deduction and Complexing of 

langanese by Gallic Acids

31 JOHN D. HEM

:HEMISTRY OF MANGANESE IN NATURAL WATER

GEOLOGICAL SURVEY WATER-SUPPLY PAPER 1667-D

Chemical reactions with organic acids 

vhich influence the behavior of 

nanganese in natural water

KITED STATES GOVERNMENT PRINTING OFFICE, WASHINGTON : 1965
UNITED STATES DEPARTMENT OF THE INTERIOR 

STEWART L. UDALL, Secretary

GEOLOGICAL SURVEY 

Thomas B. Nolan, Director

For sale by the Superintendent of Documents, U.S. Government Printing Office 

Washington, D.C. 20402 - Price 20 cents (paper cover)
CONTENTS

Page

j' bstract_____________._____________________._-_.-----_--__.____- Dl

Introduction ______________________________________________________ 1

Previous work _________________________________________________ 1

Behavior of manganese in lakes__________--__-_-__-_-----_---_-_-___ 2

Organic materials in natural water____________________________-_____- 4

Properties of tannic acid__________________________________________ 5

Digallic acid complex with manganese______________________________ 6

Determination of stability constant___-___-_-_-__-_---_----_____- 8

Properties of gallic acid____________________________________________ 10

Complexes of Mn+2 with other organic ligands_______________________ 11

 Complexes with higher oxidation states of manganese __________________ 12

Tolution of manganese dioxide by digallic acid_______________________ 13

Calculation of amount of manganese complexed___________________ 20

Solvent-extraction experiments._________________________________ 22

Stability of complex toward oxidation____-_-_____--_-__-----_-__- 22

! 'olution of manganese dioxide by gallic acid __________________________ 23

ffects of gallic acid compared with those of digallic acid_______________ 25

"mportance of reducing and complexing in natural systems.____________ 25

.literature cited.__________________________________________________ 27

ILLUSTRATION

Page 

"'IGTJRE 1. Solution of manganese by tannic acid at (A} pH 5.5 and (B) pH

8.5..________________________________________________ D15

TABLES

Pag» 

TABLE 1. Computation of stability constant for manganese digallate

complex ___ ___ _ _____________________________ __ D9

2. Manganese concentration and pH of solutions in contact with

manganese dioxide. Tannic acid absent. ________________ 14

3-6. Manganese concentration, pH, and Eh of 

3. 10-ppm tannic acid solution in contact with manganese

4. 100-ppm tannic acid solution in contact with manganese 

dioxide____-___ ----------------------------------

5. 1,000-ppm tannic acid solution in contact with manganese dioxide _ ___________________________________

6. Solutions of gallic acid in contact with manganese dioxiden r

16

16

17

24

CHEMISTRY OF MANGANESE IN NATURAL WATER

REDUCTION AND COMPLEXING OF MANGANESE BY 

GALLIC ACIDS

By JOHN D. HEM

ABSTRACT

Tannic (digallic) and gallic acids form 1:1 complexes with Mn+2 in dilute solution. The stability constant of the digallic complex was determined to be 103 -9 

by spectrophotometric studies. Water solutions of tannic and gallic acid can 

bring 100 parts per million or more of manganese into solution from solid manganese dioxide reagent. The reducing and complexing reactions are most rapid 

at pH's below 5. The manganese brought into solution was oxidized ve~y slowly 

at pH's as high as 9.8.

INTRODUCTION

The biochemical factors that are involved in the chemical behavior 

of manganese in natural environments have been touched upon in 

the earlier papers in this series. A further study of some chemical 

reactions that may influence manganese solution and deposition is 

described here. The study concerns the behavior of oxidized and 

reduced forms of manganese in the presence of the organic compounds 

tannic acid and gallic acid. The rates of reduction of tetravalent 

manganese by the organic acids, and the form and stability of the 

soluble manganese compounds formed in the reduction process, were 

investigated. The results provide some indications as to the principles which govern the circulation of manganese in natural environments at or near the land surface, where organic material is plentiful.

PREVIOUS WORK

Preliminary studies of the behavior of divalent manganese in dilute 

solutions of tannic acid were described by the writer in an earlier 

report (Hem, 1963, p. A51-A52), but no definite conclusions were 

drawn. A green color was observed in solutions containing tannic 

acid and divalent manganese, and a similar color was noted when the

Di
D2 CHEMISTRY OF MANGANESE IN NATURAL WATER

pH of the tannic acid solutions was increased in the absence of manganese. Tannic acid noticeably retarded the rate of oxidation of 

divalent manganese when the alkaline solutions wer? exposed to the 

air. In some environments, such as bogs, that arc rich in organic 

matter, the effects of organic solutes are unquestionably important. 

In this paper reduction and complexing effects have been studied in 

detail so that their importance could be intelligently evaluated. 

Some preliminary results of experiments with tannic-acid reduction 

of manganese dioxide have been reported by Rawson (1963).

BEHAVIOR OF MANGANESE IN LJ KES

One of the more common water-quality problems encountered in 

the water-supply industry is the appearance of manganese in solution 

in the water withdrawn from a supply reservoir. Tl e literature contains many references to this type of problem, with descriptions of 

conditions found and methods used for coping with the problem. The 

causes, however, have not always been clearly understood or satisfactorily explained.

Vegetation requires manganese in its life processes, and some of 

the aquatic species of plants seem to accumulate the element (Oborn, 

1964). Hutchinson (1957, p. 805) states that the mr.nganese content 

of lake water is directly related to the productivity of the lake. Lakes 

and reservoirs support a considerable volume of aquatic vegetation, 

ranging in size from diatoms to large water weeds rooted in the lake 

bottom, and during the growing season as much af 10 tons of dry 

cellulose and related material could be produced per acre of water 

surface (Lieth, 1963). This high production rate is probably more 

typical of sewage-oxidation ponds or peat-bog or swamp environments than of a clear lake, but any water impoundment can be expected to produce a considerable mass of organic nratter each year. 

In his studies of productivity in several lakes in Denmark, Nygaard 

(1953) found a production rate of organic matter of 0.62 g per day per 

square meter in relatively clear water 28 meters deep. This rate is 

equivalent to more than a third of a ton per acre in a growing season 

of 150 days and is probably lower than the usual rate for a storage 

reservoir. In large part, the organic matter produced in each growing 

season decays, and a residue of low-solubility mater al settles to the 

lake bottom.

Organic matter of aquatic vegetation contains fr-om a few hundredths to a few tenths of a percent of manganese (Oborn, 1964). 

The amount of manganese present depends partly on its availability. 

Bottom-rooted plants can extract manganese from the soil below the 

water surface, but free-floating forms can obtain it only from the 

water in which they grow.
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D3

The quantitative importance of biological productivity as a mechanism for trapping manganese in a lake can be estimated in several 

ways. The amount of manganese in solution brought into a reservoir 

during the growing season probably is normally almost all removed 

biologically. The median manganese concentration of North American 

rivers of major size is reported by Durum and Haffty (19P3) to be 

20 micrograms per liter. Thus, in 1 acre-foot of water, one might 

expect to find about 25 g of manganese. The median concentration 

of 20 micrograms per liter is equivalent to 0.02 ppm (parts per million) 

and is a higher concentration than might be expected from the rather 

sparse data for manganese content in the analytical records published 

by the U.S. Geological Survey in annual reports on quality of surface 

water of the United States. In keeping with longstanding custom, 

the dissolved constituents in river water are ordinarily determined on 

filtered or clarified samples. Durum and Haffty used samples which 

had been filtered and acidified, but the filtration did not remove all 

the particulate matter. Hence, their results may include some of the 

participate manganese that was present. Hutchinson (1957) indicates 

that particulate manganese may be utilized by plants and in any 

event is likely to be precipitated in a standing water body such as a 

lake or reservoir.

A production rate of one-quarter ton of dry vegetation per acre of 

lake surface per year, of material containing 0.10 percent mr.nganese, 

would represent a rather modest productivity and would require 250 

g of manganese per acre of lake surface. The amount of manganese 

available from lake water having 20 micrograms of manganese pe 

liter, or 25 g of dissolved manganese per acre-foot, is enough to supply 

the biochemical demand for the whole year if the photosynthetic 

process is active to a depth of 10 feet. An annual inflow of water 

sufficient to renew the manganese supply would be required to maintain the process. Bottom-rooted aquatic species growing in shallow 

water would extract some of their manganese from the soil of the lake 

bottom. When the growing season ends, the crop of vegetation dies 

back and the debris settles toward the bottom of the lake. This 

debris carries with it at least a part of the manganese originally 

incorporated in the vegetation and forms an organic-rich precipitate.

Biological processes, therefore, tend to remove manganese from 

water and transfer it to a lake-bottom precipitate in significant 

amounts. The normal circulation of aerated water through the lake 

maintains oxidizing conditions and depresses manganese solubility. 

During periods of thermal stratification, howevei, the deeper water 

may be depleted in oxygen through biochemical processes to a degree 

that induces reduction and re-solution of manganese in wr.tei near 

the bottom muds that contain the manganese placed there earlier
D4 CHEMISTRY OF MANGANESE IN NATURAL WATER

through the decay of vegetation. The volume of water, so affected, 

will generally be small compared with the volume from which this 

manganese was originally extracted. Thus, the effect of a 10-fold 

or 100-fold concentration could readily be achieved, leading to a 

manganese concentration in the bottom water of 0.2 ppm or 2.0 ppm 

instead of 20 micrograms per liter. Concentrations higher than a few 

tenths parts per million lead to problems if introduced into water 

distribution systems. Intakes may obtain such water if they are 

oriented in such a way as to intercept bottom water, or if the bottom 

water migrates to other levels.

Although the biochemical processes of manganese removal and occasional re-solution can probably account for the usual occurrences of 

manganese in water of storage reservoirs, the actual processes in the 

bottom sediments may involve some additional factors related to 

organic complexing and reduction. In a few areas, such as the piedmont region of Georgia studied by Ingols and Wilroy (1963), the manganese content of the material making up the original bottom of the 

reservoir may be high enough to be a significant source of manganese 

for solution.

ORGANIC MATERIALS IN NATURAL WATER

The nature of naturally occurring organic solutes in colored river 

and lake water has never been completely established, even though 

such materials have been extensively studied. Lamer and Goerlitz 

(1963), however, were able to identify 12 carboxylic acids in unpolluted 

surface water from streams in the States of California ard Washington. 

Substances identified included acetic, propionic, butyric, valeric, 

caproic, lactic, malonic, maleic, succinic, adipic, fum^ric, and oxalic 

acids. The concentrations observed were from 4 to 592 parts per 

billion. Other carboxylic acids of higher molecular weight were also 

isolated but could not be specifically identified by gas chromatography 

because they were relatively stable and nonvolatile. P'lch acids have 

commonly been referred to in the literature as "mimic" or "fulvic," or 

are given other names, which imply a definite chemical formula or 

structural pattern. Actually these names have been us3d with various 

connotations by many different writers to refer to convex mixtures of 

high molecular weight organic acids and polymers which impart a 

brown or yellow color to water and can be extracted by certain empirical procedures. The color and probably the structure of these 

materials are partly a function of pH and it has not been agreed 

whether they occur in true solution or as a colloidal suspension. The 

molecules of these substances may actually approach colloidal size.

Shapiro (1964) conducted extensive studies of organic coloring 

matter which he extracted from lake water which 1 e refers to as
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D5

"yellow organic acids" and has determined to be a mixture of aliphatic 

polyhydroxy carboxylic acids.

Black and Christman (1963a, b) suggested that the organic material 

responsible for color observed in many natural river waters consists of 

hydroxycarboxylic acids which, have some degree of unsaturation and 

aromaticity. Their material also contained between 1 and 2 percent 

nitrogen. Lamar and Goerlitz (1963) reported that the infrared spectrum of colored material obtained from one of their waters resembled 

that of metadigallic acid, and the studies of Black and Cl ristman 

(1963b) also indicate that such a compound can be expected ir colored 

natural water. Therefore, it is possible to obtain some useful indications of the behavior of manganese in natural waters by studying the 

reactions between manganese and digallic (tannic) acid. Tarnic acid 

is readily available in relatively pure form. Another series of experiments was performed with gallic acid, C6H2 (OH) 3COOH, in the expectation that the simpler compound would be more stable and easier 

to work with.

The results of experiments with well-defined chemical compounds 

can be expressed in relatively exact terms, with respect to the nature 

of reactions and products and to the stability of complex ions which 

are produced. However, results of experiments with the natural 

organic materials extracted from water are usually less exact.

PROPERTIES OF TANNIC ACID

Tannic acid (digallic acid) has the chemical formula (HO) 3C6COO 

C6H2 (OH) 2COOH[, and can be represented structurally as

HOOK

OH

It is formed by hydrolysis from the more complex substances called 

tannins. Tannins are commonly present in vegetation, and therefore 

natural water, which usually has some contact with plant debris, commonly contains tannic acid or chemically similar material in solution.

The reagent-grade tannic acid used in this work was a pale yellowbrown powder, which dissolved readily in water. Solutions containing 1 mg of reagent per ml of water were pale yellow, and at this concentration had a pH near 4.0.

In earlier work (Hem, 1960, p. 78) the writer estimated the first 

dissociation constant for tannic acid to be about 10~6 . The curves 

obtained in potentiometric titrations of the tannic acid witlN strong 

base gave no well-defined inflection points but seemed to suggest that

769-002 65 2
D6 CHEMISTRY OF MANGANESE IN NATURAL WATER

the carboxyl hydrogen is mostly titrated before the pher olic hydrogens 

are detached from the molecule. The amount of base consumed, 

however, suggests that some influence from the phenolic hydrogens is 

probably present at pH 8 and that this influence increases as pH 

increases.

Redox titrations of tannic acid at various pH's u?mg potassium 

permanganate and platinum and saturated calomel e^ctrodes were 

performed to evaluate the reducing ability of the reagent. In these 

titrations the manganese was at first reduced from Mn+7 to Mn+2, but 

ultimately a black flocculent precipitate, or a dark-brown colloidal 

suspension, containing Mn+4 , was produced. These experiments indicate that solutions of digallic acid are definitely capal j e of reducing 

manganese to the Mn+2 state, and that their ability for reduction is 

increased by lowering the pH.

The chemical mechanisms involved in the reduction of permanganate 

by digallic acid were not definitely ascertained. The cdor of the solution suggests that the organic compound undergoes considerable 

alteration in the process. For the present study, the processes of 

oxidation and reduction which are of interest are those involving Mn+2 

and Mn+4, and therefore no detailed investigations of the effects of 

permanganate on digallic acid seemed necessary.

