272 INTERMEDIATES IN REDUCTION OF CO2 CHAP. 10 



dissociation constant of about 6.2 X 10-» (pK 4.21) (c/. Ball 1937). (This is the 

 extrapolated dissociation constant at the ionic strength m = 0; from measurements of 

 the oxidation-reduction potential — see below — a value of 9.1 X 10~^ at fi = 0.1 was 

 obtained.) Consequently, the acid must be present in the tissues almost entirely as an 

 anion (or as a metalloorganic complex). The large dissociation constant seems at first 

 to contradict the accepted formula, since the latter shows no carboxyl group. However, 

 the group, — COH^COH — CO — , apparently has an acid character similar to that 

 of a carboxyl group (cf. Ball 1937). Ascorbic acid has an L value of 0.833, that is, it is 

 a partially oxidized sugar. Its main tendency is to oxidize itself further, by loosing two 

 or even four hydrogen atoms. In a certain pH range, this loss is reversible, particularly 

 as far as the first step is concerned. It transforms ascorbic acid into dehydroascorbic 

 acid (CeHsOe, L = 0.75, cf. Formula 10.11). 



Many attempts have been made to determine the oxidation-reduction potential of 

 the ascorbic acid-dehydroascorbic acid system, i. e. by Georgescu (1932), Wurmser and 

 Loureiro (1933), Green (1933), Borsook and Keighley (1933), Fruton (1934), Borsook, 

 Davenport, Jeffries, and Warner (1937), and Ball (1937). According to Ball (1937), 

 the system is electrochemically sluggish, so that "electrode catalysts" (for example, 

 thionine or methylene blue) must be added to accelerate the establishment of the elec- 

 trode equihbrium. Furthermore, according to Borsook and Keighley (1933) and Ball 

 (1937), the oxidant (dehydroascorbic acid) is unstable in neutral solution (pH > 5.75). 

 Therefore, reliable potentials can be obtained only in the acid range. (Above pH 6 the 

 "apparent" normal potential becomes more positive with time because the oxidant 

 gradually disappears from the system). 



Taking these comphcations into account, Ball was able to calculate the normal 

 potentials of the ascorbic acid-dehydroascorbic acid system between pH 1 and pH 8.6, 

 and obtained (for 30° C.) the values Eo = - 0.329 volt for pH 1, and Eo' = - 0.057 

 for pH 7. 



Neutral solutions of ascorbic acid reduce thionine (Eo' = — 0.06 v.), cresyl blue 

 (^0' = - 0.047 V.) and (slowly) methylene blue (Eo' = - 0.011 v.), but not Nile blue 

 or phenosafranine (Eo = + 0.252 v.). They can be titrated with 2,6-dichlorophenol- 

 indophenol (£^0' = — 0.20 v.). The oxidation by methylene blue can be accelerated 

 by fight, according to Mentzer and Vialard-Goudon (1937). Ascorbic acid also reduces 

 silver nitrate (cf. page 270), cupric ions (e. g., Fehhng solution), mercuric ions, ferric 

 ions, nitrites, quinones, indophenols, flavones, etc. (cf., for example, King 1939). The 

 reduction of dehydroascorbic acid to ascorbic acid can be achieved by hydrogen sulfide, 

 at pH 3-4. 



According to Kelfie and Zilva (1938) and Arcus and Zilva (1940), ascorbic acid in 

 solution is oxidized to dehydroascorbic acid by ultraviolet fight in the absence of oxygen. 

 The reaction must be either a reduction of water or a dismutation. Pure ascorbic acid is 

 not oxidized directly by oxygen but numerous substances catalyze this reaction (e. g., 

 Cw^"^ ions, in both the free state and in organic complexes; cf. King 1939). 



The increased production of ascorbic acid in the presence of glucose, 

 as well as its formation in seedlings before the beginning of photo- 

 synthesis {cf. Rubin and Strachitzky 1936), indicate that it is formed by 

 oxidation of sugars. Therefore, the increase in ascorbic acid concentra- 

 tion following intense assimilation (reported by Giroud 1938, Reid 1938, 

 and Moldtmann 1939) does not necessarily mean that it is an intermediate 

 of photosynthesis (although this possibility, first suggested by von Euler 

 and Klussmann in 1933, cannot be excluded). Photosynthesis may 



