THE INFLUENCE OF NEUTRAL SALTS ON THE pH OF PHOSPHATE

THE

INFLUENCE OF NEUTRAL SALTS ON THE PHOSPHATE BUFFER MIXTURES. BY

(From

the

Depadment (Received

HOWARD of

W.

ROBINSON.*

Biochemistry, Vanderbilt Medicine, Nashville.) for

publication,

pH OF

April

University

School

of

5, 1929.)

* Presented to the Graduate fulfilment of the requirements

School of Vanderbilt University in partial for the Degree of Doctor of Philosophy. 775

Downloaded from http://www.jbc.org/ by guest on March 17, 2015

The results embodied in this report were obtained as preliminary data for a systematic study of the magnitude of errors produced by the presence of various substances in the use of phenol red for calorimetric determinations of hydrogen ion concentration in biological fluids. Neutral salts, present in all body fluids, are known to have an influence on dye color. There are available many tables of salt errors of indicators. In most cases these are given for one buffer mixture and one indicator concentration. The question arises as to the constancy of the salt error with change in indicator concentration, with change in the buffer concentration, and with change in the pH. Before this can be definitely answered it is necessary to study the effect of neutral salts on the true hydrogen ion concentration (as measured by the hydrogen electrode) for the entire range of the buffer and at various dilutions. S#rensen and Palitzsch (1) drew attention to the fact that the addition of neutral salt to buffer solutions altered the electrometric pH and pointed out that the entire change as observed by a calorimetric reading could not be called a salt error of the indicator. It is a well recognized fact that the addition of small amounts of neutral salt to dilute buffer solutions increases the hydrogen ion concentration as determined by the potential of the hydrogen electrode. At the present time all other electrometric and also calorimetric, titrimetric, and gasometric methods for the determination of H+ are standardized directly or indirectly by a hydrogen electrode system. In order to note the effects produced by a

776

Neutral Salts on Phosphate Buffers

EXPERIMENTAL.

Preparation

of Xolutions.

Solutions of ~/7.5 Na2HP04, ~/7.5 KH2POI, and 10 per cent solutions (10 gm. of salt dissolved in water and made up to 100 cc.


substance on these simpler and under most conditions more convenient methods, it is necessary that their effects as measured by the hydrogen electrode be carefully considered. It was decided to study first the effects of addition of neutral salt to the phosphate buffer system because it is the one used extensively in biological studies, both as a buffering agent and as standards for calorimetric pH determinations with phenol red in the range of pH 7 to 8. S@rensen used the salts Na2HP04.2Hz0 and KHzPOd in ~/15 concentration. Some investigators have experienced trouble with Na2HP04. 2HzO in that it became more hydrated. The acid salt crystallizes out of solution easily and the anhydride is stable. Therefore, Clark and Lubs (2) recommended that the phosphate standards be prepared from KH2POI and standardized NaOH solution. Cullen (3) has used the ~/15 phosphate mixtures of Serensen. When, because of war conditions, it became impossible to import the Na2HP04.2H20 salt, he asked Merck and Company, Inc. to prepare it. They prepared instead the anhydrous salt which has been entirely satisfactory for many years. This method is more convenient for most biological laboratories as it does not involve the preparation and keeping of It has been the experience in pure, COz-free, sodium hydroxide. this laboratory that the salts as furnished by the above manufacturer are reliable and convenient for the preparation of the solutions. The values obtained for the pH of mixtures made from these salts have checked consistently within experimental errors the values as given by Hastings and Sendroy (4) when the system was standardized in the same manner. Myers and Muntwyler (5) have recently used these salts in the preparation of their standAs these ards to be used in the wedges of the Myers calorimeter. salts are being generally used, it seemed advisable to continue their use in the present study. Considerable new information has been obtained. An attempt has been made to interpret these results in accordance with the Debye-Hiickel concept of the influence of interionic forces (6, 7).

H. W. Robinson

777

Electrometric

pH Determination.