The amount of permanganate which reacted with 10 mg (3.1X10"6 

moles) of digallic acid without precipitating manganese oxide at pH 7.0 

was about 6X10~5 moles. The conversion of one MnO^ ion to Mn+2 

involves the gain of five electrons. Hence, the reduction equivalency 

of digallic acid in the conditions of this particular experiment was about 

10 times its molality. One mole of tannic acid, therefore, should be 

capable of reducing 5 moles of Mn+4 to Mn+2.

If, in being oxidized, the digallic acid molecule loses all 5 of its 

phenolic groups, and they are replaced with hydrogen, the net effect on 

the nominal charges on the carbon atoms in the 2 benzene rings is 

equivalent to the loss of 10 electrons. This mechanism would, therefore, be in accord with the results of the permanganate titration.

DIGALLIC ACID COMPLEX WITH MANGANESE

In earlier work (Hem, 1960), digallic acid appeared to form 

complexes with ferrous iron, and the possibility of similar complex 

formation with manganous ions was investigated. Aliquots were 

prepared containing 5.0 mg and 10.0 mg of digallic r.cid and small 

amounts of Mn+2 ranging from 0.02 to 1.0 mg. These aliquots were 

adjusted to pH 7.0 with 0.016-molar sodium hydroxide, diluted to 

25 ml, and the intensity of the green color which formed was measured 

with a spectrophotometer after the aliquots stood for 2 hours to
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D7

permit complete color development. Maximum absorbance occurred 

at 400 imt. A direct linear response was obtained for the lower 

concentrations of manganese (up to about 0.15 mg of manganese in 

25 ml of water). This response indicates a complex, whose behavior 

was in accord with Beers' law, in dilute solution. As the concentration 

of manganese was increased, however, a maximum color intensity 

was reached beyond which further increases in manganese content 

brought about no change.

The composition of the complex was determined from the spectrophotometric data using the procedure of Yoe and Jones (1944). In 

this procedure the absorbance is determined for various mole ratios 

of ligand to manganese. Straight parts of the resulting curve are 

extrapolated to an intersection which indicates the mole ratio of the 

predominant complex. A 1:1 mole ratio of manganese and digallic acid 

was indicated, at least for the concentrations of the reactants used 

here (about 10~4 to 10~8 molar) and at a pH near or slightly above 

neutrality.

The intensity of the green color was a function not only of manganese 

and digallic acid activities, but also of pH; the rate at which the color 

developed was also influenced by pH. The color was more intense 

at pH 8 than at pH 7 and developed more rapidly at the higher pH. 

After about 2 hours, at pH 7 and pH 8, the color stabilized and did 

not change significantly during the next 12-16 hours. Some greenish 

color was observed in solutions of digallic acid without manganese, 

when made alkaline with sodium hydroxide. The hue in these 

solutions, however, was somewhat different from that developed when 

manganese was present, and the intensity was generally much less 

when manganese was absent.

The color in these solutions is apparently associated with changes 

of the molecular configuration of the digallic acid. Such changes are 

probably brought about to some extent by neutralization with strong 

base which may displace some of the phenolic hydrogens. The fact 

that the color development is stronger in the presence of manganese 

and at higher pH's suggests that the two positive charges of the 

manganese ion are probably attracted both to the carboxyl and 

phenolic sites on the digallic acid molecule when the hydrogens in 

these positions are displaced.

If the digallate ligand is attached to the Mn+2 cation, both at the 

carboxyl and phenolic positions, the 1:1 complex would be uncharged. 

Perhaps this form would become more abundant as the pH is increased. 

However, in the absence of positive evidence as to the charge of the 

complex species, a linkage at only one position has been assumed.
D8 CHEMISTRY OF MANGANESE IN NATURAL WATER

DETERMINATION OF STABILITY CONSTANT

The stability constant of the manganous-digallate complex was 

determined from the spectrophotometric data using the following 

lines of reasoning :

The digallate ligand, represented by the symbol T~, is considered 

to be the form of tannic acid present in the solutions on removal of 

the carboxyl hydrogen, and possibly one or more of the phenolic 

hydrogens, as a result of raising the pH. Solutions used in the 

experiments were about 10~4 molar. In media having so low an 

ionic strength (near 10~4), the ion activity coefficients, represented 

by the symbol y, are near unity. The quantities in brackets are ion 

activities, the C terms are measured concentrations.

The association equilibrium

when written in mass-law form is

[Mn+2 ]|T-] M ' v '

and the constant KI is the stability constant for the 1:1 complex. 

The total amount of manganese present is either in the complexed or 

uncomplexed form

[Mn+2 ] . [M

The dissociation of digallic acid can be assumed to be very nearly 

complete at pH 8.0, because the first dissociation constant is near 

10~~6. Accordingly, the undissociated form can be ignored, and one 

may equate the total concentration of the ligand present to two forms

[MnT+] [T-] 

- ~

The optical density observed in the solutions was used as a measure 

of the amount of manganese complexed. The spectrophotometer 

readings were adjusted for the effects of absorbance by digallic acid by 

means of blanks containing no manganese, and the measurements were 

all made with the same length of light path. Under these conditions, 

for any given manganese concentration, a maximum optical density 

can be obtained by increasing the proportion of digallic acid to manganese in the solution. This maximum value represents the point at 

which essentially all the manganese has been complered. For solutions containing the same amount of manganese, but not enough
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D9

digallic acid to develop the maximum optical density, a lower opticaldensity value will be observed, and the ratio of total manganese to 

complexed manganese will follow the relation

CMII_maximum optical density 

observed optical density (4)

Experimental data for seven different laboratory samples sre given 

in table 1. They may be used to provide solutions for equations 1 to 4. 

In these solutions manganese concentrations ranged from 7.3X10" 6 to 

7.3 X10" 5 molar, and digallic acid concentrations ranged from 1.2X10"4 

to 6.2 X10~~4 molar. Maximum ionic strength in these solutions was 

4.8X10"4 . At this ionic strength, all activity coefficients rre close 

to 1.0. Values for 7Mnr and for 7r cannot be readily computed from 

the Debye-Hiickel equation; therefore, they were assumed to be 

unity. This assumption introduces an error of less than 5 percent into 

the values for KI. Values for 7Mn+2 were obtained from the DebyeHiickel equation using the concentrations of Mn+2 , Cl~, Na+, and 

T~ without complexing to calculate the ionic strength. The effect of 

complexing on the ionic strength could have been corrected for by 

further computation but this was not done because the spectrophotometric data could not be reproduced very closely. Values for K\ range 

from 1.4X104 to 5.5X103 . The mean of the seven values is 8.2X103 

or 103 - 9.

The uncertainties in the experimental data limit the accuracy of 

the constant. The pH had a considerable effect on the depth and 

stability of the color developed, and most of the range observed in K\ 

values in table 1 can probably be attributed to this factor. The true 

value of KI could be higher or lower than the value calculated here by 

a factor of 2.

TABLE 1. Computation of stability constant for manganese digallate complex

Sample

1. ._ ._

Z. ....................... _ .....

3. _. _ .....

4... ______ . ___ . ____ . ...

5.. .... .. ....... ......

6.. ...........................

7_ -

Total 

manganese 

SCuin (moles 

per liter)

3. 6X10-« 

3. 6X10-5 

7. 3X10-S 

7.3X10-6 

3.6X10-« 

7.3X10-* 

1. 5X10-»

Total 

digallic acid 

SOr (moles 

per liter)

2. 5X10-* 

6. 2X1CM 2. 5X10-4 

1.2X10-* 

1.2X10-' 

1. 2X10-4 

1.2X10-4

Absorbence

Maximum

0.44 

.44 

.88 

.46 

.79 

1.20 

.55

Observed

0.250 

.390 

.48 

.240 

.305 

.53 

.212

Ki

6.4X103 

1.4X10' 

6. 2X103 

9.6X103 

6.0X103 

9.6X103 

5. 5X103 

8. 2X103

The behavior of ferrous iron in digallic acid solution at pH 5.0 was 

considered earlier by the writer (Hem, 1960). For an assumed 1:1 

complex the stability constant was estimated to be about 1013 , based
D10 CHEMISTRY OP MANGANESE IN NATURAL WATER

on the interference of digallic acid with the formation of the pink 

complex of dipyridyl and ferrous iron.

Ferrous complexes with organic ligands are considerably stronger, 

as a rule, than the corresponding manganese complexes. However, 

the difference between the stability constants for manganese and 

ferrous complexes with digallic acid seems excessive. The composition of the ferrous complex was not determined, and it may be that 

the simple Fe!T+ structure that was assumed is incorrect.

The chemical behavior of the digallic acid reagent used in this 

series of experiments was not exemplary for exacting quantitative 

studies. The solutions of the acid were not stable in contact with 

air and slowly darkened in color on standing. Changes in the molecular configuration and complexing behavior, which might have resulted, were avoided by using freshly prepared solutiors. Complexing 

and neutralization reactions were relatively slow, and the green complex tended to change color after a day or so, probably because of 

oxidation. The amount of color development and the rate of complexing were strongly influenced by pH. Close agreement of results 

in successive sets of samples therefore was not always attained.

The strength of the digallic complex suggests tl x «,t in solutions 

containing about 10 ppm of digallic acid and a few tenths of a part 

per million or less of manganese, the manganese wo^ild be about 50 

percent complexed. Solutions of this kind may occ^r in rivers and 

lakes, although the nature of the organic matter in such water is 

probably not the same as the material studied. Natural coloring 

matter can be expected to have significant effects both as a reducing 

agent and as a complexing agent.

PROPERTIES OF GALLIC ACID

Experiments using gallic acid were compared with those using 

tannic acid. The reagent grade gallic acid used in tb^se experiments 

was a white powder, having the chemical formula (HO) 3C6H2COOHH2O. The solutions of this material were colorless at first but tended 

to darken with age. Both gallic and digallic acids seemed to serve 

as nutrients for fungi which grew luxuriantly in some of the solutions 

left in the open air. Biological growths, however, were prevented 

from forming in the solutions used in the actual experiments.

The first dissociation constant for gallic acid is reported to be 

4X10~B (Lange, 1956). The acidic properties are therefore stronger 

than those of digallic acid. Potentiometric titrations, using glass and 

saturated calomel electrodes, indicated a single well-defined inflection 

at pH 6.2. The phenolic hydrogen on the gallic acid r^olecule appears 

not to be active as a source of protons until the pH has increased
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACTDS Dll

considerably beyond 6.2, but phenolic hydrogens probably become 

available at about pH 8.

Gallic acid is an effective reducing agent and, like digallic acid, 

reacts with permanganate ions to yield Mn+2.

In the titration of gallic acid solutions with sodium hydroxide, a 

green color was observed which increased in intensity when the pH 

increased. This color probably resulted from the disruption or 

rearrangement of the gallic acid molecules. A very similar color was 

observed when gallic acid and manganese solutions were mixed and 

the pH adjusted. When manganese was present, however, the color 

appeared more rapidly and at a lower pH. The color developed in 

the presence of manganese was a more yellowish shade of green.

The nature of the complex formed between gallic acid and divalent 

manganese has been studied by Kachhawaha and Bhattacharya 

(1962). These investigators used higher concentrations of reactants 

and worked at a lower pH than in the work described here, but they 

did report a green color in some of their solutions. They reported a 

1:1 complex whose stability constant is 10*-16. This compares with 

the value 10 3 9 for the 1:1 complex of divalent manganese ard digallic 

acid determined in the present experiments.

Observations in the current study indicate that the gallic acid 

complex is weaker than the digallic complex, although perhaps not as 

much weaker as the difference between the two constants would 

indicate. A preliminary series of experiments in which the gallate 

complex was developed at pH 7 in the presence of different amounts 

of manganese and gallic acid gave results which indicated the 1:1 

association constant was between 102 and 103 . The value of Kachhawaha and Bhattacharya is in the lower part of this range. Hence, 

refined experimental work to determine the constant more closely 

was not undertaken.

COMPLEXES OF Mn+2 WITH OTHER ORGANIC LIGANDS

The compilation of Bjerrum, Schwarzenbach, and Sillen (1957) 

includes stability constants for a 1:1 complex of divalent manganese 

with several other organic acids which may occur in natural water. 

The acids and constants for the type reaction (where M is the metal 

ion and L is the organic ligand)

M+L=ML

are as follows: Oxalic acid, .K"=l02 - 55-103 - 96 ; malonic acid, K=IQ3 *9. 

It would appear that complexes of divalent manganese with tho organic 

ligands likely to occur in natural waters have similar stabilities.

\
D12 CHEMISTRY OF MANGANESE IN NATURAL WATER

Concentrations of a few tens of parts per million of organic acids would 

be necessary in order to complex half of the dissolved manganese. 

Until more information on the concentrations of manganese in natural 

water is available, the importance of complex formers such as these in 

influencing the behavior of other solutes cannot be evaluated.

COMPLEXES WITH HIGHER OXIDATION STATES OF

MANGANESE

The oxidation state of manganese, which normally is to be expected 

in natural water, is Mn+2 . Thermodynamic data in the literature 

show that oxidation of the divalent form will normally produce 

Mn+4 (Hem, 1963), and most aqueous species of trivalent manganese 

which might be formed as intermediate products are unstable. Some 

organic complexes of trivalent manganese are known, however. For 

example, complexes of oxalate ions with trivalent manganese are 

formed at an intermediate stage in the oxidation of oxalate by permanganate in acid solution. The mechanism of this oxidation reaction has been rather extensively studied and several values for 

stability constants of the oxalate complexes with trivalent manganese 

are quoted in Bjerrum, Schwarzenbach, and Sillen (1957). Although 

the complexes are strong in the sense that most trivalent manganese 

in solution is tied up on the complexes, they are very susceptible to 

oxidation. The literature is in general agreement that trivalent 

manganese is unstable in aqueous solution and can be expected to 

revert either to divalent or to quadrivalent forms. Strictly speaking, 

an equilibrium model therefore need not consider trivalent species. 

However, some consideration must always be given to metastability, 

and the rate of oxidation of divalent manganese was considerably 

retarded by the organic acids studied in the work here described. 

Some writers have expressed the view that manganese in natural 

water may often occur in the trivalent state, presumably either as a 

metastable complex or perhaps as a colloidal or moleoularly dispersed 

hydroxide.

The solid species of manganese oxide to be expected in natural 

environments were considered in an earlier paper (Hem, 1963). The 

standard free-energy values used for that report indicated that the 

hydroxide Mn(OH) 3 was unstable with respect to dehydrated species, 

but some indication of hydrated partly oxidized forms- was found in 

laboratory studies of oxidation of Mn+2 solutions in ?ir. It is possible, however, to assign a mixed-oxide composition to compounds 

such as Mn2O3 , which thus becomes MnO-MnO2 or to Mn3O4 which 

becomes 2MnO-MnO2 .

