INTERACTION OF AMINO ACIDS AND SALTS III. THE DETERMINATION OF THE ACTIVITIES OF CALCIUM, BARIUM, AND STRONTIUM CHLORIDE IN AMINO ACID SOLUTIONS BY MEANS OF ELECTRODES OF THE THIRD KIND BY NORMAN R. JOSEPH (From the Children’s Hospital Research Foundation, Department of Pediatrics, and the Department of Internal Medicine, University of Cincinnati, Cincinnati) (Received for publication, June 29, 1939) An electrometric method for the determination of calcium ions in protein solutions has recently been described (6). Calcium amalgam, which is unstable in the presence of protein, is protected from protein by a cellophane membrane, and in this condition is employed as one of the electrodes in a cell without liquid junction. The cellophane membrane protects the amalgam from proteins, but does not protect it from amino acids and diffusible cations other than calcium. The amalgam electrode is not sufficiently stable in the .presence of these substances to yield reversible electromotive forces. This fact limits the use of this type of electrode to mixtures of proteins and calcium salts. For these reasons an alternative method has been sought, which will permit the electrometric determination of calcium ions under conditions which produce irreversible potentials in the calcium amalgam electrode. To be useful in biological fluids a calcium ion electrode must yield reversible potentials in the presence not only of proteins but of a variety of diffusible substances, many of which decompose calcium amalgam. In the following pages an electrode will be described which can be employed for the determination of ionic calcium at very low concentrations in the presenceof a variety of diffusible substances. In this method the electrode is lead amalgam in contact with lead oxalate and calcium oxalate. Electrodes of this type are called electrodes of the third kind. They have been studied by 203 by guest on June 19, 2018 http://www.jbc.org/ Downloaded from
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INTERACTION OF AMINO ACIDS AND SALTS
III. THE DETERMINATION OF THE ACTIVITIES OF CALCIUM,
BARIUM, AND STRONTIUM CHLORIDE IN AMINO ACID SOLUTIONS BY MEANS OF ELECTRODES OF
THE THIRD KIND
BY NORMAN R. JOSEPH
(From the Children’s Hospital Research Foundation, Department of Pediatrics, and the Department of Internal Medicine, University
of Cincinnati, Cincinnati)
(Received for publication, June 29, 1939)
An electrometric method for the determination of calcium ions in protein solutions has recently been described (6). Calcium amalgam, which is unstable in the presence of protein, is protected from protein by a cellophane membrane, and in this condition is employed as one of the electrodes in a cell without liquid junction. The cellophane membrane protects the amalgam from proteins, but does not protect it from amino acids and diffusible cations other than calcium. The amalgam electrode is not sufficiently stable in the .presence of these substances to yield reversible electromotive forces. This fact limits the use of this type of electrode to mixtures of proteins and calcium salts.
For these reasons an alternative method has been sought, which will permit the electrometric determination of calcium ions under conditions which produce irreversible potentials in the calcium amalgam electrode. To be useful in biological fluids a calcium ion electrode must yield reversible potentials in the presence not only of proteins but of a variety of diffusible substances, many of which decompose calcium amalgam.
In the following pages an electrode will be described which can be employed for the determination of ionic calcium at very low concentrations in the presence of a variety of diffusible substances. In this method the electrode is lead amalgam in contact with lead oxalate and calcium oxalate. Electrodes of this type are called electrodes of the third kind. They have been studied by
Luther (ll), by Corten and Estermann (3), and by LeBlanc and Harnapp (9). Corten (2) has applied this type of ‘electrode to the st.udy of biological fluids. The theory of the electrode has been thoroughly treated by LeBlanc and Harnapp.
Theory
In a saturated solution of two slightly soluble salts with a com- mon anion, such as lead oxalate and calcium oxalate, the concen- tration of one of the cations, lead for example, will depend on that of the other. Thus
and (Pb++)(CzO4-) = KI (1, 4
(Ca++)(C~O~-) = KS (1, b)
where Kl and KS denote the solubility products of the two salts and the symbols in parentheses denote ionic activities.
In such a mixture
(Pb++)/(Ca++) = KI/K~ = Kc, (2)
where K. is determined by the two solubility products. Thus over the entire range of calcium concentration above the limit of cal- cium oxalate solubility (Pb*) is proportional to (Ca++). This is true under two conditions: first that there be no other anions present in sufficient concentration to form other insoluble lead salts, and second that there be no other cations present in sufficient concentration to form insoluble oxalates. Thus we should not expect the electrode to function in solutions containing phosphate or zinc ions, for example.
