Specific-conductance determinations are made in cells of predetermined cell constant at a temperature most convenient for the locality, by measuring, as described above, the resistance of a solution of the product in equilibrium water and substituting this value in eq 63. The value of the specific conductance of the equilibrium water should be subtracted from the value found for the solution before the calculation of the C-ratio is made.

If the solutions are heated or subjected to vacuum, the value of Koly will be changed from loss of CO2; if high precision is required, the value of Koly should be redetermined on a sample of the equilibrium water treated in the same manner as the solution.

The concentration of the solution recommended by different observers varies from 2.5 g of dry substance per 100 ml to 50 g per 100 ml. Since the conductivity of solutions of beet-sugar products passes through a maximum at 25 g of sucrose per 100 ml of solution [8], measurements made near this concentration will be affected less by slight errors in concentration than at other concentrations [40]. Regardless of what concentration is selected as being the most suitable, it should be the same as that used for the determination of the C-ratio.

(b) ZERBAN AND SATTLER CONDUCTANCE METHOD FOR ASH

Although the C-ratio method may be used to determine the ash content of sugars from the same source in any one season, the ratio determined for one district cannot be used for another district. This fact has resulted in the Zerban and Sattler method for ash determination, based on conductivity measurements under special conditions, and no reference need be made to the gravimetric ash.

In developing this method, they minimize all causes for variability of the C-ratio except that resulting from the composition of the dissolved salts. The inorganic salts in solution in raw sugars are principally sulfates and chlorides. Since the equivalent conductance of solutions containing anions of inorganic salts is considerably higher than that of solutions containing anions of organic salts derived from the same base, it follows that the specific conductance of a solution containing a large percentage of chlorides or sulfates in the ash is relatively greater than that of a solution containing a small percentage. Furthermore, the C-ratio will decrease as the percentage of inorganic anions in the ash increases.

If hydrochloric acid is added to a solution containing only inorganic anions, the specific conductance will increase linearly with the addition of acid. However, if the solution contains both organic and inorganic anions, the specific conductance will increase linearly only after the weaker inorganic anions have been displaced by those of the added acid.

A similar relationship exists between the specific conductance of solutions containing inorganic and organic cations if a solution of potassium hydroxide is added. However, except when the sugar product has received treatment with bone black, this latter relationship is not pronounced and may be ignored.

From these relationships, Zerban and Sattler have developed three general methods for the determination of ash applicable to: 1. Raw cane and soft sugars [41].

2. Refinery sirups and molasses produced without char treatment [42].

3. Raw and refinery sirups and molasses of unknown origin [43]. (1) RAW CANE AND SOFT SUGARS [40].-Twenty-five grams of sugar is dissolved in equilibrium water and the solution diluted to 500 ml at 20° C. The specific conductance, k, of one portion of this solution is determined and likewise, the specific conductance, k1, of another portion to which 5 ml of 0.25 N hydrochloric acid has been added to each 200 ml. Corrections are made for temperature and the conductivity of the water. The corrected specific conductances multiplied by 106 give, respectively, K for the original solution, and K1 for the acidified solution. The percentage of ash may then be computed from the empirical equations,

Raw sugar, percentage of ash=0.001757 (0.913K+193.5–0.1K1)

(74) (75)

(76)

Soft sugar, percentage of ash=0.001695X (0.913K+193.5-0.1K1) The temperature corrections for k and k, are given by the equations (k)=(k)20 [1+0.02234 (t-20)+0.0000885 (t-20)2] (k1),= (k1)20[1+0.01704 (t−20)+0.000062 (t−20)2], in which t is the temperature at which the conductivity determination is made. The correction for (k), is roughly 2.2 percent per degree centigrade if measurements are made near 20° C.

(77)

The concentration of the acid may be checked by conductivity determinations. When 5 ml of the 0.25 N acid is mixed with 200 ml of equilibrium water the corrected specific conductance is 0.002370 at 20° C.

