
10
Filtered river water
11
Filtered river water
Point of collection
Absorbance Approx in40-mm ionic cell at strength 400nm (mM)
pH value
[HCO~] (mM)
[CI-] (mM)
[SO~-] (mM)
WPRL WPRL
7.2 7-7
5" 1
0.39 2'8
0'030 0-34
0
8.4
0-22
17
WPRL Filter 30
8-0
8'2
3-0
0"31
0.22
!7
WPRL Filter 1
8'0
5-8
3'1
0"33
0.16
14
Stevenage Brook WPRL R. Ivel, Stotfold, Herts Pix Brook, Arlesey, Herts R. Hiz, Arlesey, Herts R. Lee, Luton++ R. Lee, Luton§ R. Trent, Swarkeston¶
7-7
4-4
1.8
0"59
0-07
l1
8"1
5.2
0-94
0.20
0.01
9
8"1
5"5
2-8
0-44
0.07
12
8.1
5.2
1'5
0.36
0.03
Il
8"0
4. I
3.7
0-40
0"02
12
7-9
4.8
3-6
0.68
0:05
14
7.7
2.6
1'3
0-35
0"06
8
8
* Used in fish toxicity tests. ? +2 drops I00 vol H20 2. :~Upstream of Luton sewage works outflow. § Downstream of Luton sewage works outflow. ¶ Four miles south of Derby. material from the absorbance of the solution is log
[Cd complex] [Cd 2+] 0"64 log A4oo + 0.20 pH - 1.14, 04)
where A4oo is the absorbance of the filtered solution at 400 nm in a 40-ram cell, assuming ferric ion and other interfering substances are absent or negligible.
Complexation in samples of river water and sewage effluent. The source, type and composition of the sewage and river water samples used are given in Table 3. From the concentrations of the ligands present, using the measured stability constants and equation (14), it was possible to calculate the proportion of cadmium present in complex forms and as the free ion (see Table 4). The results of a direct determination of the extent of complexafion, obtained by adding 0.1-10mgl - t of cadmium to the sample and measuring the proportion remaining as free ion, are also given in Table 4 for
comparison. The fact that there was, for most of the samples, good agreement between the calculated and observed values indicates that there was no appreciable formation of complexes other than those mentioned in Table 4. The extent of complexation, in those samples in which it was studied, was independent of the total cadmium concentration in the range stated. The extent of complexation increased with pH value and with the proportion of sewage effluent in the water (cf. especially Samples 9 and 10). Most of the complexation in sewage and sewage effluent (Samples 2--4) resulted from the presence of humic material, with the carbonate complex usually as the next most prominenL In the unpolluted water (Sample 1) approximately 90 per cent of the cadmium was present as the free ion. CONCLUSIONS
The work described is concerned only with the chemistry of cadmium in solution, where the dissolved
The chemistry of cadmium in natural water--l
29
Table 4. Extent of complexation of cadmium in borehole water, settled sewage, sewage effluents and river water samples Calculated proportion (~o)as Sample No. 1
2 3 4 5 6 7 8 9 10 11
CdOH +
CdCO3
CdCI ÷
CdSO4
Cd humic complex
1.4 l'8 3.2 3.6 2.8 6-5 4.9 5-7 4.8 3.6 2.6
3-9 9.0 t5 12 6.1 21 16 18 12 9.7 3.9
1.8 5"3 5"2 6.2 4.6 2.6 6"0 3"8 10 9-2 3"5
0-6 3.1 2"5 3.0 7.2 2.6 4-3 4-1 5-1 7.7 7.2
0 39 38 37 24 9-3 24 16 12 20 24
(mV)
Observed proportion as Cd-' ~ (O.g)
1.5 12.3 14-4 15.9 6.0 7.8 13.8 8.6 5-5 8.1 " 7.9
89 38 32 29 63 54 35 51 65 53 54
Eo-E Cd-' + 92 41 35 38 55 59 44 52 56 51 58
Total added cadmium concentration was 1'0 mg 1-t except for Samples 2-4 (0.3 mg l -t) and Sample 5 (2 mg I -t)
state is defined by the ability to pass through a 0.45/~m filter. The question of precipitation and adsorption will be dealt with in Part II of this Study. The stability constants measured compare well with published values obtained under slightly different conditions (Sill6n and Marteil, 1964), in those cases in which comparisons are possible. The results for complexation by humic substances are in good agreement with those found by Schnitzer (1971) for zinc and manganese. Equal concentrations of the extracted humic acid and natural humic material in sewage effluent seemed to complex the metal to a similar extent. The humic complexes were labile in the sense that they were substantially formed in l or 2 min, and probably less, but incorporation into a polymerized complex to give non-labile complexes of the type studied by Matson (I 968) may occur over a longer period. Whether the interaction is regarded as complex formation or adsorption will depend on the size of the humic molecule. Of the more common potential ligands not considered, phosphate and cyanide are likely