Under controlled conditions, however, lead amalgam in contact wit.h saturated lead oxalate and a solution containing calcium chloride should act according to Equation 2 as a calcium ion electrode. The electrode reaction can be represented by the equation
Pb + CaCeOl = PbGOa + Ca++ -t 2e (3, a)
the symbol e denoting an electrochemical equivalent of negative electricity.
If a silver-silver chloride electrode be in equilibrium with the same calcium chloride solution, we have
At equilibrium all the substances in Equation 4 with the exception of calcium chloride are in the solid state or in equilibrium with it, so’that the reversible E.M.F. of the cell measures the mean ionic activity of this salt referred to its standard state.
The cell reaction represented by Equation 4 can be compared with that occurring with a calcium amalgam electrode in combina- tion with a silver-silver chloride electrode.
Ca + 2AgCl = CaCla + 2Ag (5)
In both cases the free energy change involved in the passage of 2 electrochemical equivalents of electricity through the cell is the formation of 1 mole of calcium chloride, from lead, calcium oxalate, and silver chloride in one case, and from calcium and silver chloride in the other. In both cases the E.M.F. of the cell is given by the relation
E - EO = - 3RT/2F In acacll (6)
where E”, the standard potential of the cell, is determined by the temperature, the concentration of the lead amalgam, and the standard potential of the silver-silver chloride electrode. R, T, and F denote respectively the gas constant, the absolute tempera- ture, and the Faraday equivalent. The symbol aoaClz denotes the mean activity of calcium and chloride ions
acam = Q&i (7)
While the direct proportionality of (Pb++) to (Ca++) holds down to the limiting value of (Ca++) determined by Kz, (Ca++) at very low concentrations is determined not only by the composition of the fluid that is being studied but also by the dissociation of calcium oxalate and lead oxalate. Thus there is a lower limit beyond which the electrode does not determine the calcium ions of the cell electrolyte. This limit is determined by the values of Kl and Kz. LeBlanc and Harnapp (9) have developed equations giving the relation between the true calcium ion concentration of the unknown fluid and that resulting from the dissociation of the oxalates in contact with the unknown. The resulting equations
are quite complex, and are omitted here. In Fig. 1 we have plotted the relation between the lead ion activity and the activity of alkaline earth ions of the unknown solution, for the case of all three alkaline earths, calcium, barium, and strontium. In all three cases, it is evident that below a certain limit (Pb++) is no longer directly proportional to the activity of the ions of the un- known solution. This lower limit depends on the solubility prod-
-LOG
(Mt3++l
- L 0 G ( Pb”)
FIG. 1. The relation between the activity of lead ions and alkaline earth cations in mixtures of the oxalates. (Me+) denotes the activity of the alkaline earth cations.
ucts of the several oxalates. In the case of all three, however, this limit is of the order of 1 mM or less, so that the electrodes should theoretically yield valid results down to 1 mM.
EXPERIMENTAL
For the purpose of studying the behavior of electrodes of the third kind, determinations of the activity of calcium, barium, and strontium chlorides have been carried out in water and in amino
acid solutions, in which it is possible to compare the results with those of other methods. The influence of sodium, potassium, and magnesium ions has also been studied.
The cell employed is that described in an earlier paper (6). It is an H-shaped tube, in one arm of which the amalgam electrode is suspended, while the other arm contains a silver-silver chloride electrode.
Lead amalgam was prepared by the electrolysis of lead acetate against a mercury cathode. Approximately 100 gm. of amalgam are introduced into the amalgam container, the mouth of which is immediately sealed with a cellophane membrane. The amalgam container, a calcium chloride drying tube with a sealed-in platinum electrode, has been described elsewhere, as has the method of applying the membrane (6).
Freshly washed lead oxalate is applied to the outside of the membrane in the form of a paste. Lead oxalate is precipitated from a concentrated solution of lead acetate by the addition of an excess of oxalic acid. A second membrane is fixed over the mouth of the tube, so that a thin layer of the lead oxalate paste is held between two membranes. The electrode is supported in the cell by means of the bulb of the drying tube.