Studies of some of the manganese oxide precipitates obtained in 

the laboratory were made by X-ray diffraction methodr. and patterns 

corresponding to hydrohausmannite, a hydrated form of Mn3O4 , de-
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACTD8 D13

scribed by Frondel (1953) were sometimes obtained. (More recently, 

Feitknecht, Brunner, and Oswald (1962) have reported that 

material prepared and termed hydrohausmannite in earlier experiments was actually a mixture of hausmannite and beta MnO(OH).) 

The bulk of the material was too poorly crystallized to give any X-ray 

pattern. In some of the digallic and gallic acid solutions containing 

dissolved manganese, a flocculent black precipitate formed after the 

solutions had stood for several months in contact with rir. This 

material gave no distinguishable X-ray pattern. Some of the solutions took on a brown color, perhaps owing to some form of complex 

of partly oxidized manganese, but possibly also explainable as an 

organic oxidation product, or a colloidal suspension of manganese 

dioxide. The brown color could be extracted from these solutions 

with amyl alcohol, which suggests the material was organic and the 

result of oxidation of the organic acid.

In a series of solvent-extraction experiments, manganese was left 

behind in the water fraction. The water solutions were clear and 

colorless, and no precipitates of oxide formed. These facts suggest 

that the complexed manganese was originally in the Mn+2 form. 

Some additional evidence against higher oxidation states in the complexed manganese is the approximate agreement between the calculations of amounts complexed from the Eh-pH data and the results 

obtained using the stability constant where the Mn+2 state was assumed at the beginning. These calculations and the solvent extraction studies are discussed later in this paper.

This evidence is not complete proof that organic complexes of 

trivalent manganese do not occur in natural water. Under the conditions of these experiments, however, manganese evidently did not 

occur in that state.

SOLUTION OF MANGANESE DIOXIDE BY DIGALLIC ACID

The effectiveness of digallic acid as a reducing agent for manganese 

dioxide was studied in a series of experiments. Solutions containing 

10, 100, and 1,000 milligrams per liter of tannic acid reagent were 

prepared. Six aliquots of each solution were adjusted in pB by using 

potassium hydroxide or hydrochloric acid, as required, to giA^e a series 

of 18 solutions having pH's of 4.5, 5.5, 6.5, 7.5, 8.5, and 9.£. About 

500 ml of each solution was placed hi a polyethylene bottle, and 1.0 

g reagent-grade manganese dioxide powder was introduced. In addition to the 18 bottles containing tannic acid solutions, a group of six 

bottles was prepared containing distilled water, to which was added 

first enough acid or base to give the desired pH and then manganese 

dioxide. Air was not excluded from these solutions, owing to the 

permeability of polyethylene for gases.

769-002 65 3
D14 CHEMISTRY OF MANGANESE IN NATURAL WATER

The solutions were allowed to stand at a temperature of 25°C 

and, from time to time, portions were removed for determination of 

pH and of redox potential. At the same time, an aliquot was filtered 

through a plastic membrane filter having pore diameters of 0.45/i, 

and analyzed for dissolved manganese using the standard periodate 

procedure. The bottles were shaken about once a day to keep the 

contents mixed. The solutions remaining after 2 months were 

filtered into new bottles, and some further experiments v^ere performed 

on these portions.

The results of the determinations are given in tables 2-5. Some 

indication of the nature of the reactions that took place in these 

solutions can be gained by examination of these data. General 

trends can be summarized as follows:

1. Manganese in solution at first increased with time, but tended 

eventually to reach a stable plateau.

2. More manganese went into solution when the amount of tannic 

acid was increased.

3. Amounts of manganese dissolved were greatest in solutions that 

were started at low pH.

4. The solutions containing tannic acid at an initial pH above 7 

decreased in pH on standing. The more acid solutions increased 

in pH. Final pH values were generally between 6.5 and 7.0.

5. The Eh of all the solutions tended generally to increase with time, 

at first rapidly and then more slowly.

6. Very little manganese was dissolved in the absence of tannic acid.

7. Solutions containing tannic acid and manganese were colored 

green in the early part of the experiment but later a brown color 

developed which tended to obscure the original green. The 

solutions still absorbed light strongly at the wave length used 

in studies of the complex.

TABLE 2. Manganese concentration and pH of solutions in contact with 

manganese dioxide

[Tannic acid absent]

Time (hours)

0..-. _ . ____ . ___ ....

3.... _.-. .._____.__

24.........................

72..... .-. _ .

144.- _ . __ ... .. ___ ....

720... ___ .. ___ . .......

1,320......................

Manganese concentration and pH of solutions in contact with manganese 

dioxide for the initial pH indicated

4.5

Mn

0.0

.11 

.12 

.24 

.51 

.61

PH

4.50 

4.97

4.85 

4.98 

4.98 

5.20 

5.60

5.5

Mn

0.00

.11 

.12 

.17 

.33

.36

pH

5.50 

5.85

5.65 

5.72 

5.71 

5.58 

5.88

6.5

Mn

0.00

.07 

.07 

.06 

.13 

.36

PH

6.50 

6.48

6.07 

6.27 

6.20 

5.81 

5.43

7.5

Mn

0.00

.07 

.07 

.06 

.06 

.08

PH

7.50 

6.76

6.59 

6.58 

6.38 

6.20

a 01

8.5

Mn

O.C"

.C3 

.04

.OR

.o«

.08

pH

8.50 

6.90

6.77 

6.69 

6.40 

6.21 

6.03

9.5

Mn

0.00

.07 

.04 

.01 

.03 

.04

PH

9.50 

9.08 

8.80 

8.01 

7.00 

6.41 

6.32
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D15

The rate at which manganese was dissolved is indicated graphically 

in figure 1 for solutions initially brought to pH 5.5 and pH 8.5. The

0 200 400 600 800 1000 1200 1400 1600 1800 2000 2200 2400 2600 2800

TIME, IN HOURS

A. Initial pH 5.5.

0 200 400 600 800 1000 1200 1400 1600 1800 2000 2200 2400 2600 2800

TIME, IN HOURS

B. Initial pH 8.5. 

FIGURE 1. Solution of manganese from manganese dioxide by diflerent concentrations of tannic acidj
TABLE 3. Manganese concentration, pH, and Eh of 10-ppm tannic ac

Time (hours)

0- _______ ...

3. _ ....... __ . 

21 __ ..... .....

70 _ ..... ___ .

142. ...... _ .

646. . -

1,320.... ...

Manganese concentration, pH, and Eh of 10-ppm tannic acid solution in c

4.6

! 

i

0.0

".~64~ 

1.1 

1.8 

3.0 

3.6

a

4.50 

5.28 

5.55 

6.06 

6.15 

6.14 

6.19

Eh (volts)

 o 

S

O

0.58 

.65 

.56 

.56 

.61 

.60

 o 

a>

! a 

o

~o~n~

.64 

.63 

.62 

.62

5.S

!a 

!*

0.0

""53" 

1.1

2.0 

3.1 

3.5

W 

P,

5.50 

6.17 

6.10 

6.39 

6.27 

6.32 

6.60

Eh (volts)

 0

ID 

O

0.63 

.52 

.65 

.55 

.59 

.59

 a 

o

0.64 

.60 

.61 

.60

.57

6.5

a

1

0.0

.51 

.88 

1.1 

2.0 

2.4

W a

6.50 

6.60 

6.50 

6.66 

6.42 

6.32 

6.60

Eh (volts)

O

0.49 

.50 

.53 

.53 

.69 

.59

1

"3 

O

"o'ecf

.68 

.60 

.61 

.57

!a 

8

0.0

""47"

.88 

1.2 

1.8 

1.8

TABLE 4. Manganese concentration, pH, and Eh of 100-ppm tannic a

Manganese concentration, pH, and Eh of 100-ppm tannic acid solution in 

Tims (bo"rs)

0..... .....

3. ...... ......

21.. .. .

4.6

Mn (ppm)

0.0

1.7

W 

a

4.50

6.58

5.83

Eh (volts)

Observed

0.54

.52

Calculated

0.67

6.S

Mn (ppm)

0.0

1.3

W 

a

6.50 

6.02 

6.98

Eh (volts)

Observed

0.62 

.63

"S

1

a 66

6.S

1

0.0

1.2

W

Q,

6.50 

6.62 

6.30

Eh (volts)

Observed

0.49 

.51

Calculated

0.61

Mn (ppm)

0.0

1.2
70_ ....... 

142 .............

670 _ _ _ . .. 

1,320 

2,600. 

3.2 

4.9 

20 

21

5.99 

5.90 

6.45 

6.87 

6.90

.47 

.50 

.58 

.56 

.54

.64 

.66 

.56 

.50 

.50

2.8 

4.2 

14 

14 

214

5.98 

5.98 

6.58 

6.78 

6.89

.48 

.50 

.57 

.57 

.54

.64 

.63 

.55 

.52 

.51

2.3 

3.2 

13 

13

313

6.16 

6.00 

6.37 

6.91 

6.78

.49 

.50 

.57 

.56 

.54

.62 

.64 

.57 

.50 

.52

1.6 

2.3

8.7

no

9.3

6.3

6.1

6.2

6.9

6.7

Percentages of manganese as complex, calculated from stability constant and from Eh and pH, respecti

1 46 and 79. 3 53 and 98. ' 67 and 99. 

^ 59 and 88. « 61 and 99. « 66 and 99. 

TABLE 5. Manganese concentration, pH, and Eh of 1,000-ppm tannic acid 

Time (hours)

0 

24....... ....

72 

144..............

646.. 

1,320... _____ .

2,500 

Manganese concentration, pH, and Eh of 1,000-ppm tannic acid solution in co

4.5

I 

ft 

3 

a

a

0

3.6

8.4 

15 

71 

112

a

4.50

4.88

5.12 

5.50 

5.57 

5.58 

6.35 

6.68

Eh (volts)

1fa

O

0.55

.55 

.51 

.51 

.47 

.51 

.54

03

1

O

0.74 

.68 

.67 

.65 

.54

5.5

1

3 

a

0

1.3 

7.3

12 

59 

90

5.50 

5.88 

5.70 

5.5 

5.57 

5.65 

6.25 

6.59

Eh (volts)

sW3

fl

O

6.48 

.51 

.53 

.53 

.49

 a

03 "3

 3 

O

0.69 

.70 

.67 

.65

.57

6.5

1 

3 

a

a

0

1.2 

2.2 

7.1 

50 

SO 

89

W 

a

6.50 

6.49

5.96 

5.66 

5.75 

5.99 

6.40 

6.68

Eh (volts)

 a 

I

O

6.48

.51 

.54 

.53 

.51 

.54 

.55

 a

"31

"3 

0

0.65 

.68 

.66 

.60 

.55 

.52

a 

a

a

a

0

1.8 

2.2 

2.4 

32 

54 

59

M 

a

7.5

6.5

6.0

6.0

6.0

6.4

6.6

... 
D18' CHEMISTRY OF MANGANESE IN NATURAL VATER

solutions containing the lower concentrations of tannic acid appeared 

generally to reach a state of equilibrium, and mangr.nese concentration leveled off after about 30 days.

A rigorous evaluation of the results in tables 2-5 and in figure 1 is 

complicated by several factors. The results do indicate, of course, 

that manganese reduction by the digallic acid takes place and that in 

most solutions an apparent equilibrium condition was reached. Consideration of the chemical data and the appearance of the solutions 

during the experiment provide some basis for hypotheses regarding 

the reactions. If equilibrium is reached, several reactions will be involvdd. One is the reduction of manganese dioxide, liberating soluble 

manganese. Divalent manganese was probably produced rather than 

Mn+3 because the more reduced form is the more stable, thermodynamically, and because the colors which developed in the standing 

samples resembled the ones produced by mixing divrlent manganese 

solutions with digallic acid in the spectrophotometric work. The 

appearance of these colors suggests that the complexing reaction in 

which divalent manganese is combined with digallate also must be 

taken into account. The process of oxidation undoubtedly alters the 

digallic acid also, but the oxidized material may still be capable of 

complexing manganese.

Polyethylene is permeable to gases, and diffusion of oxygen and 

carbon dioxide into the bottles probably took place. The trends in 

pH observed in solutions containing no tannic acid (table 2) were 

probably the result of carbon dioxide solution. Eh values for solutions that contained no tannic acid were probably influenced by dissolved oxygen and were not reported because they are probably not 

related to dissolved manganese.

Solutions listed in table 2 reached very low concentrations of dissolved manganese because very little oxidizable material was present. 

The solutions at lowest pH did contain some chloric1 e, added in the 

process of lowering the pH with hydrochloric acid, and some dissolved 

manganese could have resulted in these solutions from the oxidation 

of chloride to chlorine. Small amounts of soluble manganese also 

could have been derived from impurities in the manganese dioxide. 

The significance of concentrations below 0.10 ppm is somewhat doubtful because the method of determination has low accuracy and precision in that range.

To interpret the results in tables 3-5, the role of the tannic acid is 

assumed to be to reduce manganese from the Mn+4 st ate to the Mn+2 

state, and to complex part of the Mn+2. These reactions proceed 

simultaneously, approaching an equilibrium in which the Eh and pH 

are those compatible with the activity of Mn+2 in the system. In
REDUCTION AND COMPLEXING OP MANGANESE BY GALLIC ACIDS D19

the earlier stages of the reaction, the Eh probably represent? a system 

involving digallic acid and dissolved oxygen.

In earlier work (Hem, 1963), a series of Eh-pH diagrams were 

prepared for manganese systems. From these diagrams an equilibrium Eh can be calculated for any activity of dissolved Mn+2 and 

hydrogen ions. Comparison of this calculated Eh with that actually 

measured gives an indication as to the degree to which equilibrium is 

reached in the solution. If the calculated Eh is greater than the 

measured Eh, the solution is unsaturated and is still capable of reducing and dissolving more manganese. If the calculated Eh is less 

than the measured value, supersaturation with manganese is indicated. This apparent supersaturation could occur in an equilibrium 

condition as the result of complexing, by which the effective concentration of Mn+2 is decreased to produce MnT*, the 1:1 complex with 

digallate ions. The dissolved manganese determined is a total of all 

species present. The apparent supersaturation was used to compute 

percentage of manganese complexed in a few of the solutions listed 

in table 4, and will be discussed later in this report.

The calculated Eh values in tables 3-5 were obtained from largescale graphs of Eh, pH, and [Mn+2] based on the Nernst equation and 

the redox couple MnO2c+4H++2e=Mn+2aq+2H2O.