The electrometric determinations of hydrogen tion were made by the customary method with Saturated calomel electrode

Saturated KC1 bridge

Electrode liquid

ion concentra-

Hz-Pt

The present study was made at 38” for the reason that it was desirable to secure data at 38” to use in studying the correctionsnecessary in the use of phenol red for calorimetric pH determinations in serum and tissue fluid. Considerable data are available for room temperature while there is but little available for blood temperature. The temperature was controlled within f0.1”. The 2.5 cc. Clark-Cullen temperature-controlled electrode vessel was used with the saturated calomel electrode. Potential measurements were made with a Leeds and Northrup “hydrogen ion” potentiometer, enclosed lamp and scale galvanometer, and a Weston standard cell. The hydrogen used was furnished by the International Oxygen Company. The details of the method, preparation and care of electrodes were as described by Cullen (3). The individual pH determinations reported here are accurate to at least 0.005 PH.


in a volumetric flask) of NaCl and KC1 were prepared from Merck’s Blue Label reagents. Stock solutions of the buffer mixtures were made for each series of determinations. For example, when the effect of a concentration of NaCl of 0.5,1, and 2 per cent in a ~/15 phosphate mixture containing 8 parts of basic salt and 2 parts of phosphate acid salt was determined, a stock solution of a M/7.5 mixture containing that ratio of salts was first prepared. To each of four 100 cc. volumetric flasks, 50 cc. of the phosphate mixtures were added by means of a delivery pipette and to the respective flasks, no NaCl, 5 cc., 10 cc., and 20 cc. of a 10 per cent solution of NaCl. The flasks were then made up to mark with distilled water. In this manner one obtains an exact phosphate ratio in each of the flasks. Approximately the same procedure was carried out in the entire series of determinations indicated in Tables I to HI. In all cases the solutions were kept in Pyrex flasks.

778

Neutral

At the beginning pared from constant and Bonner (8), was tial measured. The

Salts on Phosphate Buffers of a series of determinations 0.1 N HCl, preboiling acid obtained by the method of Hulett placed in the electrode vessel and the potenpH of 0.1 N HCl was given the value of 1.08

and the e of the system

was calculated,

E.M.F.

pHs80 = 0~169

f3

Standardization

and Use of the Term pH.

0.1 N HCI, with the pH value assigned at 1.08, has been used as the final standard of reference as suggested by Cullen (3). The reasons for using this standard were again discussed by Cullen, Keeler, and Robinson (9). It is now widely used in American The value of 1.08 at 38” was obtained by biological laboratories. the use of the activity values for 0.1 N HCl given by Noyes and Ellis (10). Owing to the lack of information on the change of activity values with temperature the value 1.08 was adopted for all temperatures between 1540”. In practically all the work that has been done in the past with the use of M/15 phosphate buffers, the pH values given to the various mixtures are those of Sorensen. His values were obtained at 18” by using the Bjerrum extrapolation for liquid junction potential and assigning to hydrochloric acid a value based on conductivity measurements. It has been previously pointed out (Cullen, Keeler, and Robinson (9), Hastings and Sendroy (4)) that the values of t(he pH of ~/15 phosphate buffer mixtures at 20”, determined electrometrically without correction for liquid junction potentials and with the Cullen standardization, check within 0.01 pH unit the pH values given by Sorensen for 18”, and that at 38” the phosphate mixtures give a pH value approximately 0.03 pH more acid. The latter difference has been obtained repeatedly and used by Cullen, Keeler, and Robinson and by Hastings and Sendroy. It seems to be In constant for the entire range of the phosphate buffer mixture.


In a This was the method used for most of the determinations. few the value of e was determined with a phosphate mixture that In every case had previously been standardized with 0.1 N HCI. the standardizing solution was again run at the end of the series in order to make sure that no change had taken place in the system. Determinations were made in duplicate and refills were made until the readings of the two cells were constant and in agreement.