(2) REFINERY SIRUPS AND MOLASSES PRODUCED WITHOUT CHAR TREATMENT [41].-A solution is made of 100 ml of hot equilibrium water and 25 g of sirup or molasses. It is filtered with vacuum through asbestos and filter-paper pulp into a 200-ml volumetric flask with repeated washing with hot equilibrium water. The filtrate is mixed thoroughly and diluted with equilibrium water to 200 ml at 20° C. To a 20-ml portion of the filtrate is added 22.5 g of pure tablet sugar. This is diluted to 500 ml at 20° C with equilibrium water. This is known as solution A. To 200 ml of solution A, 5 ml of 0.25 N hydrochloric acid is added. This is known as solution B.

The specific conductances are determined in both solutions A and B and are corrected for solvent, specific conductance of the tablet sugar, and temperature. These values multiplied by 106 are respectively K and K1, which may be substituted in the equation

Percentage of ash=0.001757 (9.13K+1935-K1).

(78)

(3) RAW AND REFINERY SIRUPS AND MOLASSES OF UNKNOWN ORIGIN [42]. The procedure is the same as that used for the preceding determinations except that three conductivity measurements are required. Normal orthophosphoric acid is added to solution A to obtain solution B. The specific conductance at 20° C is determined on each of the following solutions:

1. k of solution A, as prepared in the preceding section.

2. k1 of solution A, to which has been added 5 ml of 0.025 N potassium hydroxide per 200 ml.

3. k of solution A, to which has been added 5 ml of normal orthophosphoric acid per 200 ml.

Each of the three specific conductances so determined are corrected for the specific conductance of the equilibrium water used and for the specific conductance of the tablet sugar. The corrected values multiplied by 10° are respectively K, K2, and K. The percentage of ash is then determined by substituting these values in the equation Percentage of ash=0.0191369K-0.002249K2-0.001210 K2+3.07.

(79) The concentration of the acid may be checked by conductivity determinations. When 5 ml of the normal orthophosphoric acid is added to 200 ml of equilibrium water the corrected specific conductance is 1925 X 10-6 at 20° C.

5. REFERENCES

[1] Hugh Main, Int. Sugar J., 11, 334–9 (1911).

[2] A. E. Lange, Z. Ver. deut. Zucker-Ind. 60, 359-81 (1910).

[3] F. Toedt, Chem.-Ztg. 49, 656-7 (1925).

[4] K. Sandera, Listy Cukrovar. 48, 312-5 (1930; J. Peller, Listy Cukrovar. 49, 76-80 (1930); O. Spengler, F. Toedt, and J. Wigand, Z. Ver. deut. ZuckerInd. 82, 789-816 (1932); 83, 822-32 (1933); 84, 93-111 (1934), 84, 443-9 (1934); 84, 789-805 (1934); P. Honig and W. F. Alewijn, Arch. Suikerind. III, Mededeel. Proefsta. Java-Suikerind. (1932); 1811-224 A. Courriere, Betterave 42, 17–19 (1932); 44, 5 (1934).

[5] K. Sandera, Listy Cukrovar. 48, 312-5 (1930).

[6] V. Netuka, Listy Cukrovar. 46, 564–6 (1928).

[7] K. Sandera, Z. Zuckerind Čechoslovak. Rep. 53, 378-82 (1929).

[8] A. R. Nees, Ind. Eng. Chem. (Ind. Ed.) 19, 225–6 (1927).

[9] J. H. Zisch, Facts About Sugar 25, 741-6 (1930).

[10] V.Dejmek and F. Stern, Listy Cukrovar. 53, 353–6 (1935); Z. Zuckerind. Čechoslovak. Rep. 59, 417-21 (1935).

[11] G. Jones and R. C. Josephs, J. Am. Chem. Soc. 50, 1049-1092 (1938). [12] W. A. Taylor and S. F. Acree, Science, 44, 576-578 (1916).

[13] F. Kohlrausch and L. Holborn, Leitvermögen der Electrolyte (B. G. Teubner, Leipzig, 1898).

[14] K. Sandera, Chimie & industrie, 651, Special No. May, 1927.

[15] W. A. Taylor and S. F. Acree, J. Am. Chem. Soc. 38, 2415-2430 (1916). [16] W. A. Taylor and S. F. Acree, J. Am. Chem. Soc., 38, 2396-2403 (1916). [17] G. Jones and S. M. Christian, J. Am. Chem. Soc. 57, 272-280 (1935). [18] G. Jones and G. M. Bollinger, J. Am. Chem. Soc. 53, 411-451 (1931). [19] M. Randall and G. N. Scott, J. Am. Chem. Soc. 49, 639 (1927); F. A. Smith, J. Am. Chem. Soc. 49, 2167-2171 (1927).

