J. Biochem., 81, 163-168 (1977)

Measurement of the pH of Frozen Buffer Solutions

by Using pH Indicators

Yutaka OR‡U and Masayuki MORITA

Department of Biology, Faculty of Science , Osaka University, Toyonaka, Osaka 560

Received for publication, May 4, 1976

A method was established to estimate the pH change of several buffer solutions on freezing

by using a combination of pH indicators. Among more than 30 buffer solutions examined , almost half exhibited a pH change in the temperature range between freezing point and 220•‹K;

the results were tabulated. Glycerol was found to suppress the pH changes because of its "salt buffer" effect .

Previously Orii and lizuka (1) found that the addi precipitating compounds. Thus, the pH of the tion of a carbonyl reagent such as cyanide or unfrozen portion of slightly alkaline sodium phos

sodium bisulfite to the formyl side group of heme phate buffers at room temperature dropped to as a was accelerated appreciably when an alkaline low as pH 3.6 at -9.9•‹ (eutectic point). Since reaction mixture prepared in 1% aqueous sodium then, Chilson et al. (4) have observed activity loss carbonate was frozen, and ascribed this to a lower of aldolase [EC 4.1.2.13] and various dehydro

ing of the effective pH, by at least 4 pH units, of the genases in solutions of sodium or potassium phos

unfrozen portion where the reaction took place. phate on freezing and thawing, and based on the Variation in hydrogen-ion and salt concentra findings of van den Berg and Rose, ascribed the tions in the aqueous phase of frozen muscle juice inactivation to pH changes as well as the concentra was recognized by Finn in 1932, and he proposed tion effects of the salts. Other than in these ex that these two factors were responsible for the ceptional cases, however, the pH changes of salt denaturation of proteins contained in the juice (2). solutions on freezing have not necessarily been In 1959 van den Berg and Rose noted a marked pH appreciated. drop of mixtures of mono- and disodium phosphate In 1949, Keilin and Hartree (5) used low-tem- on freezing (3). They analyzed the composition perature spectroscopy to observe the cytochrome at room temperature and measured the pH of liquid bands in several tissues; these bands were sharp portions which had been separated from the solids ened, intensified and split at the temperature of at low temperatures. The analyses indicated that liquid air. They also observed a transition of alka as the temperature of a mixture of the two salts was line methemoglobin at pH 9 and room temperature lowered, ice and less-soluble disodium salt were to the acid form on cooling. They interpreted this precipitated, and that consequently the composition as due to the suppression of ionization of a water of the unfrozen portion approached that of an molecule bound to the sixth octahedral coordination eutectic mixture, the pH and composition of which position of methemoglobin at low temperatures. were determined by the relative solubilities of the Keilin and Hartree also observed color changes on

Vol. 81, No. 1, 1977 163 164 Y. OR‡U and M. MORITA cooling solutions of some pH indicators such as priate pH values as described in the legend to phenolphthalein, phenol red, methyl red, and Table I. GOOD buffers were also prepared as bromthymol blue, and explained these changes on indicated (12). the basis of a suppression of dissociation of these The low temperature spectra were measured dyes, which are naturally weak bases and acids. as described previously (13). The media they used were not described, but may Color changes on freezing buffer solutions have been sodium phosphate buffers, as usual. If containing pH indicators were usually observed as so, all of their findings can be ascribed to a pH fall follows: to 2 ml of a buffer solution in a test tube on freezing; the significance of this in determining of 10 mm diameter were added one to two drops the essential properties of hemoproteins has not of a pH indicator solution, and the bottom of the been recognized explicitly since then (6-8, but see, test tube was placed on the surface of liquid nitrogen for example Refs. 9, 10). in an insulated vessel. As the solution was frozen For Mossbauer studies and measurements of the tube was dipped into the liquid nitrogen, and NMR, EPR, and optical spectra of hemoproteins the color change was recorded. and related substances at cryogenic temperatures,

knowledge of the actual state of samples is abso RESULTS AND DISCUSSION lutely necessary in interpreting the data, and it

should be recognized that the effects of pH changes Preparatory Measures for Use of pH Indicators

are by far more pronounced than might be expected, -When monitoring pH changes of salt solutions

depending on the type of media and handling of the on freezing by use of pH indicators, possible ex

samples. Also, suitable media should be selected trinsic color changes due to a temperature-depend

for the storage of biological materials in frozen ent shift in pH indicator constant, salt errors and

states, in order to prevent acid or alkali denatura different solubilities of different forms of the indi tion during cooling or even during storage at around cators must be considered. -20•‹ , the temperature of most conventional freez Generally, the change in dissociation constant ers. Therefore, we examined the pH changes on of most weak acids (bases) with temperature is