H. W. Robinson

779

RESULTS.

General Efects

of Neutral

Salt on pH

of Buffer

Solutions.

The influence of neutral salt on the hydrogen ion concentration of buffer solutions has been studied by many investigators (Michaelis and Kruger (15), Michaelis and Kakinuma (16), Leeper and Martin (17), Kolthoff and Bosch (18), Morton (19)). The conclusions reached have been the same. The effect of a given concentration of salt is increased if the concentration of buffer is decreased. The effect is not the same in different, buffer systems. Michaelis and Kakinuma find that the effect of NaCl upon the pH of 0.01 M acetic acid and 0.01 M sodium acetate is considerably less than Michaelis and Kruger found for phosphates.


view of the present lack of knowledge of the absolute value of the hydrogen ion activity in a reproducible standard solution, and since it is hoped that the present data will prove useful in connection with previous work, it seems logical to continue to use the HCI standard and to retain the term pH instead of adopting the term pan+ suggested by Serensen and Linderstrgm-Lang (11). At the present time there is a confusion in regard to t,his latter term. It is important to realize that the pH values given in this paper are identical with the pH values of Cullen, Keeler, and Robinson (9), the pH values of Hastings and Sendroy (4) (1924), the pan+ values of Hastings and Sendroy (12, 13) (1925, 1926) the pan+ values of Stadie and Hawes (14) (1928), but are not identical with the pan+ values of Sorensen and Linderstrom-Lang. Their present value for pan+ is about 0.04 pH units higher than the pH values given here or the pan+ values of the above authors. All the pH values referred to are those based on the hydrogen electrode. While we agree with the suggestion of Sorensen and Linderstr@m-Lang that the problem of standardization should be treated according to the Lewis-Bjerrum activity theory, it appears better to accept Clark’s suggestion (2) that the use of the term pH be continued. This avoids confusion in connecting new work with what has been done in the past. Hastings, Murray, and Sendroy (13) also emphasize the confusion in terms. The confusion that must be avoided is the apparent correction of values such as those of Cullen, Keeler, and Robinson by subtracting -0.04 when the difference is in the notation rather than in values.

780

Neutral

Salts on Phosphate

Buffers

The effect is much greater as the valence of the buffer increases. The cation effect seems to be the dominant factor in this change which increases in the order K < Na< Li. The present data show that these effects at 38” are similar to those found by the above

~PH

Change

I

of pH

in

Presence

of

Salt

at 38’


-#I --we-

Adztic

Acid-Acefate

Mixlure

Plosph~te,~fjxtureM/rS

750

a FIG. 1. Relative influence of various salts on the pH of ~/15 phosphate buffers at 38”; also comparative effect of NaCl on PO1 and acetate systems.

authors in the neighborhood of 18”. Fig. 1 shows the relative reduction, A pH, in the pH due to the presence of sodium chloride in phosphate and acetic acid-acetate mixtures when the salt-acid ratio is 1: 1, and also the comparative effects of sodium chloride,

H. W. Robinson

Variation

of pH Values with Salt-Acid Ratio and Concentration Phosphate Buffers.

Summaries

of the experiments

pH of phosphate

solutions

of

on the effect of the changes in

of various

ratios of sd

due to the pres-

ence of 1,2, and 3 per cent of sodium chloride and potassium chloride are given in Tables I and II. The concentration of PO1 was varied from ~/15 to ~/60. From Experiments 1 to 3 of Tables I and II it is evident that the effect of salt on the pH of a phosphate buffer containing the same ratios of di- and mono-salts is greater in the weaker phosphate solutions than in the stronger, also that the A pH for KC1 is less than for NaCl. From the experiments of equal PO1 concentrations but different salt-acid ratios, in Experiments 1,4, and 6 of Table II, it is evident that the change in pH is not the same throughout the entire range of buffer. Thus in the ~/15 PO4 system, 1 per cent NaCl gives a A pH of 0.178 when the 2 salt 8 salt ratio is 2 but a A pH of 0.220 when the f ratio is 3. - This acid acid difference far exceeds experimental error. With each total PO, salt cormentration the A pH increases as T ratio decreases. acid In analyzing these A pH values it was noticed that the A pH values for NaCl in a ~/15 phosphate mixture containing a ratio of NaJtP04 = 0.25 were approximately the same as that produced KHzPO4 by the same amounts of salt in a ~/30 phosphate solution con-