to give complexation of cadmium only if their concentrations are exceptionally high (Sill6n and Martell, 1964) e.g. 0-3 mg 1- ~ of free cyanide at pH 8.0. Carboxylic and hydroxycarboxylic acids would also need to be at abnormally hi'gh concentrations, e.g. citrate at about 2 mg 1- t. Amino acids could give slight complexation, but for most of them, giving complexes of stability similar to those of glycine, the concentration would have to be higher than that commonly encountered in sewage-effluent. In the sewage-effluent samples of Table 4. little or no contribution to the proportion complexed was detectable from this source. No evidence has so far been produced for naturally occurring alkyl complexes
of cadmium like those of mercury, and in any case alkylating organisms would not have passed the filter in these experiments. It is necessary to consider whether, at cadmium concentrations in the/~g 1-~ region, the extent of complexation would be similar to that found at the mg 1- t concentrations in this work. Contrary to a widespread misconception, the ratio of complexed to uncomplexed cadmium depends only on the stability constant and the concentration of ligand, not on the total cadmium concentration, provided the ligand remains in excess. Consequently the extent of complexation by such ligands as carbonate and chloride will be the same at the lower cadmium concentration. In unpolluted water with a very low concentration of humic material (as in Sample 1, Tables 3 and 4), the ratio of free to complexed cadmium is certain to remain high at total cadmium concentrations of about I pg 1-~. In polluted water and otherwise unpolluted water containing a high concentration of humic material, this may not be the case. As humic acid is a mixture, and different fractions may have different complexing properties, it is possible that very low concentrations of metal ion may be more highly complexed at a given humic concentration than at higher metal concentration when less active sites may be employed. Highly complexing ligands at concentrations in the #g 1-t region may also become important. Their effect would obviously be negligible if, as in the above experiments, the metal was in large excess. For a ligand to be effective at a concentration of, say, 10-8 M however, the stability constant of its complex with cadmium would have to be about 108 M -~ or above. This excludes all naturally occurring ligands except any free cysteine which may be present and perhaps certain highly active humic acid fractions.
30
J. GAROISeR
Synthetic chelating agents like NTA and EDTA. the second of which is at present found in sewage effluent. could be important in this context; but other factors arise, such as competition between metal ions for the available iigand, which it is hoped will be dealt with in a later paper. Only at E D T A concentrations of 20/~g 1or above is complexation likely to be of importance. (The concentration in sewage effluent is expected to be up to 100/zg 1- ~.) It can consequently be predicted that although the proportion of free cadmium ion may be reduced at the /~g 1- t level it is likely to form a substantial fraction of the total cadmium concentration in polluted fresh water, and a major fraction in unpolluted fresh water with little humic material present. The behaviour of cadmium in this respect is therefore very different from that of copper (Stiff, 1971b), which is invariably highly complexed in polluted and unpolluted fresh water. Because a substantial proportion of the total dissolved cadmium is expected to be present as the free ion, the time of survival of fish in most natural waters containing cadmium.can be assessed with reasonable accuracy from a knowledge of the total cadmium concentration, without any reference to the degree of complexation, unless the toxicity of cadmium complexes present is very much greater than the toxicity of the free ion. Most experimental work on the toxicity of cadmium to fish has been performed in relatively pure fresh ground water in which most of the cadmium must have been present as the free ion. This would help to explain why, unlike copper..the toxicity of cadmium is relatively independent of hardness (Brown and Shurben, 1973).
Acknowledgement--Crown copyright. Reproduced by permission of the Controller, H.M. Stationery Office.
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