freezing several buffer solutions currently used by small, and data presently available indicate that the observing color changes of a group of pH indi apparent pH change of some indicators on raising

cators. The results may be valuable to investiga the temperature from 18 to 70•‹ is in the range of

tors in the field of biology, biochemistry, biophysics, about -0.4 pH unit at most (14). Although no

and other related topics. For other salt solutions data are available for the change on cooling to which are not included in this report, the present cryogenic temperatures, it may be about the same

simple method can be applied to investigate the pH extent but with opposite sign. This judgement is

changes. supported by the observation that aqueous solu tions of potassium chloride containing an indicator

showed no color change except in the cases of MATERIALS AND METHODS bromthymol blue (BTB) and phenol red (PR), pH Indicators used could be grouped as 1) phtha which exhibited some alkaline colors due to an leins or 2) sulfonaphthaleins or 3) azo dyes, as de acid-base error. Therefore, the temperature de scribed in the legend to Table I. They were selected pendence of color changes seems negligible in so as to cover a pH range between 0.5 and 10.5, and practice. the indicator solutions were prepared as described The salt error due to concentration of salts in in Ref. 11. Two types of universal indicators were the liquid phase of frozen solutions might be serious, also used for convenience. One was a mixture of if it occurs, since this displaces the equilibrium be methyl yellow, methyl red, bromthymol blue, tween acid and base forms of indicators. Accord thymol blue, and phenolphthalein as described in ing to Kolthoff (14) colorimetric measurements Ref. 11, and covered the pH range between 1 and with an indicator should yield too high a pH as 10. The other was a universal indicator solution salt concentrations become higher, especially when (pH 4-10) obtained commercially from Merck. acidic indicators are used. Although only BTB Buffer solutions were prepared to give appro and PR among the indicators employed showed

J. Biochem. pH CHANGES OF FROZEN BUFFER SOLUTIONS 165 this tendency, color change of BTB was ascribed to the acid-base error since it was observed in weakly buffered solutions. PR became red after freezing aqueous potassium chloride made acidic with HCI. Therefore, BTB was used with some reservations, and PR was omitted from routine use. Color changes which could be ascribed to dif fering solubilities of the two forms of an indicator were not observed. Consequently, the indicators used in the present investigations appeared to be practically free from extrinsic color changes for the present purpose. Validity of pH Indicators as Monitors of pH Changes on Freezing of Buffers-In order to con Fig. 1. Absorption spectra of mixed pH indicators in firm the previous conclusion that the effective pH frozen 1o% aqueous sodium carbonate. To 2.0 ml of 1% of aqueous sodium carbonate solution was lowered aqueous sodium carbonate was added 0.015 ml of a appreciably by freezing (1), the pH change was universal pH indicator mixture containing methyl yel followed as a spectral change of pH indicator(s) low, methyl red, bromthymol blue, thymol blue, and added to the solution. phenolphthalein (see Ref. 11). This mixture was To 2 ml of 1 % aqueous sodium carbonate was transferred to the sample cuvette of a twin-cuvette as added 0.015 ml of a universal pH indicator mixture, sembly; carbonate solution without indicators was prepared as described in " MATERIALS AND placed in the reference cuvette. The cuvette assembly was then dipped into liquid nitrogen, and the absorption METHODS." As the pH of this solution was low spectra were recorded at the temperatures indicated. ered from the original pH of 10.79 by dropwise The estimated pH values based on spectral comparison addition of 1-6 N HCI, the spectrum was recorded are also given. -, 151•‹K and pH 3.5; -, 223•‹K at each pH at room temperature. The color of and pH 6.7; -, 247•‹K and pH 8; and -, 286•‹K and

the original solution was purple, and this finally pH 10.8. The ordinate shows arbitrary units. changed to pale orange red at pH 3.33. Accord ingly, characteristic spectra were recorded corre spondine to each pH. pH changes in the following investigations. The original solution was transferred to a sam pH Changes of Burffer Solutihns on Freezing- ple cuvette in a twin-cuvette assembly for low The pH of buffer solutions to which a pH indicator temperature spectrophotometry and carbonate solu had been added did not change as checked with a tion without pH indicators was placed in a reference pH meter, and the pH values determined color cuvette. The cuvette assembly was dipped into metrically coincided with those determined electro liquid nitrogen and the sample solution was frozen. metrically within the experimental accuracy. The absorption spectrum was recorded, then the In Table I the color changes of each buffer sample temperature was raised slowly with a heater on freezing are summarized. Among the buffers and the spectra were recorded several times at examined, the pH drop on freezing was significant intervals. From a comparison of these spectra (more than 3 pH units) with (borax-monopotas with those obtained at room temperature at various sium phosphate), (boric acid-NaOH), (sodium pH's, the effective pH's of the frozen solution at carbonate-sodium bicarbonate), and (monosodium various temperatures were roughly estimated as phosphate-disodium phosphate) buffers. (Mono summarized in Fig. 1. potassium phosphate-disodium phosphate) showed