potassium chloride, lithium chloride, potassium bromide, and potassium sulfate on the 1:l phosphate mixture. The bromide and chloride effects are practically the same, which confirms what has been previously reported. Potassium sulfate cannot be called a true neutral salt and it is felt that its deviation is due to the secondary hydrolytic reaction SO*- + H+ = HSOa. In line with this observation, Kolthoff and Bosch found that the effect of K&04 on the H+ concentration in the carbonate-bicarbonate system was much less than that of KC1 although thecalculated ionic strength of the latter was only 0.5 and that of KzSO4 was 0.75, and Morton (19) also not’ed deviations with this salt.

Neutral

782

Salts on Phosphate Buffers TABLE

Change

in pH

of Phosphate

I.

Solutions

in

I

“Gzz m41 NO.

mols per 1.

[NMHPOII [XHtPOal

Lwr

/ WC11)

Presence

of NaCl. -

I

PH

j APHI

PK’

!BE$,

r

0.0667

4

7.372 0.602 0 6.770 0.173 0.086 7.267 0.105 6.665 0.258 0.171 7.194 0.178 6.592 0.344 0.342 7.090 0.282 6.488 0.515

0.416 / 0.508 , 0.586 I 0.717

2

0.0334

4

0.602 0 7.452 6.850 0.086 0.086 7.312 0.140 6.710 0.171 0.171 7.220 0.232 6.618 0.257 0.342 7.100 0.352 6.498 0.428

0.293 0.413 0.507 0.654

3

0.0167

4

0.602 0 7.515 6.913 0.043 0.086 7.335 0.180 6.733 0.128 0.171 7.235 0.280 '6.633 0.214 0.342 7.100 0.415 6.498 0.385

0.207 0.358 0.463 0.620

4

0.066i

1

o.ooo 0 6.794 6.794 0.133 0.086 6.682 0.112 6.682 0.219 0.171 6.607 0.187 6.607 0.304 0.342 6.495 0.299 6.495 0.475

0.365 0.468 0.551 0.689

5

0.0334

0.25

-0.602 0 6.307 6.909 0.046 0.086 6.135 0.172 6.737 0.132 0.171 6.045 0.262 6.647 0.217 0.342 5.902 0.405 6.504 0.388

0.214 0.363 0.466 0.622

6

0.0667

0.25

-0.602 0 6.232 6.834 0.093 0.086 6.096 0.136 6.698 0.178 0.171 6.012 0.220 6.614 0.264 0.342 5.890 0.342 6.492 0.435

0.305 0.422 0.514 0.659

7

0.0667

0.954 0 7.708 6.753 0.186 0.086 7.611 0.097 6.657 0.271 0.171 7.543 0.165 6.589 0.357 0.342 7.434 0.274 6.480 0.528

0.431 0.520 0.597 0.726

8

0.0334

0.954 0 7.789 6.835 0.093 0.086 7.651 0.138 6.697 0.178 0.171 7.567 0.222 6.613 0.264 0.342 7.449 0.340 6.495 0.435

0.305 0.422 0.514 0.659

-


1

783

H. W. Robinson TABLE

Experi merit No.

[pod

h3r

[NaCll

pH

A pH

PK'

&$h

------

Wd8

pa 2.

9

[NazHPOaI [KFLPOII

I-~OnChded.

0.066i

r

7.48

m01.3 per 1.

