ACIDOSIS 53 



to have caused an abnormal state when it has either increased the hydrion 

 concentration of the blood or lowered its alkali reserve below the extreme 

 normal limit. 



The results of acid retention, both clinically and on the physico- 

 chemical state of the blood, are qviite different according to whether the 

 retained acid is carbonic or other than carbonic. Retention of carbonic 

 acid causes increase in the hydrion concentration of the blood (or decrease 

 in the pH). 2 It also raises the bicarbonate concentration by causing a 

 shift to this form of some of the alkali which at normal pH is bound in 

 other buffer salts. 



Retention of acids other than carbonic, on the contrary, lowers the bi- 

 carbonate, a chemical equivalent of which is decomposed by the retained 

 acid. Unless nearly all of the bicarbonate is thus decomposed, however, 

 the pH is not changed from the normal, so long as there is no hindrance, 

 mechanical or nervous, to the excretion of CO 2 by the lungs. The acidosis 

 caused by retention of non-volatile acids, without fall in pH except as 

 a terminal phenomenon, is the form commonly met in metabolic diseases, 

 and until recently was the only form that had been observed clinically. 



The bicarbonate is only one of a number of buffers in the blood, and 

 in the body, exclusive of the blood, which are of importance in regulating 

 the neutrality. The bicarbonate is the only important buffer, however, 

 which can be used to neutralize acids without any change whatever in the 

 blood pH. Furthermore the interreactions among the blood buffers are 

 such that, as will be shown later, the bicarbonate, taken together with the 

 pH, indicates the state of the entire system. We shall discuss below the 



"The term pH was introduced by Soerensen (1912) as a convenient symbol for ex- 

 pressing minute hydrogen ion concentrations without resort to decimals, and for plot- 

 ting changes over great ranges of hydrogen ion concentration. It is the negative 



N 

 logarithm of the hydrogen ion concentration, e.g., for -yrwr hydrogen ion concentration, 



CTT + = 1 x 10- a . The logarithm is therefore 2, the pH merely 2. The pH of water is 



approximately 7, i.e., the hydrogen ion concentration is 0.000,000,1 normal. The pH 

 of solutions' more acid than water is "less than 7, that of solutions more alkaline is 

 greater than 7, and each change of 1 in pH means a ten-fold change in hydrogen ion 

 concentration, e.g., pH denotes ten-fold the hydrogen ion concentration of water, 

 pH 5 denotes one hundred-fold. On the other side, pH 8 denotes 0.1, pH 9 denotes 

 0.01 the hydrogen ion concentration of water. For pH 7.35, that of blood, the hydrogen 

 ion concentration is calculated as follows: 7.35=: 8 + .65, 0.65 = log of 4.5 .'. Hy- 

 drogen ion concentration = 4.5 x 10- 8 . 



The hydroxyl ion concentration at pH 7 is equal to that of the hydrogen ions, the 

 dissociation constant of water being at ordinary temperatures approximately 10- 14 . 



C H + XC OH' = 10J4 - 



Consequently the log of the OH' concentration is calculated from the pH by subtracting 

 14 from the pH, e.g., pH=12; log OH' concentration = 12 14 = 2. OH' concen- 

 tration =0.01 N. 



It follows that anv OH' concentration may be expressed in terms of H + concentra- 

 tion or vice versa, values in terms of pH affording a convenient means of covering 

 the entire field of reaction, both acid and alkaline. 