The electrode has been studied in aqueous solutions of calcium, barium, and strontium chlorides. Equilibrium is reached gen- erally within an hour or less, although occasionally much longer time is necessary. The system is considered to be at equilibrium when five or six readings taken over a period of at least 15 minutes agree to f0.2 millivolt. For example, in a solution of 0.0024 M
CaClz a freshly prepared electrode showed an initial E.M.F. of 0.5560 volt against a silver-silver chloride electrode. In 1 hour and 5 minutes this reached a final constant value of 0.5375 volt. At 10 minute intervals following the initial reading the E.M.F.'S
were 0.5533, 0.5487, 0.5464, 0.5440, 0.5397, 0.5393, 0.5375, 0.5375. This final E.M.F. remained constant within f0.2 millivolt for 1 hour. When immersed in a solution of 0.0036 M CaClz it reached a final constant value of 0.5235 volt in 35 minutes. At 10 minute intervals the readings were 0.5370, 0.5262, 0.5250, 0.5244, 0.5235, 0.5235. This potential remained constant for 24 hours. The difference between these two final E.M.F.'s, 14 millivolts, agrees with the theoretical difference estimated from Equation 6.
Schmidt and Greenberg (16) have reported difficulty in attaining equilibrium with the zinc oxalate electrode. Their experience corresponds to that of LeBlanc and Harnapp (9) with that elec- trode, but not with that of Corten and Estermann (3), who were able to obtain the theoretical E.M.F. within half an hour. The
TABLE I
Activity Coejicient of Calcium, Barium, and Strontium Chlorides in Water at 8P
-(E - E*) is the E.M.F. of the given solution subtracted from that of the most dilute solution of the series. y* is the activity coefficient of the most dilute solution obtained from the data of Scatchard and Tefft (15) for calcium and barium chloride and from the data of Lucasse (10) for strontium chloride.
(a), activity coefficient estimated from E.M.F.; (5) activity coefficient estimated from data of Scatchard and Tefft and from that of Lucasse.
-Log y/y* is calculated from the relation, -log y/y* = 11.27 (E - E*) + log m/m*, where m* is the molality of the most dilute solution of, the series.
method of preparation of the zinc oxalate has not been specified by any of these authors. In the present study freshly precipitated, finely divided lead oxalate is employed, and equilibrium is usually attained within an hour.
The results of the measurements on pure salt solutions are presented in Table I. Activity coefllcients are computed from
the data, and are compared with those of Scatchard and Tefft (15). The agreement is good down to very low concentrations of salt,
With a given lead amalgam, the standard potential, E”, of the cell should vary for the different alkaline earth chlorides with the values of the solubility products of the alkaline earth oxalates. The standard free energy change, AFO, of the reaction Equation 3, a is given by
= -RT In KJK2 03)
TABLE II
Estimation of Solubility Products of Oxalates from Electrode Potentials
salt (0.01 M) E.?d.F. 4 - Bca LogKn+lO Kz (from E.M.P.) KS (observed)
mu CaClz. 0.5451 0.0000 1.2504 1.78 X 1O-9 SrClz . . 0.5888 0.0437 2.7490 5.34 x 10-s 5.61 X 10-a BaClz.. 0.5971 0.0520 3.0414 1.02 x 10-1 1.10 x lo-’
Kz is estimated from the E.M.F. data by the use of the equation E - Eta = 0.02957 log K2/K2(ca). The value of K~cQ), the solubility product of calcium oxalate, is taken from Kohlrausch (8). The values of the sol- ubility products of the other oxalates are estimated from the E.M.F. data, Kohlrausch’s value for K2(ca) being employed.
Similar equations hold for the strontium and barium electrodes. Thus it is possible from E.M.F. measurements in solutions of the three alkaline earth chlorides to correlate the solubility products of the three oxalates. Table II contains the values of the solu- bility products of strontium and barium oxalates obtained from E.M.F. data on 0.01 M salt solutions, the value of Kz for calcium oxa- late being taken as 1.78 X 10mg at 18”, as determined by Kohlrausch (8). The results show that the standard potentials of the elec- trodes vary in the manner predicted by the theory, and that the values of the solubility products obtained in this way agree with those determined by Kohlrausch, who employed the conductivity method.
In order to test the applicability of the electrode to solutions containing salts other than calcium chloride, a series of determina-
tions has been carried out in solutions of calcium chloride con- taining respectively sodium chloride, potassium chloride, and magnesium chloride. The results are presented in Table III, and in Fig. 2 the mean ionic activity coefficient of calcium chloride in these mixtures is plotted against the ionic strength, together with the values in pure calcium chloride solution. It is evident from these results that the values of the activity coefficient of calcium chloride in these mixtures are very near the values in pure
TABLE III Effect of Various Electrolytes on Activity of Calcium Chloride
AE is the change of E.M.F. produced by the addition of the second salt to the calcium chloride solution.