P

0.874 0 0.086 0.171 0.34i

6.759 0.184 7.633 7.538 0.095 6.664 0.269 7.468 0.165 6.595 0.355 7.358 0.275 6.485 0.526

0.429

0.518 0.596 0.725

0.0667

6.69

0.826 0 0.086 0.171 0.342

7.581 6.755 0.183 7.483 0.098 6.657 0.268 7.411 0.170 6.585 0.354 7.306 0.275 6.480 0.525

0.427 0.518 0.594 0.724

11

0.0667

2.33

0.368 0 0.086 0.171 0.342

7.142 6.774 0.160 7.040 0.102 6.672 0.245 6.964 0.178 6.596 0.331 6.854 0.288 6.486 0.502

0.400 0.495 0.575 0.708

Na2HP04 taining a ratio of KIIP04 = 9 (seeExperiments 6 and 7, Table I).

These mixtures have the same ionic strength. “Ionic strength” is a concept introduced by Lewis and Randall (20) when studying activity coefficients in dilute solutions. They make the generalization that “in dilute solutions the activity coefficients of a given strong electrolyte are the same in all solutions of the same ionic strength.” In the present paper the symbol p, called ionic strength, is defined in mols per liter of solution instead of mols per 1000 gm. of water. p = t (m121~ + m222+ m322+ . . . . . . . . . .etc.) ml, m2,ms... . . . . . . . = molsper liter of ions, Al, AZ, Aa,. . . . . . . . 21, 22, 23.. * . . . . . . . . . = valence of ions, Al, As, At.. . . . . . . . . . . . . . This is the quantity that is employed in the Debye-Hiickel equation. Disodium hydrogen phosphate and dihydrogen potassium phosphate are strong electrolytes and in the dilute solutions are assumed to be completely ionized. The basic salt gives rise to three ions Bf, B+, and HP04”, whereas the acid salt gives rise to two ions Bf and HzP04- (when B = Na or K). From the above relationship it can be calculated that the ionic strength of a


10

TABLE

Change in E;h”t NO.

[PO41

- nol.3

pH

[NazHPOJ [KHsPOII r

per 1.

II.

of Phosphate Solutions in Presenceof KCl. Log +

[KC11

PH

APH

PK'

st;$l

-- ----mols per 1.

P

0.0667

2.57

0.411 0 0 7.178 6.768 0.163 0.067 7.112 0.066 6.702 0.229 0.134 7.063 0.115 6.653 0.296 0.268 6.994 0.184 6.584 0.431

0.403 0.479 0.544 0.656

2

0.0334

2.57

0.411 0 0 7.267 6.857 0.082 0.067 7.163 0.104 6.753 0.148 0.134 7.100 0.167 6.690 0.216 0.268 7.017 0.250 6.607 0.350

0.286 0.385 0.464 0.591

3

0.0167

2.57

0.411 0 0 7.340 6.930 0.041 0.067 7.196 0.144 6.786 0.108 0.134 7.123 0.217 6.713 0.175 0.268 7.030 0.310 6.620 0.309

0.202 0.328 0.418 0.556

4

0.0667

1.00

0.001 0 0 6.789 6.789 0.133 0.067 6.710 0.079 6.710 0.200 0.134 6.661 0.128 6.661 0.268 0.268 6.591 0.197 6.591 0.402

0.365 0.447 0.517 0.634

5

0.0334

0.25

-0.602 0 6.304 6.906 0.046 0.067 6.171 0.133 6.773 0.113 0.134 6.097 0.207 6.699 0.180 0.268 6.007 0.297 6.609 0.314

0.214 0.336 0.425 0.561

0

0.0667

0.25

-0.60:2 0 6.232 6.834 0.093 0.067 6.137 0.095 6.739 0.160 0.134 6.076 0.156 6.678 0.227 0.268 5.996 0.236 6.598 0.361

0.305 0.400 0.477 0.601

7

0.0334

9.00

0.954 0 7.789 6.835 0.093 0.067 7.693 0.096 6.739 0.160 0.134 7.635 0.154 6.681 0.227 0.268 7.555 0.234 6.601 0.361