Because the ionic strength of the solutions was generally low, between 10~3 and 10~4, an activity coefficient of unity was assumed for 

Mn+2, except in the solutions which were highest in digallic acid and 

in manganese, where the activity coefficient for Mn+2 was estimated 

on the basis of the Debye-Hiickel law to be 0.8. The error introduced 

by this treatment is small compared to uncertainties in the measurement of Eh.

Reproducibility of Eh measurements in the solutions containing 

digallic acid was probably ±0.02 volt or more. Difference^ between 

observed and calculated Eh amounting to more than 0.02 volt in the 

solutions containing digallic acid may therefore be significant. At 

constant pH a change of 0.02 volt in Eh is equivalent to a change in 

manganese activity by a factor of about 2. Hence, the Eh is not a 

very sensitive indicator of manganese behavior. Some indications of 

interest, however, can be obtained by comparing observed and 

calculated Eh values listed in tables 3-5.

The observed Eh values in table 3 were at first considers bly below 

the calculated ones, but later in the experiment the calculated and 

observed values reached agreement, within 0.02 volt. This1 apparent 

equilibrium was reached earlier with the solutions started at a high 

pH. In general, more than 70 hours was required to reach equilibrium. In all but the one containing the smallest amount of dissolved 

manganese, the Eh data seem to show consistent supersaturation at
D20 CHEMISTRY OF MANGANESE IN NATURAL WATER

the time of the last observation (1,320 hr). The difference between 

observed and calculated Eh, however, never reached s, value greater 

than 0.02 volt in any of these solutions.

The data in table 4 represent systems where enough digallic acid 

was present to bring about more extensive reduction and complexing 

effects. The calculated and observed Eh values are m approximate 

agreement at 670 hours. The observations at 1,320 hours show that 

the amount of manganese in solution is greater than could be explained 

by assuming only Mn+2 ions are present. The solutiors were filtered 

after that series of observations and allowed to stand with no manganese dioxide present until the final set of analyses were made after a 

total elapsed time of 2,600 hours. All except two of the solutions 

showed some supersaturation at that time.

Table 5 lists results for the solutions containing 1,000 ppm of 

tannic acid. The amounts of manganese brought into solution were 

greater than in corresponding solutions containing lesser amounts of 

tannic acid. At the two lower pH levels the saturation point had not 

been reached at 1,320 hours. A later observation on these two solutions did not give a usable value for manganese concentration because 

the fractions to be used for this purpose were lost and no more solution was left. Equilibrium, however, was probably reached in these 

solutions between the 1,320- and 2,500-hour observations. The 

solutions having initial pH's of 6.5 or higher reached apparent equilibrium at 1,320 hours, and indicated supersaturation at 2,500 hours.

The general rend of the data is also indicated by fgure 1. The 

solutions containing lesser amounts of tannic acid gave curves which 

generally began to level off sooner than the more concentrated solutions. The illustrations also show the increased capacity for reductions of the solutions that had lower initial pH's. Because the 

graphs represent systems influenced by several variable^-, they cannot 

be very rigorously interpreted as indications of reaction rates, but 

they do show general trends in the establishment of equilibrium in 

the solutions.

CALCULATION OF AMOUNT OF MANGANESE COMPLEXED

The amount of manganese complexed in these solutions can be 

calculated from the available data and the stability constant for the 

1:1 complex, if certain assumptions are made. The equations required are:

[Mn+23[77 -]=1°3'9> (5)

2[Mn]=[Mn+2 ]+[Mn77+], (6)

2[r]-[T-]+[Mn77+]+[H.T], (7)
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC AC7DS D21

and

[HT] 

All activity coefficients are assumed to be unity, and the initial 

concentration of digallic acid is used as S[TJ. Neither of these 

assumptions is strictly correct, but for the dilute solution? used in 

these experiments the effect of activity coefficient corrections would 

be small. The effect of oxidation on the complexing properties of the 

digallate ion is not easily evaluated. An amount of digallic acid 

equivalent to perhaps two tenths of the manganese reduced would be 

oxidixed, but the oxidized digallate might still be capable of complexing manganese. Refinement of the assumption that the availability of digallate for complexing is not seriously affected by oxidation-reduction reactions seems impracticable. The dissociation 

constant for digallic acid, 10~6 -°, was estimated in earlier studies 

(Hem, 1960) and is only approximate.

Simultaneous solution of these equations is possible for any of the 

sets of observations reported in tables 3-5, and will give values for 

activity of complexed manganese as well as other dissolved species 

within the limits imposed by the assumptions. The computation, 

however, was made for only six observations, reported in table 4, 

where approximately constant dissolved manganese concentrations 

had been reached. For the six observations the extent of complex 

formation can also be estimated by using pH and Eh data.

In samples where measured Eh exceeded the calculated Eh, a 

situation indicating apparent supersaturation with Mn"1" 2, the activity 

of Mn+2 at equilibrium can be calculated by using measured pH 

and Eh values and the Nernst equation. The difference between 

the calculated activity of divalent manganese and the total dissolved 

manganese concentration determined by analysis can then be assigned 

to the complexed form. This method of evaluating the complexing 

effect has a rather low order of accuracy because small differences in 

Eh can bring about major differences in Mn+2 activity. Pesults of 

computations lay both methods are listed in table 4 in terms of percent of total manganese present as the complex.

The proportion of complexing of the manganese ranged from 46 to 

66 percent when calculated from the stability constant of the postulated complex. The degree of complexing calculated from the redox 

potentials was generally somewhat greater and ranged from 77 to 99 

percent. Although the values do not check very closely, agreement is 

about as good as could be expected in view of the uncertairty in the 

measured Eh and thus they lend some further support to the value for 

KI, the stability constant, which was calculated in this study.
D22 CHEMISTRY OF MANGANESE IN NATURAL W/TEB

SOLVENT-EXTRACTION EXPERIMENTS

A solvent-extraction technique was used in an attempt to isolate 

the digallic-manganese complex. The solutions which had stood in 

contact with solid manganese dioxide long enough to dissolve considerable Mn+2 were filtered and portions were extracted in separatory 

funnels with different organic solvents having low solubility in water. 

The solvents most effective in removing the color from the solutions 

containing complexed manganese were butyl and amyl alcohol. 

Carbon tetrachloride, benzene, xylene, hexane, amyl acetate, 4-methyl 

propanone, and chloroform all were tried but were not capable of 

removing either the color or the manganese from the v^ater fraction. 

Although the butyl and amyl alcohol extractions seemed to remove the 

colored material, the manganese remained behind the water layer.

The solvent-extraction experiments indicate that the complex does 

not have a strong enough linkage between the organic molecule and 

the manganese to survive the extraction, and this approach was 

abandoned.

STABILITY OF COMPLEX TOWARD OXIDATION

After the digallic acid solutions had attained a maximum concentration of dissolved manganese, a selected group representing different 

reagent concentrations and starting pH values were filtered through 

plastic membrane filters with pores 0.45 micron in diameter. Air was 

bubbled through the samples for 10 hours and they were again filtered 

and analyzed for manganese. No change in manganese content had 

occurred. The pH of the samples was then raised to 9.00 with sodium 

hydroxide and again air was bubbled through them for 10 hours. This 

time, after filtration, the manganese in solution had decreased by 

about 10 percent.

In previous experiments (Hem, 1963), rates of oxidation of divalent 

manganese were measured at pH 9.0. In the absence of sulfate and 

bicarbonate, one-half of the manganese was oxidized in 1^2 minutes. 

Half times of 60-80 minutes were observed in the presence of sulfate 

and bicarbonate. Obviously, the reaction is very mucb slower in the 

presence of digallic acid.

Experiments were made using solutions of manganese complexed 

with digallic and gallic acids in which the pH was raisei to about 9.5 

and was maintained at a selected value. Samples were periodically 

removed, filtered through a plastic membrane filter, and analyzed for 

manganese. The results indicated that the rate of loss of manganese 

attributable to the destruction of the complex by oxidation is appreciable. The half time indicated for pH 9.8 was about 8 hours for the 

manganese-digallate complex. The stability of the gallic acid complex 

toward oxidation was similar.
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDP D23 

SOLUTION OF MANGANESE DIOXIDE BY GALLIC ACID

A series of equilibrium experiments like the ones used for digallic 

acid was performed to evaluate the reducing and complexing- action 

of gallic acid on manganese dioxide. Because these organic acids 

are chemically similar to each other, a less elaborate series of experiments for gallic acid appeared to be reasonable and capable of demonstrating points of similarity or differences in behavior. Accordingly, a 

series of four 500-ml samples was prepared; two contained 50 ppm 

and two contained 200 ppm of gallic acid. The pH of one of the 

50-ppm gallic acid and that of one of the 200-ppm gallic acid solutions 

was raised to about 7.75. The other two solutions were not adjusted 

but were left at the pH near 4.0, reached as a result of the dissociation 

of the gallic acid. To each solution was added 1.0 g of reagent-grade 

manganese dioxide. The solutions were allowed to stand in polyethylene bottles, at 25°C, for about 5 months. During this time the 

pH, Eh, and manganese, which passed through plastic filter membrane 

with 0.45-ju diameter pores, were measured on three occasions1 , The 

results are listed in table 6. The data do not give enough points to 

prepare graphs but can be compared approximately with the results 

obtained in experiments with digallic acid.

In the first 18 hours all four solutions reduced some of the manganese and brought it into solution. As was observed in the earlier 

experiments with digallic acid, the reducing action was stronger at 

the lower pH, and the strong solution of gallic acid brought more 

manganese into solution than the weaker one. The observation at 

310 hours showed a considerable further increase in dissolved manganese. This dissolution followed the same pattern observed at the 

end of the first 18 hours. The final observation was made mo^e than 

4% months after the experiment was started and showed a smaller 

amount of dissolved manganese than the 310-hour measurement. 

On very long standing it would seem the gallic acid and the complex 

are destroyed or altered by oxidation by atmospheric oxygen.

The pH of the solutions tended to converge, the final value being 

above 7.0 and appreciably higher than was noted in solutions of 

digallic acid.

The calculated values for Eh reported in table 6, as in tables 2-5, 

are based on the equilibrium involving manganese dioxide and 

divalent manganese without complexing. Approximate agreement 

between observed and calculated Eh was obtained at 310 hours, 

except for one sample which seemed to reach this agreement earlier.

The calculated Eh was considerably less than the observed Eh 

at the time of the final observation; this fact indicates some degree of
D24 CHEMISTRY OF MANGANESE IN NATURAL WATER

TABLE 6. Manganese concentration, pH, and Eh of solutions of gallic acid in 

contact with manganese dioxide

Time (hours)

0 _ 

18 - 

310 __ _

3,380.. . ____

0_-____..__-

18 _ 

310 __ __..

3,380__.____

Manganese concentration, pH, and Eh of 50-ppm and 200-ppm gallic acid solutions 

lor the initial pH indicated

50 ppm gallic acid

Initial pH, 4.1

Mn 

(ppm)

5. 9 

13

9. 7

PH

4. 10 

5.33

7. 25 

7.80

Eh (volts)

Observed

0.49 

.49 

.46 

. 54

Calculated

0. 71 

.46 

.46

7.7

Mn 

(ppm)

3.2

8. 0 

4. 8

pH

7. 70 

6. 78 

7.38 

7. 45

Eh (volts)

Observed

0. 35 

.44

.47 

. 53

Calculated

0. 54 

.46 

. 46

200 ppm gallic acid

3.8

Mn 

(ppm)

12 

46 

39

pH

3.80 

4.38 

7. 01 

7.70

Eh (volts)

Observed

0.54 

.51

.48 

.52

Calculated

0. 81 

.48 

.40

7.8

Mn 

(Ppm)

6.3

24 

17

PH

7.80 

7.14 

7.45 

7.90

Eh (volts)

Observed

0. 32 

.37 

.47 

. 50

Calculated

0. 49 

.44 

.38

supersaturation. If it were assumed that the difference between 

observed and computed Eh could all be ascribed to effects of complexing of divalent manganese, nearly all the dissolved manganese would 

be assigned to the complexed form.

The amount of complexing to be expected in these solutions was 

computed by using equations analogous to the ones used for digallic 

acid. The stability constant of 102 - 16 for the 1:1 complex (Kachhawaha and Bhattacharaya, 1962) and the first dissociation constant 

for gallic acid of io~4 - 40 (Lange, 1956) were used. The computations 

gave a maximum of some 10 percent of total manganese complexed 

in the solutions containing 50 ppm gallic acid and atout 20 percent 

complexed in the solutions containing 200 ppm gallic p.cid.

One might assume that the 310-hour observation for the 200-ppm 

gallic acid solution started at pH 7.8 represents a condition less complicated by oxidation and related alterations that probably affected 

all the observations at 3,380 hours. If this is assumed then the 

proportion of manganese complexed calculated from the pH and Eh
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDF D25

data is still close to 90 percent of the total dissolved manganese. 

The computation based on the reported stability constant indicates 

that only 13 percent of the manganese is complexed.

Although these computations are inexact and do not seem sufficient 

for calculating a new stability constant for the manganese-gallate 

complex, they suggest that the complex is somewhat more sta^e than 

the published data would indicate; a value nearer 103 would agree more 

closely with results obtained in this work.

EFFECTS OF GALLIC ACID COMPARED WITH THOSE OF

DIGALLIC ACID

Comparison of the results of the experiments with gallic acid and 

with digallic acid suggests that gallic acid is more effective than digallic acid as a reducing agent, both with respect to the rate of reaction 

and to the amount of manganese reduced per unit weight of organic 

acid. However, digallic acid evidently forms a complex with divalent 

manganese that is more resistant to oxidation than the one formed 

with gallic acid.

The amount of manganese brought into solution after 18 hours, in 

solutions containing 50 ppm gallic acid, was greater than amounts in 

solution after a similar length of time in solutions at similar pH containing 100 ppm or even 1,000 ppm of digallic acid. Upon standing, 

a loss of manganese was observed in solutions containing gallic acid. 

This effect was generally absent or less pronounced in solutions containing digallic acid. The more rapid increase in dissolved manganese 

in early stages of the experiment, and the attainment of highe" manganese content per milligram of organic acid originally present when 

gallic acid was the reducing agent, demonstrate the superiority of 

gallic acid as a reducing agent. However, the apparent loss of 

manganese from solution by oxidation and the behavior of gallic acid 

and manganese in spectrophotometric experiments show that the 

complexing action of the digallic acid is stronger than that of gallic 

acid.

IMPORTANCE OF REDUCING AND COMPLEXING IN

NATURAL SYSTEMS

As this paper and its three predecessors dealing with the chemistry 

of manganese in natural water have pointed out, organic factors of 

various types are frequently important. Except for the water from 

some thermal springs and water which has been polluted, surface or 

ground water containing more than 0.10 ppm manganese may well 

have been affected by these organic factors. For example, the move-
D26 CHEMISTRY OF MANGANESE IN NATURAL VATER

ment of ground water into and out of sediments rich in organic debris 

provides an excellent mechanism for producing a lo^ Eh, reducing 

manganese oxides, and retaining the dissolved manganese in solution 

by complex formation.