The effective pH changed, especially between a smaller drop (around 2 pH units), and (barbital

220•‹K and room temperature, in accordance with sodium-sodium acetate-HCI), (monopotassium

the previous observations (1). The observed color phosphate-dipotassium phosphate), and (glycine

change paralleled the spectral change during the - NaOH) showed a pH drop of about 1 pH unit. On

transition, so the former was employed to follow the other hand, (Tris-acid maleate-NaOH), (Tris-

Vol. 81, No. 1, 1977 166 Y. ORII and M. MORITA

TABLE I. pH Changes of buffer solutions on freezing. Color changes of pH indicators were recorded as described in the text. In the column under each pH indicator the color before freezing is given on the left of the hyphen and that after freezing on the right. A, B, and AB represent acidic, basic, and intermediate colors, respec tively. pH Changes estimated colorimetrically by using universal pH indicators are expressed in a similar way.

J. Biochem. pH CHANGES OF FROZEN BUFFER SOLUTIONS 167 nitric acid), and (acid sodium maleate-NaOH) concentrations of glycerol were added to 0.1m showed a pH increase of around I pH unit. monosodium phosphate-disodium phosphate (pH GOOD buffers except for CAPS generally 6.85), and the mixture was frozen after adding pH showed a pH increase of 1 unit on freezing, based indicators. Even with 1 % (v/v) glycerol the pH on the color change with BTB. Although this drop was appreciably reduced, since the lowest pH dye showed a basic color on freezing in weakly buf value registered colorimetrically was pH 5; 2.5% fered solutions because of an acid-base error, this glycerol was even more effective. However, above error should be negligible in well-buffered systems this concentration no further improvement was and the effective pH changes observed seem achieved. In the presence of 5 % glycerol the pH of reasonable. sodium carbonate-sodium bicarbonate (pH 10.0) Consequently, the use of buffers showing a dropped only to pH 9 on freezing, compared to large pH change on freezing should be avoided for pH 6 in its absence. Lovelock already suggested experiments at cryogenic temperatures or storage of that glycerol acts as a "salt buffer" to reduce the

biological materials in the frozen state. However, salt concentration effect when solutions are frozen

if the time spent by biological samples in the tem (15). The present result is in conformity with this

perature range between freezing point and 220•‹K, mechanism and the preferential concentration of where liquid phase still exists (1), is short (rapid one of the component salts of buffers on freezing

freezing and thawing), damage would be minimized. seems to be appreciably inhibited by the presence of

On the other hand, selective denaturation of some glycerol.

proteins could be carried out by slow freezing and However, when glycerol was added to boric thawing in these buffers. acid- borax (pH 8.5), a peculiar phenomenon oc As an application of the present technique, the curred; in the presence of 5 % glycerol the pH meter effect of adding glycerol to buffer solutions was registered pH values between 6.8 and 6.3 even at examined. Glycerol as well as sugars and other room temperatures, and on freezing the pH dropped additives are generally used to protect biological to as low as pH 2 to 3, as determined colorimetri materials from damage caused by freezing and cally. This phenomenon might be due to esterifi thawing, and this protective effect could be ascribed cation of glycerol with boric acid, since the addition to overcoming pH changes, if they occur. Different of glycerol to other buffers such as succinic acid-

Footnotes of Table I * After freezing , CPR showed a pink color different from the ordinary basic purple color. ** Faint blue color remained. The pH indicators used were metacresol purple (MCP, acidic and alkaline, ‡U), methyl yellow (MY, ‡V), bromcresol green (BCG, ), chlorophenol red (CPR, ‡U), bromthymol blue (BTB, ‡U), phenolphthalein (PP, I), and thymol phthalein (TP, I), where I, ‡U, and ‡V indicate phthalein, sulfonphthalein, and azo dyes, respectively.