0.305 3.400 3.476 0.600

8

0.0667

7.47

0.874 0 7.633 6.759 0.184 0.067 7.570 0.063 6.696 0.251 0.134 7.525 0.108 6.651 0.319 0.268 7.455 0.178 6.581 0.453

0.429 0.501 0.564 0.673

9

0.066'7

0.20

-0.69:8 0 6.127 6.825 0.089 10.067 6.033 0.094 6.731 0.156 I0.134 5.973 0.154 6.671 0.223 I0.268 5.891 0.236 6.589 0.357 734

0.298 0.395 3.472 9.597

-


1

H. W. Robinson

785

phosphate mixture is equal to the molal concentration of the uni-univalent BH2P04 and 3 times that of the uni-bivalent B~HPOI. A phosphate mixture consisting almost entirely of BgHPOr will


Fm. 2. Influence

of p of phosphate

on NaCl effect on phosphate

pH.

have a much higher ionic strength than a mixture that consists mostly of BHzPOl although the total PO4 concentration is the same. The changes in the concentrations of the ions HPOd and HzP04- due to the change of the hydrogen ion concentration

786

Neutral

Salts on Phosphate Buffers

(HPOh= + H+ 5 HzP04-) are negligible overthe pH range studied here because the changes in H+ are negligible in comparison to the concentration of phosphate ions. The last column in Tables I and II gives the square root of the ionic strength of the solutions. It is clearly shown by the results that the change in pH due to the presence of a particular electrolyte in a phosphate mixture is

bp~o &f’f’ect

of KC/ on PO, &dV’et-

pH Downloaded from http://www.jbc.org/ by guest on March 17, 2015

FIG. 3. Influence

of p of phosphate

on KC1 effect on phosphate

pH.

related to the ionic strength of the buffer solution and not directly to the total PO4 concentration. The A pH varies over the entire range of the buffer because the ionic strength varies with the ratio of uni-bivalent BzHP04 in&univalent BH2P04 ’ Figs. 2 and 3 show the A pH effects of NaCl and KC1 when present up to concentrations of 0.3 N in solutions of phosphate mixtures of known ionic strength.

H. W. Robinson Interchangeability

787

of K and Na in the S@ensen PO, System.

In the above experiments

with the S#rensen PO1 system, two cations, sodium and potassium, are present. It was necessary to see whether the effects were different when only one cation was present. The experiments were repeated with KzHPO*-KH2P04 mixtures. The latter was made by adding standardized potassium hydroxide to the KH2P04, so that the total concentration of In Table III a typical experiment is given. PO4 was known. III.

Solution w-J1

A.

PB

Solution API-I

[KC11

0

---

APH

N

N

0.067 0.171 0.342

PB

B.

7.789 7.693 7.635 7.555

0

0.096 0.154 0.234

0.067 0.171 0.342

7.727 7.632 7.573 7.493

0.095 0.154 0.234

This experiment is given because it represents the condition at higher pH, when there is the greater proportion of Na. From the results the differences in A pH obtained from the presence of KC1 in the K2HPOa-KH2P04 mixtures are, within experimental error, the same as those obtained in the Na2HPOk-KH2P04 mixtures. DISCUSSION.

Efect of Salt on the Second pK’ of Phosphoric Acid.

About 20 years ago Henderson (21) and Washburn (22) simultaneously characterized the equilibrium of a buffer solution in terms of the law of mass action, assuming that the concentration of the negative ion approached the total concentration of salt and that the concentration of the undissociated acid was equal to the total concentration of the acid. The general expression is [H+] = K’s


TABLE

Showing That Depression of pH by KC1 Is the Same in the Systems NazHPOd and KzHPO~ -. KHzPOd KHzPO, Total concentration of POS in both cases is 0.03334 M. Solution A = mixture of Na*HPOd and K~HPOI. ‘I B= “ “ KtHPOa + KOH.