Published data on composition of river water g:;ve few direct 

indications of the influence of organic matter in the occurrence of 

manganese in natural water. One might expect, for example, a positive correlation between manganese content and the amount of organic coloring matter present in river waters, at least in regions where 

manganese is readily available in the soil. Although only a few data 

exist for testing this hypothesis, because so few waters have been 

analyzed for manganese, no strong correlation is indicated. In 

pools adjacent to the channel of the Cacapon River in West Virginia, Slack (1964) found that increased manganeso concentration 

was due to leaching of leaves which had fallen or were otherwise 

brought into the pools, and the amount of manganese as well as some 

other elements was greatest in the most highly colored water. The 

leaves themselves were probably the source of manganese at this location. The data on composition of river waters of the world collected 

by Durum, Heidel, and Tison (1960) show no well-defined relation between color intensity of water and manganese content for the whole 

mass of data, but for some individual streams there is an indication 

of higher manganese content at times when the water is more highly 

colored.

Manganese determinations for samples of streams in the Ohio River 

basin show that frequently the higher manganese concentrations occur 

in the more highly colored waters, but there are many exceptions to 

this general trend (U.S. Geol. Survey, 1960), especially in streams 

where large amounts of manganese are brought in by disposal of 

industrial waste or by mine drainage.

The principal importance of processes involving organic materials 

in the occurrence of manganese in natural water is probably in connection with the reduction of manganese oxide. Absorption of 

manganese from soil by growing plants is no doubt a major factor in 

making the element available for solution in runoff. Tv e reduction of 

oxide by organic matter in soil moisture or in the bottom mud of lakes 

is probably also a common process. It is not diflBc^ilt to visualize 

mechanisms by which water might be induced to move from organicrich environments through zones containing oxidized manganese, as for 

example in collector wells near streams. The manganese content 

of water pumped from some wells might thus tend to increase with 

time.

The aim of the work which has been described was to present in a 

systematic way a summary of the chemical factors wl ich control the
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIF^ D27

behavior of manganese in relatively pure natural water at and near the 

land surface. The behavior of manganese in hydrothermal environments and in brines, or factors which might bring about accumulation 

of manganese minerals, may be related to some of the topics which 

have been discussed. These topics, however, represent areas which 

were not directly investigated.

Most of the theoretical concepts discussed in this series of papers 

were tested in the laboratory. Field investigations of manganese 

behavior in ground water, giving particular attention to mineral 

sources of the element in rocks and soils, and studies of arees where 

solution and deposition of manganese occur in beds of rivers and creeks 

would be desirable as a further test of the concepts. Colleagues in 

the Geological Survey and others who might be able to do so are urged 

to undertake these studies.

LITERATURE CITED

Bjerrum, Jannik, Schwarzenbach, Gerold and Sillen, L. G., 1957, Stability constants, part 1, Organic ligands: Spec. Pub. no. 6, Chem. Soc. (London), 1957, 

105 p.

Black, A. P. and Christman, R. F., 1963a, Characteristics of colored surface waters: 

Am. Water Works Assoc. Jour., v. 55. p. 753-770.

 1963b, Chemical characteristics of fulvic acids: Am. Water Works Assoc. 

Jour., v. 55, p. 897-912.

Durum, W. H. and Haffty, Joseph, 1963, Implications of the minor element content of some major streams of the world: Geochim. et Cosmochim. Acta, 

v. 27, p. 1-11.

Durum, W. H., Heidel, S. G. and Tison, L. J., 1960, World-wide runoff of dissolved solids: Pub. 51 Internat. Assoc. of Sci. Hydrology, p. 618-628.

Feitknecht, W., Brunner, P. and Oswald, H. R., 162, Uber den Einfluss der 

Feuchtigkeit auf die Oxydation von Manganhydroxid durch molecularen 

Sauerstoff: Zeitschr. anorg. u. allg. Chem., v. 316, p. 154-160.

Frondel, Clifford, 1953, New manganese oxides hydrohausmannite and woodruffite: Am. Mineralogist, v. 38, p. 761-769.

Hem, J. D., 1960, Complexes of ferrous iron with tannic acid: U.S. Gecl. Survey 

Water-Supply Paper 1459-D, p. 75-94.

 1963, Chemical equilibria and rates of manganese oxidation: U.S. Geol.

Survey Water-Supply Paper 1667-A, 64 p. 

Hutchinson, G. E., 1957, Treatise on Limnology: New York, John Wile7 & Sons,

Geography, Physics and Chemistry, v. 1, 1015 p. 

Ingols, R. S. and Wilroy, R. D., 1963, Mechanism of manganese solution in lake

waters: Am. Water Works Assoc. Jour., v. 55, p. 282-290.

Kachhawaha, M. S. and Bhattacharya, A. K., 1962, Gallate complex of manganese (II): Indian Acad. Sci. Proc., v. 32a, p. 376-382. 

Lamar, W. L. and Goerlitz, D. F., 1963, Characterization of carboxyli". acids in

unpolluted streams by gas chromatography: Am. Water Works ASFOC. Jour.,

v. 55, p. 797-802. 

Lange, N. A., 1956, Handbook of chemistry: Sandusky, Ohio, Handb. Publishers

Inc., 1969 p. 

Lieth, Helmut, 1963, The role of vegetation in the carbon dioxide content of the

atmosphere: Jour, of Geophys. Research, v. 68, p. 3887-3898.
D28 CHEMISTRY OF MANGANESE IN NATURAL WATER

Nygaard, Gunnar, 1953, On the productivity of five Danish voters: Internat.

Assoc. of Theoretical and Appl. Limnology Proc., p. 123-13?. 

Oborn, E. T., 1964, Intracellular and extracellular concentration of manganese

and other elements by aquatic organisms: U.S. Geol. Survey Water-Supply

Paper 1667-C, 18 p. 

Rawson, Jack, 1963, Solution of manganese dioxide by tannii acid: Art. 117

in U.S. Geol. Survey Prof. Paper 475-C, p. C218-C219. 

Shapiro, Joseph, 1964, Effect of yellow organic acids on iron and other metals in

water: Am. Water Works Assoc. Jour., v. 56, p. 1062-1082. 

Slack, K. V., 1964, Effect of tree leaves on water quality in the Cacapon River,

West Virginia: Art. 161 in U.S. Geol. Survey Prof. Paper 475-D, p. D181-

D185. 

U.S. Geological Survey, 1960, Quality of Surface Waters of th*. United States,

1957, parts 1-4, North Atlantic Slope Basins to St. Lawrence River Basin:

U.S. Geol. Survey Water-Supply Paper 1520, 641 p. 

Yoe, J. H. and Jones, A. L., 1944, Colorimetric determination of iron with disodium-1, 2-dihydroxybenzene-3,5,disulfonate: Indus, and Eng. Chem. Anal.

Ed., v. 16, p. 111-115.
Reduction and Complexing of 

Manganese by Gallic Acids

By JOHN D. HEM

CHEMISTRY OF MANGANESE IN NATURAL WATER

GEOLOGICAL SURVEY WATER-SUPPLY PAPER 1667-D

Chemical reactions with organic acids 

which influence the behavior of 

manganese in natural water

UNITED STATES GOVERNMENT PRINTING OFFICE, WASHINGTON : 1965
UNITED STATES DEPARTMENT OF THE INTERIOR 

STEWART L. UDALL, Secretary

GEOLOGICAL SURVEY 

Thomas B. Nolan, Director

For sale by the Superintendent of Documents, U.S. Government Printing Office 

Washington, D.C. 20402 - Price 20 cents (paper cover)
CONTENTS

Page

Abstract__ ______________________________________________________ Dl

Introduction._____________________________________________________ 1

Previous work____________________________________________________ 1

Behavior of manganese in lakes.____________________________________ 2

Organic materials in natural water.....______________________________ 4

Properties of tannic acid__________________________________________ 5

Digallic acid complex with manganese______________________________ 6

Determination of stability constant-_____-_-___-_--____-__-__-___ 8

Properties of gallic acid____________________________________________ 10

Complexes of Mn+2 with other organic ligands_______________________ 11

Complexes with higher oxidation states of manganese __________________ 12

Solution of manganese dioxide by digallic acid________________________ 13

Calculation of amount of manganese complexed___________________ 20

Solvent-extraction experiments._________________________________ 22

Stability of complex toward oxidation._______-_-__--_______-_-___ 22

Solution of manganese dioxide by gallic acid __________________________ 23

Effects of gallic acid compared with those of digallic acid_______________ 25

Importance of reducing and complexing in natural systems_____________ 25

Literature cited ___________________________________________________ 27

ILLUSTRATION

Pa.ee 

FIGURE 1. Solution of manganese by tannic acid at 04) pH 5.5 and (B*) pH

8.5___---_-----________--_-___-----____------_-_--__- D15

TABLES

Page 

. Computation of stability constant for manganese digallate

complex._____________________________________________ D9

2. Manganese concentration and pH of solutions in contact with

manganese dioxide. Tannic acid absent_________________ 14

3-6. Manganese concentration, pH, and Eh of 

3. 10-ppm tannic acid solution in contact with manganese

dioxide __________________________________________ 16

4. 100-ppm tannic acid solution in contact with manganese

dioxide ___________________-___-_--____--_-_--_--_ 16

5. 1,000-ppm tannic acid solution in contact with manganese dioxide..-.--_-----_-----_-------_-.-_-------- 17

6. Solutions of gallic acid in contact with manganese dioxide. 24

m

CHEMISTRY OF MANGANESE IN NATURAL WATER

REDUCTION AND COMPLEXING OF MANGANESE BY 

GALLIC ACIDS

By JOHN D. HEM

ABSTRACT

Tannic (digallic) and gallic acids form 1:1 complexes with Mn+2 in d''ute solution. The stability constant of the digallic complex was determined to be 103 -9 

by spectrophotometric studies. Water solutions of tannic and gallic acid can 

bring 100 parts per million or more of manganese into solution from solid manganese dioxide reagent. The reducing and complexing reactions are most rapid 

at pH's below 5. The manganese brought into solution was oxidized very slowly 

at pH's as high as 9.8.

INTRODUCTION

The biochemical factors that are involved in the chemical behavior 

of manganese in natural environments have been touched upon in 

the earlier papers in this series. A further study of some chemical 

reactions that may influence manganese solution and deposition is 

described here. The study concerns the behavior of oxidized and 

reduced forms of manganese in the presence of the organic compounds 

tannic acid and gallic acid. The rates of reduction of tetravalent 

manganese by the organic acids, and the form and stability of the 

soluble manganese compounds formed in the reduction process, were 

investigated. The results provide some indications as to the principles which govern the circulation of manganese in natural environments at or near the land surface, where organic material is plentiful.

PREVIOUS WORK

Preliminary studies of the behavior of divalent manganese in dilute 

solutions of tannic acid were described by the writer in an earlier 

report (Hem, 1963, p. A51-A52), but no definite conclusions were 

drawn. A green color was observed in solutions containing tannic 

acid and divalent manganese, and a similar color was noted when the

Dl
D2 CHEMISTRY OF MANGANESE IN NATURAL WATER

pH of the tannic acid solutions was increased in the absence of manganese. Tannic acid noticeably retarded the rate of oxidation of 

divalent manganese when the alkaline solutions were, exposed to the 

air. In some environments, such as bogs, that are rich in organic 

matter, the effects of organic solutes are unquestionably important. 

In this paper reduction and complexing effects have been studied in 

detail so that their importance could be intelligently evaluated. 

Some preliminary results of experiments with tannic-acid reduction 

of manganese dioxide have been reported by Rawson (1963).

BEHAVIOB OF MANGANESE IN LAKES

One of the more common water-quality problems encountered in 

the water-supply industry is the appearance of manganese in solution 

in the water withdrawn from a supply reservoir. Th^ literature contains many references to this type of problem, with descriptions of 

conditions found and methods used for coping with the problem. The 

causes, however, have not always been clearly understood or satisfactorily explained.

Vegetation requires manganese in its life processes, and some of 

the aquatic species of plants seem to accumulate the element (Oborn, 

1964). Hutchinson (1957, p. 805) states that the manganese content 

of lake water is directly related to the productivity of the lake. Lakes 

and reservoirs support a considerable volume of aquatic vegetation, 

ranging in size from diatoms to large water weeds rooted in the lake 

bottom, and during the growing season as much as 10 tons of dry 

cellulose and related material could be produced per acre of water 

surface (Lieth, 1963). This high production rate is probably more 

typical of sewage-oxidation ponds or peat-bog or swamp environments than of a clear lake, but any water impoundment can be expected to produce a considerable mass of organic matter each year. 

In his studies of productivity in several lakes in Dermark, Nygaard 

(1953) found a production rate of organic matter of 0.62 g per day per 

square meter in relatively clear water 28 meters deep. This rate is 

equivalent to more than a third of a ton per acre in a growing season 

of 150 days and is probably lower than the usual rate for a storage 

reservoir. In large part, the organic matter produced in each growing 

season decays, and a residue of low-solubility material settles to the 

lake bottom.

Organic matter of aquatic vegetation contains from a few hundredths to a few tenths of a percent of manganese (Oborn, 1964). 

The amount of manganese present depends partly on its availability. 

Bottom-rooted plants can extract manganese from the soil below the 

water surface, but free-floating forms can obtain it only from the 

water in which they grow.
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D3

The quantitative importance of biological productivity as a mechanism for trapping manganese in a lake can be estimated in several 

ways. The amount of manganese in solution brought into a reservoir 

during the growing season probably is normally almost all removed 

biologically. The median manganese concentration of North American 

rivers of major size is reported by Durum and Haffty (1903) to be 

20 micrograms per liter. Thus, in 1 acre-foot of water, one might 

expect to find about 25 g of manganese. The median concentration 

of 20 micrograms per liter is equivalent to 0.02 ppm (parts p<?r million) 

and is a higher concentration than might be expected from the rather 

sparse data for manganese content in the analytical records published 

by the U.S. Geological Survey in annual reports on quality of surface 

water of the United States. In keeping with longstanding custom, 

the dissolved constituents in river water are ordinarily determined on 

filtered or clarified samples. Durum and HafTty used samples which 

had been filtered and acidified, but the filtration did not r?move all 

the particulate matter. Hence, their results may include some of the 

participate manganese that was present. Hutchinson (1957) indicates 

that particulate manganese may be utilized by plants and in any 

event is likely to be precipitated in a standing water body such as a 

lake or reservoir.

