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

3. k2 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 1925X10-6 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) $393.

[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 carliest 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, 1/100=10-2, . . . 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.

A diagram of the Leeds & Northrup type K potentiometer is shown in figure 53, which, with the description that follows, illustrates the potentiometric principle and the operation of the instrument. In figure 54 the instrument is shown in use with a hydrogen gas-calomel cell. The standard cell, galvanometer, and dry cells also are shown. In the diagram the portions of the bridge by which measurements are made consist of fifteen 5-ohm coils in series in the circuit AD (contact made with M), and in series with them the extended wire DB, the resistance of which is also 5 ohms. The scale of DB, shown in the photograph, reads from 0 to 1,100. Contact with the extended wire is made by the moving contact, M'. Current from the battery, W, flows through these resistances and may be made exactly 0.02 ampere by means of the regulating rheostat, R. This is done by setting the double-throw switch to "std. cell", which connects the standard cell in series with the galvanometer, G, and the tapping keys,

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Courtesy of Leeds & Northrup Co.

FIGURE 53.-Wiring diagram of type K potentiometer

R1, R2, and R3, and with the point .5 on AD, to which one point on the double-throw switch is wired. Between A and O there is a series of resistances with a sliding contact, T. The resistance between the points .5 and A is exactly that which corresponds with the emf of 1 volt, and between 1.0166 and .5 a sufficient resistance is added to make the resistance between these points correspond exactly with an emf of 1.0166 volts. The small circular slide wire connected at 1.0166 makes contact with the standard cell through the contact, T. The resistance value of this slide wire is such that a practically continuous variation can be obtained in the standard-cell circuit voltages from 1.0166 to 1.0194, a range corresponding with the variations in different standard cadmium cells. To adjust the current to 0.02 ampere, with the switch thrown to the "std. cell" position, the contact, T, is set to correspond with the standard cell voltage and rheostat, R, is regulated until the galvanometer shows no deflection. The unknown emf is

now measured by throwing the switch to the emf position and adjusting the resistances with M and M' by touching the contact keys, R2, R1, and R, until there is again no galvanometer deflection. After this measurement is made the working current may again be checked against the standard cell, as described.

In figure 55 is shown a Leeds & Northrup potentiometer-electrometer with range from 0 to 1.100 volts. This portable instrument contains a stage of amplification and is suitable for pH measurements

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FIGURE 54.-Leeds & Northrup type K potentiometer with hydrogen-gas electrode,

calomel electrode, etc.

with all types of electrodes, including glass. Other types of potentiometers are illustrated in figure 59, A and B.

(b) CALOMEL ELECTRODES

Various forms of calomel electrode vessels are shown in figure 56. The bulb at the bottom of the cell (C, D, E) is filled with a layer of pure mercury. On this rests a layer of pure calomel mixed with mercury, and the filling of the cell is completed with a solution of potassium chloride having a definite concentration and saturated with calomel. One of three concentrations of potassium chloride is customarily used, either 0.1 M, 1.0 M, or saturated, and in reference to these concentrations the terms "tenth-normal", "normal", or "saturated" calomel electrodes are used as abbreviations. By means of side-tubes the cells communicate with a reservoir of saturated potassium chloride which by turning the stopcock at the top of the cell may be allowed to fill the side arm and form the necessary salt bridge between the calomel

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