Experiment

788

Neutral

Salts on Phosphate Buffers

Influence

I

I

0.2

0.3

.

3.

.

Difufion

I

I

0.4

0.5

FIG. 4. al, u2, u3, ad, ~/60 POc containing 0, 0.5, bl, bz, bs, b4, ~130 IL ” 0, 0.5, Cl, Cz, Cs, C4, M/15 “ “ 0, 0.5, Line clbla,, dilution with water of n1/15 ~2, bz, ~2, ~/60, ~/30, ~/15 PO* containing as, bf, c3, ~/60, 430, ~/15 “ “ a+ 64, ~4, ~/60, ~/30, ~/15 “ “

on pK’

I

1

II

I I

I

r

1, 2 per cent NaCl respectively. 1, 2 “ “ “ “ 1, 2 “ “ “ “ PO4 to ~/30 and ~/60. 0.5 per cent (0.086 N) NaCl. 1.0 “ “ (0.172 “) “ 2.0 “ “ (0.344 1‘) “

I


\I\

ffa C/and

of

H. W. Robinson

789


in which the bracketed expressions represent the concentration of hydrogen ion, weak acid, and the salt of the weak acid and K’ is the proportionality constant, called the apparent dissociation constant, which approaches the dissociation constant, K, with dilution. With the general use of Sorensen’s term pH, Hasselbalch first used the above relationship in its logarithmic form, pH = pK’ + [salt] This equation applied to dilute phosphate mixture is log __ [acid]’ only an approximation. Michaelis and Kruger showed t,hat on dilution of a phosphate mixture with water the reaction became more alkaline. This would not occur if the reaction were only due to the ratio of the constituents and the pK’ remained constant. Cohn (23) recalculated Sorensen’s E.M.F. measurements of phosphate mixtures and shows that the variation in pK’, while small in comparison with the change in pH, is 10 times greater than the probable error. The pK’ changes with the ionic strength of the solution and will approach a true pK value at infinite dilution when the phosphate mixture obeys the simple laws of an ideal solution. With phosphate mixtures this change of pK’ with the ionic strength of the solution is not a st,raight line. This is observed when the data of Table I or Table II are plotted in a graph. This fact also is shown over a wider range of concentration in Fig. 2 of Cohn’s paper. When neutral salt is added to the phosphate buffer system, the curves showing change in pK’ are distinctly different for varying quantities of saIt. pK’ vaIues are Iower and the curves flatter with increase of salt concentration. This is shown in Fig. 4 where the pK’ values of Experiments 1 to 3 of Table I are plotted against ,uLtotai. The line c1czc3c4(Fig. 4) represents the change in pK’ with ionic strength in a ~/15 phosphate mixture due to the presence of sodium chloride. The line cl&al represents the change of pK’ with ionic strength by dilution of the same ~/15 phosphate mixture up to ~/60 with water. The same curve would have been followed if the ionic strength of the ~/15 phosphate mixtures had been changed by changing the ratio of basic salt to acid salt. The x points are points taken from other experiments in Table I. The points uz, bz, and q represent M/~O, ~/30, and M/E respectively,

790

Neutral

Salts on Phosphate

Buffers

Application

of Debye-Hiickel Containing

Equation to Phosphate Neutral Salt.

Bu$ers

Cohn (23) has studied the phosphate buffer solutions according to the modern conception of act,ivity and has been able to account for the deviations by means of the Debye-Hiickel equation, which assumes that in completely ionized strong electrolytes the deviation from the ideal law of mass action in dilute solutions is caused mainly by the interionic electrical effects. The mass law equation for the second dissociation constant for phosphoric acid may be written aH+

aHPOr

=K

(1)

aH2POr

where (In+ = activity of H+, onpoe’ = activity of HPOaZ, and o&pod- = activity of HzPOd-; or in logarithmic form pH -

%po4=

log -

%*PO,-

= PI