A production rate of one-quarter ton of dry vegetation per acre of 

lake surface per year, of material containing 0.10 percent manganese, 

would represent a rather modest productivity and would require 250 

g of manganese per acre of lake surface. The amount of riangaiiese 

available from lake water having 20 micrograms of manganese pe 

liter, or 25 g of dissolved manganese per acre-foot, is enough to supply 

the biochemical demand for the whole year if the photosynthetic 

process is active to a depth of 10 feet. An annual inflow of water 

sufficient to renew the manganese supply would be required to maintain the process. Bottom-rooted aquatic species growing in shallow 

water would extract some of their manganese from the soil of the lake 

bottom. When the growing season ends, the crop of vegetation dies 

back and the debris settles toward the bottom of the kke. This 

debris carries with it at least a part of the manganese originally 

incorporated in the vegetation and forms an organic-rich precipitate.

Biological processes, therefore, tend to remove manganese from 

water and transfer it to a lake-bottom precipitate in significant 

amounts. The normal circulation of aerated water through the lake 

maintains oxidizing conditions and depresses manganese solubility. 

During periods of thermal stratification, howevei, the deeper water 

may be depleted in oxygen through biochemical processes to a degree 

that induces reduction and re-solution of manganese in \ratei near 

the bottom muds that contain the manganese placed there earlier
D4 CHEMISTRY OF MANGANESE IN NATURAL "W^TER

through the decay of vegetation. The volume of water, so affected, 

will generally be small compared with the volume from which this 

manganese was originally extracted. Thus, the effect of a 10-fold 

or 100-fold concentration could readily be achieved, leading to a 

manganese concentration in the bottom water of 0.2 ppm or 2.0 ppm 

instead of 20 micrograms per liter. Concentrations higl er than a few 

tenths parts per million lead to problems if introduced into water 

distribution systems. Intakes may obtain such water if they are 

oriented in such a way as to intercept bottom water, or if the bottom 

water migrates to other levels.

Although the biochemical processes of manganese removal and occasional re-solution can probably account for the usual occurrences of 

manganese in water of storage reservoirs, the actual processes in the 

bottom sediments may involve some additional factors related to 

organic complexing and reduction. In a few areas, such as the piedmont region of Georgia studied by Ingols and Wilroy (IS63), the manganese content of the material making up the original bottom of the 

reservoir may be high enough to be a significant source of manganese 

for solution.

ORGANIC MATERIALS IN NATURAL WATER

The nature of naturally occurring organic solutes in colored river 

and lake water has never been completely established, even though 

such materials have been extensively studied. Lamar and Goerlitz 

(1963), however, were able to identify 12 carboxylic acids in unpolluted 

surface water from streams in the States of California and Washington. 

Substances identified included acetic, propionic, butyric, valeric, 

caproic, lactic, malonic, maleic, succinic, adipic, fumarc, and oxalic 

acids. The concentrations observed were from 4 to 592 parts per 

billion. Other carboxylic acids of higher molecular weight were also 

isolated but could not be specifically identified by gas chromatography 

because they were relatively stable and nonvolatile. Such acids have 

commonly been referred to in the literature as "humic" or "fulvic," or 

are given other names, which imply a definite chemical formula or 

structural pattern. Actually these names have been used with various 

connotations by many different writers to refer to complex mixtures of 

high molecular weight organic acids and polymers which impart a 

brown or yellow color to water and can be extracted by certain empirical procedures. The color and probably the structure of these 

materials are partly a function of pH and it has not been agreed 

whether they occur in true solution or as a colloidal susp ension. The 

molecules of these substances may actually approach colloidal size.

Shapiro (1964) conducted extensive studies of organic coloring 

matter which he extracted from lake water which he refers to as
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D5

"yellow organic acids" and has determined to be a mixture of aliphatic 

polyhydroxy carboxylic acids.

Black and Christman (1963a, b) suggested that the organic material 

responsible for color observed in many natural river waters consists of 

hydroxycarboxylic acids which have some degree of unsaturation and 

aromaticity. Their material also contained between 1 and 2 percent 

nitrogen. Lamar and Goerlitz (1963) reported that the infrared spectrum of colored material obtained from one of their waters resembled 

that of metadigallic acid, and the studies of Black and Christman 

(1963b) also indicate that such a compound can be expected in colored 

natural water. Therefore, it is possible to obtain some useful indications of the behavior of manganese in natural waters by studying the 

reactions between manganese and digallic (tannic) acid. Tannic acid 

is readily available in relatively pure form. Another series of experiments was performed with gallic acid, C6H2 (OH) 3COOH, in the expectation that the simpler compound would be more stable ard easier 

to work with.

The results of experiments with well-defined chemical compounds 

can be expressed in relatively exact terms, with respect to the nature 

of reactions and products and to the stability of complex ions which 

are produced. However, results of experiments with the natural 

organic materials extracted from water are usually less exact.

PROPERTIES OF TANNIC ACID

Tannic acid (digallic acid) has the chemical formula (HO) 3C6COO 

C6H2 (OH) 2COOH, and can be represented structurally as

HOOK

It is formed by hydrolysis from the more complex substances called 

tannins. Tannins are commonly present in vegetation, and therefore 

natural water, which usually has some contact with plant debris, commonly contains tannic acid or chemically similar material in solution.

The reagent-grade tannic acid used in this work was a pale yellowbrown powder, which dissolved readily in water. Solutions containing 1 mg of reagent per ml of water were pale yellow, and at this concentration had a pH near 4.0.

In earlier work (Hem, 1960, p. 78) the writer estimated the first 

dissociation constant for tannic acid to be about 10~6 . The curves 

obtained in potentiometric titrations of the tannic acid with strong 

base gave no well-defined inflection points but seemed to suggest that

769-002 65 2
D6 CHEMISTRY OF MANGANESE IN NATURAL WATER

the carboxyl hydrogen is mostly titrated before the pherolic hydrogens 

are detached from the molecule. The amount of base consumed, 

however, suggests that some influence from the phenolic hydrogens is 

probably present at pH 8 and that this influence increases as pH 

increases.

Redox titrations of tannic acid at various pH's uring potassium 

permanganate and platinum and saturated calomel electrodes were 

performed to evaluate the reducing ability of the reagent. In these 

titrations the manganese was at first reduced from Mn+7 to Aln+2 , but 

ultimately a black flocculent precipitate, or a dark-brown colloidal 

suspension, containing Mn+4, was produced. These experiments indicate that solutions of digallic acid are definitely capal^e of reducing 

manganese to the Mn+2 state, and that their ability for reduction is 

increased by lowering the pH.

The chemical mechanisms involved in the reduction of permanganate 

by digallic acid were not definitely ascertained. The color of the solution suggests that the organic compound undergoes considerable 

alteration in the process. For the present study, tte processes of 

oxidation and reduction which are of interest are those involving Mn+2 

and Mn+4, and therefore no detailed investigations of the effects of 

permanganate on digallic acid seemed necessary.

The amount of permanganate which reacted with 10 mg (3.1X10~5 

moles) of digallic acid without precipitating manganese oxide at pH 7.0 

was about 6X10~5 moles. The conversion of one MnO^ ion to Mn+2 

involves the gain of five electrons. Hence, the reduction equivalency 

of digallic acid in the conditions of this particular experirient was about 

10 times its molality. One mole of tannic acid, therefore, should be 

capable of reducing 5 moles of Mn+4 to Mn+2.

If, in being oxidized, the digallic acid molecule losa,s all 5 of its 

phenolic groups, and they are replaced with hydrogen, the net effect on 

the nominal charges on the carbon atoms in the 2 benzene rings is 

equivalent to the loss of 10 electrons. This mechanism would, therefore, be in accord with the results of the permanganate titration.

DIGALLIC ACID COMPLEX WITH MANGANESE

In earlier work (Hem, 1960), digallic acid appeared to form 

complexes with ferrous iron, and the possibility of similar complex 

formation with manganous ions was investigated. Aliquots were 

prepared containing 5.0 mg and 10.0 mg of digallic acid and small 

amounts of Mn+2 ranging from 0.02 to 1.0 mg. These aliquots were 

adjusted to pH 7.0 with 0.016-molar sodium hydroxide, diluted to 

25 ml, and the intensity of the green color which formec1 was measured 

with a spectrophotometer after the aliquots stood for 2 hours to
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D7

permit complete color development. Maximum absorbance occurred 

at 400 m/z. A direct linear response was obtained for the lower 

concentrations of manganese (up to about 0.15 mg of manganese in 

25 ml of water). This response indicates a complex, whose behavior 

was in accord with Beers' law, in dilute solution. As the concertration 

of manganese was increased, however, a maximum color ir tensity 

was reached beyond which further increases in manganese content 

brought about no change.

The composition of the complex was determined from the ppectrophotometric data using the procedure of Yoe and Jones (1944). In 

this procedure the absorbance is determined for various mole ratios 

of ligand to manganese. Straight parts of the resulting curve are 

extrapolated to an intersection which indicates the mole ratic of the 

predominant complex. A 1:1 mole ratio of manganese and digallic acid 

was indicated, at least for the concentrations of the reactants used 

here (about 10~4 to 10~6 molar) and at a pH near or slightly above 

neutrality.

The intensity of the green color was a function not only of manganese 

and digallic acid activities, but also of pH; the rate at which th?- color 

developed was also influenced by pH. The color was more intense 

at pH 8 than at pH 7 and developed more rapidly at the higher pH. 

After about 2 hours, at pH 7 and pH 8, the color stabilized and did 

not change significantly during the next 12-16 hours. Some greenish 

color was observed in solutions of digallic acid without manganese, 

when made alkaline with sodium hydroxide. The hue in these 

solutions, however, was somewhat different from that developed when 

manganese was present, and the intensity was generally much less 

when manganese was absent.

The color in these solutions is apparently associated with changes 

of the molecular configuration of the digallic acid. Such charges are 

probably brought about to some extent by neutralization with strong 

base which may displace some of the phenolic hydrogens. The fact 

that the color development is stronger in the presence of manganese 

and at higher pH's suggests that the two positive charges of the 

manganese ion are probably attracted both to the carboxyl and 

phenolic sites on the digallic acid molecule when the hydrogens in 

these positions are displaced.

If the digallate ligand is attached to the Mn+2 cation, both at the 

carboxyl and phenolic positions, the 1:1 complex would be uncharged. 

Perhaps this form would become more abundant as the pH is increased. 

However, in the absence of positive evidence as to the charge of the 

complex species, a linkage at only one position has been assumed.
D8 CHEMISTRY OF MANGANESE IN NATURAL WATER 

DETERMINATION OF STABILITY CONSTANT

The stability constant of the manganous-digallate complex was 

determined from the spectrophotometric data using the following 

lines of reasoning:

The digallate ligand, represented by the symbol T~, is considered 

to be the form of tannic acid present in the solutions on removal of 

the carboxyl hydrogen, and possibly one or more of the phenolic 

hydrogens, as a result of raising the pH. Solutions used in the 

experiments were about 10~4 molar. In media having so low an 

ionic strength (near 10~4), the ion activity coefficients, represented 

by the symbol 7, are near unity. The quantities in brackets are ion 

activities, the C terms are measured concentrations.

The association equilibrium

when written in mass-law form is

[Mn!T+]

[Mn+2 ][r- (1)

and the constant KI is the stability constant for the 1:1 complex. 

The total amount of manganese present is either in the complexed or 

uncomplexed form

[Mn+2 ] [MnT+]

The dissociation of digallic acid can be assumed to be very nearly 

complete at pH 8.0, because the first dissociation constant is near 

10~6. Accordingly, the undissociated form can be igrored, and one 

may equate the total concentration of the ligand presert to two forms

The optical density observed in the solutions was use°I as a measure 

of the amount of manganese complexed. The spectrophotometer 

readings were adjusted for the effects of absorbance by digallic acid by 

means of blanks containing no manganese, and the measurements were 

all made with the same length of light path. Under these conditions, 

for any given manganese concentration, a maximum optical density 

can be obtained by increasing the proportion of digallic acid to manganese in the solution. This maximum value represents the point at 

which essentially all the manganese has been complexed. For solutions containing the same amount of manganese, but not enough
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D9

digallic acid to develop the maximum optical density, a lower opticaldensity value will be observed, and the ratio of total manganese to 

complexed manganese will follow the relation

<?Mn_maximum optical density 

M&T observed optical density (4)

Experimental data for seven different laboratory samples are given 

in table 1. They may be used to provide solutions for equations 1 to 4. 

In these solutions manganese concentrations ranged from 7.3X10" 6 to 

7.3X 10~ B molar, and digallic acid concentrations ranged from 1.2 X 10~4 

to 6.2X10"4 molar. Maximum ionic strength in these solutions was 

4.8X10"4. At this ionic strength, all activity coefficients are close 

to 1.0. Values for TMnr and for 7T cannot be readily computed from 

the Debye-Hiickel equation; therefore, they were assumed to be 

unity. This assumption introduces an error of less than 5 percent into 

the values for KI. Values for TMn+2 were obtained from tlx e DebyeHtickel equation using the concentrations of Mn+2, Cl~, Na+, and 

T~ without complexing to calculate the ionic strength. Tbe effect of 

complexing on the ionic strength could have been corrected for by 

further computation but this was not done because the spectrophotometric data could not be reproduced very closely. Values for KI range 

from 1.4X104 to 5.5 X103 . The mean of the seven values i? 8.2X103 

or 103 - 9.

The uncertainties in the experimental data limit the accuracy of 

the constant. The pH had a considerable effect on the depth and 

stability of the color developed, and most of the range observed in KI 

values in table 1 can probably be attributed to this factor. The true 

value of KI could be higher or lower than the value calculated here by 

a factor of 2.

TABLE 1. Computation of stability constant for manganese digallate complex

Sample

1. ... .... __ .................

2.. . . _ .... _ ......

3. _ . .--

4. _____________ . . ......

5................ _ .............

6 .

7. ................... ... .........

Total 

manganese 

2CMn (moles 

per liter)

3. 6X10~5 

3.6X10-' 

7.3X10-" 

7.3X10-6 

3. 6X1(H 

7. 3X10-« 

1. 6X10-5

Total 

digallic acid 

2CT (moles 

per liter)

2. 5X10-4 

6. 2X10-4 

2. 5X10-4 

1.2X10-4 

1. 2X10-4 

1.2X10-4 

1. 2X10-4

Absorbence

Maximum

0.44 

.44 

.88 

.46 

.79 

1.20 

.55

Observed

0.250 

.390 

.48 

.240 

.305 

.53 

.212

Ki

6. 4X10" 

1.4X104 

6. 2X103 

9.6X103 

6.0X103 

9.6X103 

5. 5X103 

8. 2X103

The behavior of ferrous iron in digallic acid solution at pH 5.0 was 

considered earlier by the writer (Hem, 1960). For an assumed 1:1 

complex the stability constant was estimated to be about 1013 , based
D10 CHEMISTRY OF MANGANESE IN NATURAL WATER

on the interference of digallic acid with the formation of the pink 

complex of dipyridyl and ferrous iron.