[20] H. C. Parker, J. Am. Chem. Soc. 45, 1366, 2017 (1927).

[21] F. W. Zerban and L. Sattler, Facts about Sugar 21, 1158-1162 (1926). [22] J. Reilly and W. N. Rae, Physico-Chemical Methods, p. 707 (D. Van Nostrand Co., Inc., New York, N. Y.) 1932.

[23] G. Jones and D. M. Bollinger, J. Am. Chem. Soc. 57, 280-284 (1935) [24] C. W. Davies, The Conductivity of Solutions, 2d. ed., p. 56 (John Wiley & Sons, Inc., New York, N. Y., 1933).

[25] F. Kohlrausch and Heydweiller; Z. Physik. Chem. 14, 326 (1894).

[26] Edna H. Fawcett and S. F. Acree, J. Bact. 17, 163–204 (1929).

[27] J. Kendall, J. Am. Chem. Soc. 38, 2460-2466 (1916).

[28] F. W. Zerban and L. Sattler, Facts about Sugar 21, 1158-1159, 1162-6 (1926).

[29] G. Jones and M. J. Prendergast, J. Am. Chem. Soc. 59, 731-736 (1937). [30] C. W. Davies, J. Chem. Soc. 59, 432-436 (1937).

[31] Theodore Shedlovsky, J. Am. Chem. Soc. 54, 1424 (1932).

[32] C. R. Johnson and G. A. Hulett, J. Am Chem Soc. 57, 258 (1935).

[33] Theodore Shedlovsky, J. Am. Chem. Soc. 54, 1410 (1932).

[34] W. A. Taylor and S. F. Acree, J. Am. Chem. Soc. 38, 2403–2415 (1916). [35] G. Jones and B. C. Bradshaw, J. Am. Chem. Soc. 55, 1791 (1933). [36] H. Thiene, Glas, p. 130, (G. Fischner, Jena, Germany, 1931.) [37] C. G. Peters and C. S. Cragoe, Sci. Pap. BS 16, 449 (1920) S393. [38] G. Jones and D. M. Bollinger, J. Am. Chem. Soc. 57, 280–284 (1935). [39] F. W. Zerban and L. Sattler, Ind. Eng. Chem., Anal. Ed. 3, 43 (1931). [40] H. Lunden, Centr. Zuckerind, 33, 204–5 (1925); B. Lazar, Listy Cukrovar. 49, 480-5 (1931); M. Sanderova and K. Sandera, Listy Cukrovar. 53, 389-6 (1935); F. Majer, Listy Cukrovar, 54, 341-5 (1936).

[41] F. W. Zerban and L. Sattler, Facts about Sugar 22, 990-994 (1927). [42] F. W. Zerban and L. Sattler, Ind. Eng. Chem., Anal. Ed. 2, 32–35 (1930). [43] F. W. Zerban and L. Sattler, Ind. Eng. Chem., Anal. Ed. 3, 38-40 (1931). [44] S. F. Acree, Edward Bennett, G. H. Gray, and Herald Goldberg, J. Phys. Chem. 42, 871-896 (1938).

XVIII. MEASUREMENT OF HYDROGEN-ION

CONCENTRATION

1. INTRODUCTION

Bates and Associates [1] in 1920 reported results giving the hydrogenion concentration of aqueous solutions of a number of commercial sugars. In 1922 Brewster and Raines [2] published an account of the use of hydrogen-ion methods for reaction-control in cane juices in the experimental manufacture of sugar. Balch and Paine [28] in 1925 described a scheme for the automatic recording of hydrogen-ion concentration in lime-treated cane juice. These appear to be the earliest recorded instances of the application of such methods (now regarded as indispensable) to commercial sugar products and their manufacture. These were followed shortly by the development of apparatus for the automatic dosage of cane or beet juices with lime. or carbon dioxide to a desired end point controlled by electrodes reacting to the pH of the liquid, the resulting pH being simultaneously recorded.

