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.