-Log y/‘y* is calculated by means of the relation, -AE = 0.05915 log Cl/Cl* + 0.08872 log y/y*, where Cl* and y* are respectively the chlo- ride ion concentration and the mean ionic activity coefficient of calcium chloride in the pure calcium chloride solution. The latter values are taken from Table I.
calcium chloride solutions of the same ionic strength. The magnitude of the deviations is within the range that would be predicted by the Debye-Hiickel theory, if the different ionic radius terms for the different salt mixtures are taken intoaccount. It can be concluded that these added cations, all of which form soluble oxalates, do not interfere with this type of electrode, even when present in excess.
In Table IV are presented the results of determinations of the
effects of glycine and alanine on the activity coefficients of calcium, barium, and strontium chlorides. The results are calculated by the relation
- log ~a/?; = 11.27 (E - Eo) (9) where y3 and 7: are respectively the mean ionic activity coefficients of the salt in the amino acid solution and in an isomolal salt solu- tion. E is the E.M.F. observed with the amino acid-salt mixture,
I I I I 0 .05 JO .I5
IONIC STRENGTH
FIG. 2. The mean ionic activity coefficient of calcium chloride in water and in the presence of other salts. The various mixtures are denoted by the symbols: 0 CaCL in water; + 0.0012 M CrtCl~ + KCl; X 0.0012 M
CaCL + N&l; 0 0.012 M CaC12 + NaCl; 0 0.012 M CaClz + KCl; A 0.012 M CaC12 + MgC12.
and Eo is that of the pure salt solution. 11.27 is the reciprocal of the ratio (3 X 2.303 RT)/2F at 25”.
In addition to the values of -log ys/yi, Table IV contains values of -log y2/& where “/z and -$ are the activity coefficients of the amino acids in the salt solution and in water. These values are calculated as in earlier papers (4, 5) by application of the thermo- dynamic equation
can be taken as a constant, and -log rz/$ is found to be directly proportional to tiz. This result agrees with our earlier findings and with those of Cohn (l), who in ethanol-water mixtures finds -log yz to be proportional to the ionic strength of the added salt.
In Fig. 3 are presented the values of -log 72/y: for glycine in solutions of the alkaline earth chlorides, plotted against p, the ionic strength of the salt solution. The results are compared graphically with the solubility results of Pfeiffer and Wurgler (13). Our results indicate somewhat greater values for -(a log y2)/8c(
-LOG
72
I 234 6 6
“- I 2 I0NI.C STRENGTH
FIG. 3. The effect of various salts on the activity coefficient of glycine in water. Curve 1, N&l (from E.M.F. data at 1.4” (5)); Curves 2, 3, 4, B&L, CaC$, SrClz, respectively at 25” (E.M.F. data, Table IV); Curve 5, BaCL (from solubility data, approximately 3 M glycine (13)); Curve 6, CaCL, SrClz (from solubility data, approximately 3 M glycine (13)).
than we calculate from the solubility data, but they are clearly of the same order of magnitude, and vary with respect to the different salts in the same manner. The rather small differences may pos- sibly be accounted for by the different experimental conditions, the solubility results having been obtained with saturated glycine solutions (approximately 3 M), and the electrometric measure- ments having been carried out in 1 M and 0.5 M glycine solutions.
The p,otentiometric results can also be correlated with the freezing point results of Pfeiffer and Angern (12). As we have
Calculations based on Equation 11,~ and the following values of - (a logra)a mz, estimated from data of Table IV: BaClz 0.16, CaCl, 0.14, and SrClz 0.125.
where A23, A,, and A3 are respectively the freezing point depressions of the three-component system, the isomolal amino acid solution, and the isomolal salt solution.
In Table V are entered the values of the freezing point effects for the systems calcium chloride-glycine, barium chloride-glycine, and strontium chloride-glycine, as estimated from E.M.F. data by means of Equation 11, a. These values are compared with the freezing point data of Pfeiffer and Angern (12). Although the freezing point data in dilute solutions are hardly adequate for a thorough comparison, the results are in good agreement as far as they extend. The deviations in the more dilute strontium chloride solutions appear to be considerably greater than the error of either method, but they may possibly be accounted for by the difference of temperature in the two sets of obsetvations.