Ferrous complexes with organic ligands are considerably stronger, 

as a rule, than the corresponding manganese complexes. However, 

the difference between the stability constants for manganese and 

ferrous complexes with digallic acid seems excessive. The composition of the ferrous complex was not determined, and it may be that 

the simple ¥&T+ structure that was assumed is incorrect.

The chemical behavior of the digallic acid reagent used in this 

series of experiments was not exemplary for exacting quantitative 

studies. The solutions of the acid were not stable in contact with 

air and slowly darkened in color on standing. Change? in the molecular configuration and complexing behavior, which might have resulted, were avoided by using freshly prepared solutions, Complexing 

and neutralization reactions were relatively slow, and the green complex tended to change color after a day or so, probably because of 

oxidation. The amount of color development and the rate of complexing were strongly influenced by pH. Close agreement of results 

in successive sets of samples therefore was not always attained.

The strength of the digallic complex suggests that in solutions 

containing about 10 ppm of digallic acid and a few tenths of a part 

per million or less of manganese, the manganese would be about 50 

percent complexed. Solutions of this kind may occur in rivers and 

lakes, although the nature of the organic matter in such water is 

probably not the same as the material studied. Natural coloring 

matter can be expected to have significant effects both as a reducing 

agent and as a complexing agent.

PROPERTIES OF GALLIC ACID

Experiments using gallic acid were compared with those using 

tannic acid. The reagent grade gallic acid used in these experiments 

was a white powder, having the chemical formula (HO) 3C6H2COOHH2O. The solutions of this material were colorless at first but tended 

to darken with age. Both gallic and digallic acids seamed to serve 

as nutrients for fungi which grew luxuriantly in some of the solutions 

left in the open air. Biological growths, however, were prevented 

from forming in the solutions used in the actual experiments.

The first dissociation constant for gallic acid is reported to be 

4X10"5 (Lange, 1956). The acidic properties are therefore stronger 

than those of digallic acid. Potentiometric titrations, using glass and 

saturated calomel electrodes, indicated a single well-defined inflection 

at pH 6.2. The phenolic hydrogen on the gallic acid molecule appears 

not to be active as a source of protons until the pH has increased
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS Dll

considerably beyond 6.2, but phenolic hydrogens probably become 

available at about pH 8.

Gallic acid is an effective reducing agent and, like digrllic acid, 

reacts with permanganate ions to yield Mn+2.

In the titration of gallic acid solutions with sodium hydroxide, a 

green color was observed which increased in intensity when the pH 

increased. This color probably resulted from the disruption or 

rearrangement of the gallic acid molecules. A very similar color was 

observed when gallic acid and manganese solutions were mixed and 

the pH adjusted. When manganese was present, however, the color 

appeared more rapidly and at a lower pH. The color developed in 

the presence of manganese was a more yellowish shade of green.

The nature of the complex formed between gallic acid anc1 divalent 

manganese has been studied by Kachhawaha and Bhattacharya 

(1962). These investigators used higher concentrations of reactants 

and worked at a lower pH than in the work described here, but they 

did report a green color in some of their solutions. They reported a 

1:1 complex whose stability constant is 102 -16. This compares with 

the value 10 3 ' 9 for the 1 :1 complex of divalent manganese and digallic 

acid determined in the present experiments.

Observations in the current study indicate that the gallic acid 

complex is weaker than the digallic complex, although perhs ps not as 

much weaker as the difference between the two constants would 

indicate. A preliminary series of experiments in which the gallate 

complex was developed at pH 7 in the presence of different amounts 

of manganese and gallic acid gave results which indicatec1 the 1 : 1 

association constant was between 102 and 103 . The value of Kachhawaha and Bhattacharya is in the lower part of this range. Hence, 

refined experimental work to determine the constant mor<? closely 

was not undertaken.

COMPLEXES OF Mn+2 WITH OTHER ORGANIC LIGANDS

The compilation of Bjerrum, Schwarzenbach, and Sillen (1957) 

includes stability constants for a 1 :1 complex of divalent nranganese 

with several other organic acids which may occur in natural water. 

The acids and constants for the type reaction (where M is the metal 

ion and L is the organic ligand)

are as follows: Oxalic acid, l£=102 - 55-103 -96 ; malonic acid, K=1Q3 - . 

It would appear that complexes of divalent manganese with tH organic 

ligands likely to occur in natural waters have similar stabilities.
D12 CHEMISTRY OF MANGANESE IN NATURAL WATER

Concentrations of a few tens of parts per million of organic acids would 

be necessary in order to complex half of the dissolved manganese. 

Until more information on the concentrations of manganese in natural 

water is available, the importance of complex formers such as these in 

influencing the behavior of other solutes cannot be evaluated.

COMPLEXES WITH HIGHER OXIDATION SPATES OF

MANGANESE

The oxidation state of manganese, which normally is to be expected 

in natural water, is Mn+2. Thermodynamic data in the literature 

show that oxidation of the divalent form will nornally produce 

Mn+4 (Hem, 1963), and most aqueous species of trivalent manganese 

which might be formed as intermediate products are unstable. Some 

organic complexes of trivalent manganese are known, however. For 

example, complexes of oxalate ions with trivalent manganese are 

formed at an intermediate stage in the oxidation of oxalate by permanganate in acid solution. The mechanism of this oxidation reaction has been rather extensively studied and several values for 

stability constants of the oxalate complexes with trivalent manganese 

are quoted in Bjerrum, Schwarzenbach, and Sillen (1957). Although 

the complexes are strong in the sense that most trivalent manganese 

in solution is tied up on the complexes, they are very susceptible to 

oxidation. The literature is in general agreement that trivalent 

manganese is unstable in aqueous solution and can be expected to 

revert either to divalent or to quadrivalent forms. Strictly speaking, 

an equilibrium model therefore need not consider trivalent species. 

However, some consideration must always be given to metastability, 

and the rate of oxidation of divalent manganese war considerably 

retarded by the organic acids studied in the work here described. 

Some writers have expressed the view that manganese in natural 

water may often occur in the trivalent state, presumaFy either as a 

metastable complex or perhaps as a colloidal or molecularlv dispersed 

hydroxide.

The solid species of manganese oxide to be expected in natural 

environments were considered in an earlier paper (Hem, 1963). The 

standard free-energy values used for that report indicated that the 

hydroxide Mn(OH) 3 was unstable with respect to dehydrated species, 

but some indication of hydrated partly oxidized forms was found in 

laboratory studies of oxidation of Mn+2 solutions in air. It is possible, however, to assign a mixed-oxide composition to compounds 

such as Mn2O3 , which thus becomes MnO-MnO2 or to Mn3O4 which 

becomes 2MnO-MnO2 .

Studies of some of the manganese oxide precipitates obtained in 

the laboratory were made by X-ray diffraction methods, and patterns 

corresponding to hydrohausmannite, a hydrated form of Mn3O4 , de-
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D13

scribed by Frondel (1953) were sometimes obtained. (More recently, 

Feitknecht, Brunner, and Oswald (1962) have reported that 

material prepared and termed hydrohausmannite in earlier experiments was actually a mixture of hausmannite and beta MrO(OH).) 

The bulk of the material was too poorly crystallized to give any X-ray 

pattern. In some of the digallic and gallic acid solutions containing 

dissolved manganese, a flocculent black precipitate formed after the 

solutions had stood for several months in contact with air. This 

material gave no distinguishable X-ray pattern. Some of the solutions took on a brown color, perhaps owing to some form of complex 

of partly oxidized manganese, but possibly also explainalJe as an 

organic oxidation product, or a colloidal suspension of manganese 

dioxide. The brown color could be extracted from these solutions 

with amyl alcohol, which suggests the material was organic and the 

result of oxidation of the organic acid.

In a series of solvent-extraction experiments, manganese was left 

behind in the water fraction. The water solutions were clear and 

colorless, and no precipitates of oxide formed. These factn suggest 

that the complexed manganese was originally in the Mr+2 form. 

Some additional evidence against higher oxidation states in the complexed manganese is the approximate agreement between the calculations of amounts complexed from the Eh-pH data and tbe results 

obtained using the stability constant where the Mn+2 state was assumed at the beginning. These calculations and the solvent extraction studies are discussed later in this paper.

This evidence is not complete proof that organic complexes of 

trivalent manganese do not occur in natural water. Under the conditions of these experiments, however, manganese evidently did not 

occur in that state.

SOLUTION OP MANGANESE DIOXIDE BY DIGALLIC ACID

The effectiveness of digallic acid as a reducing agent for manganese 

dioxide was studied in a series of experiments. Solutions containing 

10, 100, and 1,000 milligrams per liter of tannic acid reagent were 

prepared. Six aliquots of each solution were adjusted in pH by using 

potassium hydroxide or hydrochloric acid, as required, to give a series 

of 18 solutions having pH's of 4.5, 5.5, 6.5, 7.5, 8.5, and 9.5. About 

500 ml of each solution was placed in a polyethylene bottle, and 1.0 

g reagent-grade manganese dioxide powder was introduced. In addition to the 18 bottles containing tannic acid solutions, a group of six 

bottles was prepared containing distilled water, to which was added 

first enough acid or base to give the desired pH and then manganese 

dioxide. Air was not excluded from these solutions, owing to the 

permeability of polyethylene for gases.

769-002 65 3
D14 CHEMISTRY OF MANGANESE IN NATURAL WATER

The solutions were allowed to stand at a temperature of 25 °C 

and, from time to time, portions were removed for determination of 

pH and of redox potential. At the same time, an aliquot was filtered 

through a plastic membrane filter having pore diameters of 0.45/z, 

and analyzed for dissolved manganese using the standard periodate 

procedure. The bottles were shaken about once a day to keep the 

contents mixed. The solutions remaining after 2 months were 

filtered into new bottles, and some further experiments were performed 

on these portions.

The results of the determinations are given in tabler 2-5. Some 

indication of the nature of the reactions that took pT ace in these 

solutions can be gained by examination of these data. General 

trends can be summarized as follows:

1. Manganese in solution at first increased with tune, but tended 

eventually to reach a stable plateau.

2. More manganese went into solution when the amount of tannic 

acid was increased.

3. Amounts of manganese dissolved were greatest in solutions that 

were started at low pH.

4. The solutions containing tannic acid at an initial pH above 7 

decreased in pH on standing. The more acid solutions increased 

in pH. Final pH values were generally between 6.5 and 7.0.

5. The Eh of all the solutions tended generally to kicrea.se with tune, 

at first rapidly and then more slowly.

6. Very little manganese was dissolved in the absence of tannic acid.

7. Solutions containing tannic acid and manganese were colored 

green in the early part of the experiment but later a brown color 

developed which tended to obscure the original green. The 

solutions still absorbed light strongly at the wave length used 

in studies of the complex.

TABLE 2. Manganese concentration and pH of solutions in contact with 

manganese dioxide

[Tannic acid absent]

Time (hours)

0 .-.-.-..._ ...___...._

3.... _ . __ . ______ ..

24.._.__.__________________

72... _______ ..........

144. __ ... __ .. __ ....

720... ____ ... ___ . ...

1,320.............. __ ....

Manganese concentration and pH of solutions in contac*; with manganese 

dioxide for the initial pH indicated

4.5

Mn

0.0

.11 

.12 

.24 

.51 

.61

PH

4.50 

4.97

4.85 

4.98 

4.98 

5.20 

5.60

5.5

Mn

0.00

.11 

.12 

.17 

.33 

.36

pH

5.50 

5.85

5.65 

5.72 

5.71 

5.58 

5.88

6.5

Mn

0.00

.07 

.07 

.06 

.13 

.36

pH

6.50 

6.48

6.07 

6.27 

6.20 

5.81 

5.43

7.5

Mn

0.00

.07 

.07 

.06 

.06 

.08

PH

7.50 

6.76

6.59 

6.58 

6.38 

6.20 

6.01

8.5

Mn

0.00

.03 

.04 

.05 

.06 

.08

pH

8.50 

6.90

6.77 

6.69 

6.40 

6.21 

6.03

9.5

Mn

0.00

.07 

.04 

.01 

.03 

.04

PH

9.50 

9.08 

8.80 

8.01 

7.00 

6.41 

6.32
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACID*" D15

The rate at which manganese was dissolved is indicated graphically 

in figure 1 for solutions initially brought to pH 5.5 and pH 8.5. The

0 200 400 600 800 1000 1200 1400 1600 1800 2000 2200 2400 2600 2800

TIME, IN HOURS

A. Initial pH 5.5.

0 200 400 600 800 1000 1200 1400 1600 1800 2000 2200 2400 2600 2800

TIME, IN HOURS

B. Initial pH 8.5. 

FIGURE 1. Solution of manganese from manganese dioxide by diflerent concentrations of tannic acidj
TABLE 3. Manganese concentration, pH, and Eh of 10-ppm tannic aci

Time (hours)

0 .............

3 ... ... . 

21 ________ .

70 

142... ....... __

646..............

1,320.... . ....

Manganese concentration, pH, and Eh of 10-ppm tannic acid solution in co

4.5

1

a 

a 

3

0.0

".~64~ 

1.1 

1.8 

3.0 

3.6

w0,

4.50 

5.28 

5.55 

6.06 

6.15 

6.14 

6.19

Eh (volts)

 a

CD 

0

0.68 

.55 

.56 

.56 

.61 

.60

0

"a72~ 

.64 

.63 

.62 

.62

5.5

!a 

3

0.0

"""53" 

1.1

2.0 

3.1 

3.5

W 

ft

5.50 

6.17 

6.10 

6.39 

6.27 

6.32 

6.60

Eh (volts)

S

,a 

O

0.53 

.52 

.55 

.55 

.59 

.59

1

 a 

o

0.64 

.60 

.61 

.60 

.57

6.5

?

a*

a

3

0.0

.51 

.88 

1.1 

2.0 

2.4

W 

ft

6.50 

6.60 

6.50 

6.66 

6.42 

6.32 

6.60

Eh (volts)

O

0.49 

.50 

.53 

.53 

.59 

.59

1

 3

O

0.60 

.58 

.60 

.61 

.57

!a

3

0.0

.47 

.88 

1.2 

1.8 

1.8

7

7

6

6

6

6

6

TABLE 4. Manganese concentration, pH, and Eh of 100-ppm tannic ac

Time (hours)

0.--_ . 

3 ... ..........

21 ____ . ......