Prior to the advent of hydrogen-ion methods, the adjustment of reaction in sugar juices was based upon the titration of a quantity with standard acid or alkali to an indicator end point. This gave a measure of the quantity of acid or alkali present. Hydrogen-ion methods, on the other hand, give information regarding the intensity of reaction of acid or alkali due to the concentration of hydrogen (or hydroxyl) ions, which influence the rate of inversion of sucrose, and the clarification, filtration, and decolorization of juices, as is well known. The nomenclature of hydrogen-ion concentration is based upon the normal weight (1.008 g) of ionized hydrogen in a liter of solution. This may be expressed fractionally as 1/1 N. A solution containing 0.01008 g of hydrogen ions per liter would be 1/100 N, and so on. If these fractions are expressed as powers of 10, we have 1/1-10°,1/10=10', 1/100=102, . . . 1/1,000,000=10-6, and So on. The negative exponent, which may or may not be a whole number, is the logarithm of the reciprocal of the hydrogen-ion concentration, log 1/[H+] (the brackets indicate normality). To this has been assigned the symbol pH. The numerical value of pH is sometimes called the hydrogen-ion exponent. Instead of using fractions, we therefore write pH=1, pH=6, pH=8.4, etc. Also one may say that the pH value of a solution is the logarithm of the number of liters that contain 1 gramion of hydrogen. Thus in pure water, at the ordinary temperature, it requires 10 million liters to yield 1.008 g of ionized hydrogen. The log of 10,000,000 is 7. Therefore the pH of pure water is 7.

In all neutral aqueous solutions, as well as in pure water, pH+ POH=14 at the ordinary temperature, and the products of dissociation, H and OH- ions, are present in equal amounts, so that pH=pOH=7. The expression, pOH, which might be taken for the purpose of expressing alkalinity, ordinarily is not used, since, as may be seen by reference to the two equations, values of pH below 7 indicate an excess of H+ over OH- ions, and the solution is said to be acid. Similarly, when pH values are above 7, the OH- ions are in excess and the solution is said to be alkaline.

The so-called strength of a pure acid or alkali in solution depends upon the degree of ionization and the hydrogen-ion concentration of such solutions may be calculated from the ionization constant. If a strong acid or alkali be added little by little to water, the hydrogen-ion concentration changes enormously with each addition. If, however, a soluble salt of the acid or base be present, the ionization is depressed and upon the addition of the one or the other the change in hydrogenion concentration is gradual and may be controlled within certain limits. This resistance to change in pH due to the presence of a salt is called buffer action. Buffer action also results from the presence of salts of bases with weak acids as encountered in most plant juices and accounts for the very gradual increase in hydrogen-ion concentration of cane juice when the latter is treated with sulfur dioxide. Advantage is taken of the buffer action of certain salts in the preparation of standard solutions for colorimetric pH methods, as described later.

Methods for the measurement of hydrogen-ion concentration may be divided into two categories (a) potentiometric methods, whereby is measured, under proper conditions, the potential of a concentration cell, of which the unknown solution is a part, and (b) colorimetric methods depending upon the use of indicators, the color or shade of color of which changes with variation in hydrogen-ion concentration. It is advantageous to have both potentiometric and colorimetric methods available.

2. POTENTIOMETRIC METHODS

(a) POTENTIOMETERS

The electromotive force (emf) of a concentration cell is measured by balancing against it a measurable potential from an external source. At balance no current passes through the cell, as indicated by the null-point instrument. The apparatus employed consists of a potentiometer; a galvanometer as indicating instrument; an outside source of emf, such as that furnished by dry cells or a storage cell, the value of which is known by reference to a standard cell, the lastnamed being usually a Weston cadmium-mercury cell.

Potentiometers are available in several forms, some of which are highly accurate and used in research and for calibrating other instruments. Others intended for industrial laboratory and plant use, although having sufficient accuracy, are portable, with galvanometer, standard cell, and measuring device in a single housing, or with all accessories, including glass and reference electrodes, amplifying tubes, and voltage supply housed within a small space. The choice of apparatus is dictated by the accuracy required and by the nature of the material in which measurements are to be made.