VLa-6 MANUAL OF METHODS FOR PURE CULTURE STUDY 



In solutions more alkaline than about pH 9, the 015 glass electrode 

 responds also to cations other than H ions, the potential being in- 

 fluenced by the activity and kind of such cations. Sodium and 

 lithium ions produce the most marked effects, potassium and bivalent 

 cations smaller effects. When working under these conditions, it is 

 advisable to standardize the electrode with known buffer solutions of 

 about the same composition and of pH closely above and below the 

 pH of the sample being tested. 



The standardization for linearity of response from pH 1 to 9 is a 

 necessary check on the operation of the glass electrode, since its re- 

 sults are comparative, not absolute. The slope, -AEh/ApH, 

 should be not merely constant at any temperature but also equal or 

 closely equal to 0.000,198,322 T (the values for this constant are 

 shown under A on p. 4). Obviously, a "pH-meter" with its pH 

 scale adjusted to the theoretical slope for a given temperature cannot 

 give correct readings at all points from pH 1 to 9 if its glass electrode 

 follows a significantly different slope at the same temperature. For 

 a brief discussion of the effects of temperature, see Clark (1948). 



Cleaning of the glass surface, by immersion in a hot mixture of concentrated nitric 

 and sulfuric acids followed by soaking in water, may restore a sluggish or erratic 

 electrode to normal functioning. A somewhat drastic procedure that may be effective 

 is to dip the glass electrode for a second or two in dilute HF or in a 20% solution of 

 ammonium bifluoride and then to wash it thoroughly in water. If the electrode still 

 behaves erratically, it should be discarded. For such an emergency, it is highly 

 advisable to have available a reserve electrode. This may obviate any mistaken 

 tendency to carry on with an electrode of doubtful reliability. 



The instructions accompanying the various glass-electrode "pH-meters" now on 

 the market are usually sufficient to aid the user in tracing out sources of trouble and 

 error in operation. A major source of trouble is electrical leakage due to accumulation 

 of films of moisture at critical parts of the circuit; and perhaps the most frequent sites 

 of such accumulation are the electrode support and lead, both of which are apt to be 

 spattered with water or salt solution during careless manipulation. 



The glass electrodes now available are fairly rugged and easily adaptable to use 

 under a variety of conditions and on difiFerent types of biological material (e.g., liquid 

 and "solid" culture media). Measurements with an accuracy of 0.05 pH may be 

 made rapidly in poorly buffered, colored, or turbid solutions, and in blood or serum. 

 The monograph by Dole (1941) discusses many of its uses. 



THE COLORIMETRIC METHOD 



The colorimetric method of measuring pH makes use of acid-base 

 indicators, which, within certain limits, vary in color with the pH of 

 the solution. Such indicators are compounds capable of existing in 

 solution as conjugate proton (H-ion) donor and proton acceptor, with 

 one of the conjugate pair differing in color from the other. The re- 

 lation of these two forms to pH is defined by the equation 



[proton acceptor] 



pH=pK'+log (5) 



[proton donor] 



in which brackets represent concentrations, and pK' (= - log K') 

 is called the apparent ionization exponent of the indicator's proton 

 donor-acceptor system. Simple calculations, using, for example, 

 0.8, 0.5 and 0.3 as values for the ratio [proton acceptor]/ [proton donorl 

 at each of the pK' values 3, 6, and 9, will show that indicators with 

 different pK' values cover different ranges of pH. (See Fig. 1). For 