In addition to the measurements on amino acid-salt mixtures a series of determinations of calcium chloride activity in egg albumin solutions has been carried out. The protein employed in these determinations was an electrodialyzed preparation. The specific conductivity of a 5 per cent solution was 18 X 1O-6 reciprocal ohm. The results are presented in Table VI. They show that in the presence of egg albumin, the activity of calcium chloride is diminished to a small extent. The magnitude of the effect, which is quite small, is about the same as that previously observed with carboxyhemoglobin and with serum albumin, and considerably less than that obtained with pseudoglobulin and gelatin (6).
TABLE VI
Activity of Calcium Chloride in Presence of Egg Albumin
AE is the increase in the E.M.F. produced by the addition of protein to the salt solution.
DISCUSSION
It is evident from the foregoing results that electrodes of t,he third kind are applicable to the study of the activities of the al- kaline earth chlorides both in water and in the presence of amino acids or of other inorganic salts.
The values of the salting-in coefficient (a log +/a~) agree satisfactorily with those obtained by the freezing point and solu- bility methods. These values lie considerably below the theo- retical limiting slope for amino acids, estimated by Cohn (1) to be 0.32 from solubility data in ethanol-water mixtures. As we have pointed out elsewhere (5), deviations from this limiting slope are to be considered as a measure of the salting-out effect in aqueous solutions. Here because of the high dielectric constant
of the solvent, the Coulomb forces are of a smaller order of magni- tude than in ethanol-water mixtures, and the salting-out forces thereby become relatively more significant.
Kirkwood (7) has recently elucidated this effect from the theo- retical point of view. He gives the equation, -log yz = (B: - Bi)I’, where the coefficient BP represents the salting-in coefficient, Bj represents the salting-out coefficient, and I’ denotes the ionic strength. For sodium chloride and water he estimates Bi to be approximately 25 per cent of By. From the foregoing experimental dat,a on the alkaline earth chlorides we should estimate Bi to be of the order of 50 per cent of Bi. A salting- out effect of this magnitude is quite consistent with the results obtained by Randall and Failey (14) for the salting-out of acetic acid and its chloride derivatives.
SUMMARY
1. The electrode PbHg 1 PbGO+ CaC&OJ, Gaff has been studied in pure calcium chloride solutions and in the presence of sodium, potassium, and magnesium ions. It has been found to come to equilibrium in 1 hour or less, and to yield stable, reproducible potentials.
2. The electrode has been used to study the activities of calcium, barium, and strontium chlorides in the presence of glycine and alanine.
3. From the data the effect of these salts on the activity coef- ficients of the amino acids has been estimated and compared with the effects observed by the solubility and freezing point methods.
4. It is concluded that the electrode behaves as a reversible calcium ion electrode not only in the presence of pure calcium chloride but also in the presence of other neutral salts or of amino acids.
BLBLIOGRAPHY
1. Cohn, E. J., Naturwissenschajten, 20,663 (1932); in Luck, J. M., Annual reviews of biochemistry, Stanford University, 4, 93 (1935); Chem.
Rev., 19, 241 (1936). 2. Corten, M. H., Centr. a&. Path. u. path. Anat., 44, 144 (1928). 3. Corten, M. H., and Estermann, I., 2. physik. Chem., 136, 228 (1928). 4. Joseph, N. R., J. Biol. Chem., 111,479 (1935).
6. Joseph, N. R., J. Biol. Chem., 126, 389 (1938). 7. Kirkwood, J. G., Chem. Rev., 24, 233 (1939). 8. Kohlrausch, F., 2. physik. Chem., 64, 166 (1908).
9. LeBlanc, M., and Harnapp, O., Z. physik. Chem., Abt. A, 166, 321 (1933).
10. Lucasse, W. W., J. Am. Chem. Xoc., 47, 743 (1925).
11. Luther, R., Z. physik. Chem., 27, 364 (1899). 12. Pfeiffer, P., and Angern, O., 2. physiol. Chem., 136,16 (1924).
13. Pfeiffer, P., and Wurgler, J., Z. physiol. Chem., 97, 128 (1916). 14. Randall, M., and Failey, C. F., Chem. Rev., 4, 271, 285, 291 (1927). 15. Scatchard, G., andTefft, R. F., J. Am. Chem. Sot., 62,2265,2272 (1930). 16. Schmidt, C. L. A., and Greenberg, D. M., Physiol. Rev., 16, 297 (1935).