Universal indicator solution (UniA) was obtained commercially from Merck and UniB was prepared as described in Ref. 11. Buffer solutions of various pH values as indicated were prepared from the following stock solutions. 1, 0.2 M KCI and 0.2 N HCI; 2, 0.1 m glycine and 0.1 M NaCl in 0.1 N HCI; 3, 0.1 M potassium biphthalate and 0.1 N HCI; 4, 0.05 M borax and 0.05 M succinic acid; 5, 0.2 M acetic acid and 0.2 M sodium acetate; 6, 0.1 M citric acid and 0.2 M disodium phosphate; 7, 0.2 M succinic acid and 0.2 N NaOH; 8, 0.1 M potassium biphthalate and 0.1 N NaOH; 9, 1/7 M barbital sodium, 1/7 M sodium acetate and 0.1 N HCI; 10, 1/15 M monopotassium phosphate and 1/15 M disodium phosphate; 11, 0.2 M Tris acid maleate and 0.2 N NaOH; 12, 0.05 M borax and 0.1 M monopotassium

phosphate; 13, 0.2 M Tris and 0.2 N HCI; 14, 0.1 M barbital sodium and 0.1 N HCI; 15, 0.1 M boric acid, 0.1 M KCI, and 0.1 N NaOH; 16, 0.2 M sodium carbonate (anhydrous) and 0.2 M sodium bicarbonate; 17, 0.1 M glycine, 0.1 M NaCI, and 0.1 N NaOH; 18, 0.1 M citric acid and 0.1 M sodium citrate; 19, 0.2 M acid sodium maleate and 0.2 N

NaOH; 20, 0.2 M boric acid and 0.2 M borax; 21, 0.1 M monosodium phosphate and 0.1 M disodium phosphate; 22, 0.1 M monopotassium phosphate and 0.1 M dipotassium phosphate; 23,0J M Tris and 0.1 N HNO3; MOPS, 0.1 M 3-(N-morpholino)propanesulfonic acid and 0.1 N NaOH; MES, 0.1 M 2-(N-morpholino)ethanesulfonic acid and 0.1 N NaOH; TES, 0.1 M tris(hydroxymethy])methyl-2-aminoethanesulfonic acid and 0.1 N NaOH; HEPES, 0.1 M N-2-hydroxyethylpiperazine-N'-2-ethanesulfonic acid; TAPS, 0.1 M tris(hydroxymethyl)methylaminopropanesulfonic acid and 0.1 N NaOH; Tricine, 0.1 M N-tris(hydroxymethyl)methylglycine and 01 N NaOH; Bicine, 0.1 M N,N-bis-

(hydroxyethyl)glycine and 0.1 N NaOH; CAPS, 0.1 M cyclohexylaminopropanesulfonic acid and 0.1 N NaOH.

Vol. 81, No. 1, 1977 168 Y. ORII and M. MORITA

NaOH and succinic acid-borax did not affect the 7. Ehrenberg, A. & Estabrook, R.W. (1966) Acta pH values in this way. Chem. Scand. 20, 1667-1672 Unlike the pH change on freezing, the change 8. Curti, B., Massey, V., & Zmudka, M. (1968) J. Biol. in effective ionic activity in super-cooled mixtures Chem. 243, 2306-2314 -of salt solutions and glycerol has been investigated 9. Yonetani, T., Wilson, D.F., & Seamonds, B. (1966) J. Biol. Chem. 241, 5347-5352 extensively by Douzou et al. (16, 17). 10. lizuka, T. & Kotani, M. (1969) Biochim. Biophys. Acta 181, 275-286 The authors thank Dr. T. lizuka for the use of a spectro- 11. Sober, H. (ed.) (1970) Handbook of Biochemistry 2nd photometer at cryogenic temperatures. ed., The Chemical Rubber Co., Cleveland 12. Good, N.E., Winget, G.D., Winter, W., Connolly, REFERENCES T.N., Izawa, S., & Singh, R.M.M. (1966) Biochem 1. Orii, Y. & Iizuka, T. (1975) J. Biochem. 77, 1123- istry 5, 467-477 1126 13. Hagihara, B. & lizuka, T. (1971) J. Biochem. 69, 2. Finn, D.B. (1932) Proc. Roy. Soc. London Bill, 355-362 396-411 14. Kolthoff, I.M. (1937) Acid-Base Indicators The 3. van den Berg, L. & Rose, D. (1959) Arch. Biochem. Macmillan Company, New York Biophys. 81, 319-329 15. Lovelock, J.E. (1953) Biochim. Biophys. Acta 11, 28- 4. Chilson, O.P., Costello, L.A., & Kaplan, N.O. 36 (1964) Federation Proc. 24, suppl. 15, 55-65 16. Hui Bon Hoa, G. & Douzou, P. (1973) J. Biol. Chem. 5. Keilin, D. & Hartree, E.F. (1949) Nature 164, 254- 248,4649-4654 259 17. Douzou, P. (1973) Mol. Cell. Biochem. 1, 15-27 6. Camerino, P.W. & King, T.E. (1965) Biochim. Biophys. Acta 96,18-27

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