Manganese concentration, pH, and Eh of 100-ppm tannic add solution in co

4.5

Mn (ppm)

0.0

1.7

£

4.50 

5.58 

5.83

Eh (volts)

I 

O

0.54 

.52

0

a 67

5.5

Mn (ppm)

0.0

1.3

s

5.50 

6.02 

5.98

Eh (volts)

Observed

0.52 

.63

0

0.66

6.5

a

0.0

1.2

W 

ft

6.50 

6.52 

6.30

Eh (volts)

s

o

0.49 

.51

O

0.61

1

a

0.0

1.2

7

7

6
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIES D17

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D1S CHEMISTRY OF MANGANESE IN NATURAL W/TER

solutions containing the lower concentrations of tannic acid appeared 

generally to reach a state of equilibrium, and manganese concentration leveled off after about 30 days.

A rigorous evaluation of the results in tables 2-5 and in figure 1 is 

complicated by several factors. The results do indicate, of course, 

that manganese reduction by the digallic acid takes place and that in 

most solutions an apparent equilibrium condition was reached. Consideration of the chemical data and the appearance of the solutions 

during the experiment provide some basis for hypotheses regarding 

the reactions. If equilibrium is reached, several reactions will be involved. One is the reduction of manganese dioxide, liberating soluble 

manganese. Divalent manganese was probably produced rather than 

Mn+3 because the more reduced form is the more stable, thermodynamically, and because the colors which developed in the standing 

samples resembled the ones produced by mixing divalent manganese 

solutions with digallic acid in the spectrophotometric work. The 

appearance of these colors suggests that the complexing reaction in 

which divalent manganese is combined with digallate also must be 

taken into account. The process of oxidation undoubtedly alters the 

digallic acid also, but the oxidized material may still be capable of 

complexing manganese.

Polyethylene is permeable to gases, and diffusion cf oxygen and 

carbon dioxide into the bottles probably took place. The trends in 

pH observed in solutions containing no tannic acid (table 2) were 

probably the result of carbon dioxide solution. Eh values for solutions that contained no tannic acid were probably influenced by dissolved oxygen and were not reported because they are probably not 

related to dissolved manganese.

Solutions listed in table 2 reached very low concentrations of dissolved manganese because very little oxidizable material was present. 

The solutions at lowest pH did contain some chloride, added in the 

process of lowering the pH with hydrochloric acid, and some dissolved 

manganese could have resulted in these solutions from the oxidation 

of chloride to chlorine. Small amounts of soluble manganese also 

could have been derived from impurities in the manganese dioxide. 

The significance of concentrations below 0.10 ppm is somewhat doubtful because the method of determination has low accuracy and precision in that range.

To interpret the results in tables 3-5, the role of the tannic acid is 

assumed to be to reduce manganese from the Mn+4 stat e to the Mn+2 

state, and to complex part of the Mn+2. These reactions proceed 

simultaneously, approaching an equilibrium in which tl 3 Eh and pH 

are those compatible with the activity of Mn+2 in th<s system. In
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D19

the earlier stages of the reaction, the Eh probably represents a system 

involving digallic acid and dissolved oxygen.

In earlier work (Hem, 1963), a series of Eh-pH diagrams were 

prepared for manganese systems. From these diagrams an equilibrium Eh can be calculated for any activity of dissolved Mn+2 and 

hydrogen ions. Comparison of this calculated Eh with that actually 

measured gives an indication as to the degree to which equilibrium is 

reached in the solution. If the calculated Eh is greater than the 

measured Eh, the solution is unsaturated and is still capable of reducing and dissolving more manganese. If the calculated Eh is less 

than the measured value, supersaturation with manganese is indicated. This apparent supersaturation could occur in an equilibrium 

condition as the result of complexing, by which the effective concentration of Mn+2 is decreased to produce MnT+, the 1:1 complex with 

digallate ions. The dissolved manganese determined is a total of all 

species present. The apparent supersaturation was used to compute 

percentage of manganese complexed in a few of the solutions listed 

in table 4, and will be discussed later in this report.

The calculated Eh values in tables 3-5 were obtained from largescale graphs of Eh, pH, and [Mn+2] based on the Nernst equation and 

the redox couple MnO2c+4H++2e=:Mn+2aq+2H2O.

Because the ionic strength of the solutions was generally low, between 10~3 and 10~4, an activity coefficient of unity was assiimed for 

Mn+2, except in the solutions which were highest in digallic acid and 

in manganese, where the activity coefficient for Mn+2 was estimated 

on the basis of the Debye-Hiickel law to be 0.8. The error introduced 

by this treatment is small compared to uncertainties in the measurement of Eh.

Reproducibility of Eh measurements in the solutions containing 

digallic acid was probably ±0.02 volt or more. Differences between 

observed and calculated Eh amounting to more than 0.02 vcH in the 

solutions containing digallic acid may therefore be significant. At 

constant pH a change of 0.02 volt in Eh is equivalent to a change in 

manganese activity by a factor of about 2. Hence, the Eh is not a 

very sensitive indicator of manganese behavior. Some indications of 

interest, however, can be obtained by comparing observed and 

calculated Eh values listed in tables 3-5.

The observed Eh values in table 3 were at first considerably below 

the calculated ones, but later in the experiment the calculated and 

observed values reached agreement, within 0.02 volt. This apparent 

equilibrium was reached earlier with the solutions started »,t a high 

pH. In general, more than 70 hours was required to reach equilibrium. In all but the one containing the smallest amount of dissolved 

manganese, the Eh data seem to show consistent supersaturation at
D20 CHEMISTRY OF MANGANESE IN NATURAL VATER

the time of the last observation (1,320 hr). The difference between 

observed and calculated Eh, however, never reached a value greater 

than 0.02 volt in any of these solutions.

The data in table 4 represent systems where enough digallic acid 

was present to bring about more extensive reduction and complexing 

effects. The calculated and observed Eh values are in approximate 

agreement at 670 hours. The observations at 1,320 tours show that 

the amount of manganese in solution is greater than could be explained 

by assuming only Mn+2 ions are present. The solutions were filtered 

after that series of observations and allowed to stand with no manganese dioxide present until the final set of analyses were made after a 

total elapsed tune of 2,600 hours. All except two cf the solutions 

showed some supersaturation at that tune.

Table 5 lists results for the solutions containing 1,000 ppm of 

tannic acid. The amounts of manganese brought into solution were 

greater than in corresponding solutions containing lesser amounts of 

tannic acid. At the two lower pH levels the saturation point had not 

been reached at 1,320 hours. A later observation on these two solutions did not give a usable value for manganese concentration because 

the fractions to be used for this purpose were lost and no more solution was left. Equilibrium, however, was probably reached in these 

solutions between the 1,320- and 2,500-hour observations. The 

solutions having initial pH's of 6.5 or higher reached apparent equilibrium at 1,320 hours, and indicated supersaturation at 2,500 hours.

The general rend of the data is also indicated by figure 1. The 

solutions containing lesser amounts of tannic acid gave curves which 

generally began to level off sooner than the more concentrated solutions. The illustrations also show the increased capacity for reductions of the solutions that had lower initial pH's. Because the 

graphs represent systems influenced by several variables, they cannot 

be very rigorously interpreted as indications of reaction rates, but 

they do show general trends in the establishment of equilibrium in 

the solutions.

CALCULATION OF AMOUNT OF MANGANESE CGMPLEXED

The amount of manganese complexed in these solutions can be 

calculated from the available data and the stability constant for the 

1:1 complex, if certain assumptions are made. The equations required are:

[Mn+2][T-r

2[Mn]=[Mn+2]+[MnT+], (6)

(7)
D21

and

+ - ,o , } ~ W ' (b) [RT]

All activity coefficients are assumed to be unity, and the initial 

concentration of digallic acid is used as S[TJ. Neither of these 

assumptions is strictly correct, but for the dilute solutions used in 

these experiments the effect of activity coefficient corrections would 

be small. The effect of oxidation on the complexing properties of the 

digallate ion is not easily evaluated. An amount of digallic acid 

equivalent to perhaps two tenths of the manganese reduced would be 

oxidixed, but the oxidized digallate might still be capable of complexing manganese. Refinement of the assumption that tte availability of digallate for complexing is not seriously affected by oxidation-reduction reactions seems impracticable. The dissociation 

constant for digallic acid, 10~6 -°, was estimated in earlier studies 

(Hem, 1960) and is only approximate.

Simultaneous solution of these equations is possible for any of the 

sets of observations reported in tables 3-5, and will give values for 

activity of complexed manganese as well as other dissolved species 

within the limits imposed by the assumptions. The computation, 

however, was made for only six observations, reported in table 4, 

where approximately constant dissolved manganese concentrations 

had been reached. For the six observations the extent of complex 

formation can also be estimated by using pH and Eh data.

In samples where measured Eh exceeded the calculated Eh, a 

situation indicating apparent supersaturation with Mn+2, the activity 

of Mn+2 at equilibrium can be calculated by using measured pH 

and Eh values and the Nernst equation. The difference between 

the calculated activity of divalent manganese and the total dissolved 

manganese concentration determined by analysis can then be assigned 

to the complexed form. This method of evaluating the complexing 

effect has a rather low order of accuracy because small differences in 

Eh can bring about major differences in Mn+2 activity. F.esults of 

computations by both methods are listed in table 4 in terms of percent of total manganese present as the complex.

The proportion of complexing of the manganese ranged from 46 to 

66 percent when calculated from the stability constant of the postulated complex. The degree of complexing calculated from the redox 

potentials was generally somewhat greater and ranged from 77 to 99 

percent. Although the values do not check very closely, agreement is 

about as good as could be expected in view of the uncertainty in the 

measured Eh and thus they lend some further support to the value for 

KI, the stability constant, which was calculated in this study.
D22 CHEMISTRY OF MANGANESE IN NATURAL WATER

SOLVENT-EXTRACTION EXPERIMENTS

A solvent-extraction technique was used in an attempt to isolate 

the digallic-manganese complex. The solutions which had stood in 

contact with solid manganese dioxide long enough to dissolve considerable Mn+2 were filtered and portions were extracted in separatory 

funnels with different organic solvents having low solubility in water. 

The solvents most effective in removing the color from the solutions 

containing complexed manganese were butyl and r.myl alcohol. 

Carbon tetrachloride, benzene, xylene, hexane, amyl acetate, 4-methyl 

propanone, and chloroform all were tried but were not capable of 

removing either the color or the manganese from the water fraction. 

Although the butyl and amyl alcohol extractions seemed to remove the 

colored material, the manganese remained behind the water layer.

The solvent-extraction experiments indicate that the complex does 

not have a strong enough linkage between the organic molecule and 

the manganese to survive the extraction, and this approach was 

abandoned.

STABILITY OF COMPLEX TOWARD OXIDATION

After the digallic acid solutions had attained a maximum concentration of dissolved manganese, a selected group represerting different 

reagent concentrations and starting pH values were filtered through 

plastic membrane filters with pores 0.45 micron in diameter. Air was 

bubbled through the samples for 10 hours and they were again filtered 

and analyzed for manganese. No change hi manganese content had 

occurred. The pH of the samples was then raised to 9.00 with sodium 

hydroxide and again air was bubbled through them for 1C hours. This 

time, after filtration, the manganese in solution had decreased by 

about 10 percent.

In previous experiments (Hem, 1963), rates of oxidation of divalent 

manganese were measured at pH 9.0. In the absence of sulfate and 

bicarbonate, one-half of the manganese was oxidized in 1^ minutes. 

Half times of 60-80 minutes were observed in the presence of sulfate 

and bicarbonate. Obviously, the reaction is very much slower in the 

presence of digallic acid.

Experiments were made using solutions of manganese complexed 

with digallic and gallic acids in which the pH was raised to about 9.5 

and was maintained at a selected value. Samples were periodically 

removed, filtered through a plastic membrane filter, and analyzed for 

manganese. The results indicated that the rate of loss of manganese 

attributable to the destruction of the complex by oxidation is appreciable. The half tune indicated for pH 9.8 was about 8 hours for the 

manganese-digallate complex. The stability of the gallic acid complex 

toward oxidation was similar.
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIE * D23 

SOLUTION OF MANGANESE DIOXIDE BY GALLIC ACID

A series of equilibrium experiments like the ones used for digallic 

acid was performed to evaluate the reducing and complexing action 

of gallic acid on manganese dioxide. Because these organic acids 

are chemically similar to each other, a less elaborate series of experiments for gallic acid appeared to be reasonable and capable of demonstrating points of similarity or differences in behavior. Accordingly, a 

series of four 500-ml samples was prepared; two contained 50 ppm 

and two contained 200 ppm of gallic acid. The pH of one of the 

50-ppm gallic acid and that of one of the 200-ppm gallic acid solutions 

was raised to about 7.75. The other two solutions were not adjusted 

but were left at the pH near 4.0, reached as a result of the dissociation 

of the gallic acid. To each solution was added 1.0 g of reagent-grade 

manganese dioxide. The solutions were allowed to stand in polyethylene bottles, at 25°C, for about 5 months. During this time the 

pH, Eh, and manganese, which passed through plastic filter membrane 

with 0.45-/I diameter pores, were measured on three occasion^. The 

results are listed in table 6. The data do not give enough points to 

prepare graphs but can be compared approximately with the results 

obtained in experiments with digallic acid.

In the first 18 hours all four solutions reduced some of the manganese and brought it into solution. As was observed in the earlier 

experiments with digallic acid, the reducing action was stronger at 

the lower pH, and the strong solution of gallic acid brought more 

manganese into solution than the weaker one. The observation at 

310 hours showed a considerable further increase in dissolved manganese. This dissolution followed the same pattern observed at the 

end of the first 18 hours. The final observation was made more than 

4% months after the experiment was started and showed a smaller 

amount of dissolved manganese than the 310-hour measurement. 

On very long standing it would seem the gallic acid and the complex 

are destroyed or altered by oxidation by atmospheric oxygen.

The pH of the solutions tended to converge, the final value being 

above 7.0 and appreciably higher than was noted in solutions of 

digallic acid.

The calculated values for Eh reported in table 6, as in tables 2-5, 

are based on the equilibrium involving manganese dioxide and 

divalent manganese without complexing. Approximate agreement 

between observed and calculated Eh was obtained at 31C hours, 

except for one sample which seemed to reach this agreement earlier.

The calculated Eh was considerably less than the observed Eh 

at the time of the final observation; this fact indicates some degree of
D24 CHEMISTRY OF MANGANESE IN NATURAL W/TER

TABLE 6. Manganese concentration, pH, and Eh of solutions of gallic acid in 

contact with manganese dioxide

Time (hours)

0 ________

18 _____

310_. _______

3,380_______

0-__________

18____-____.

310__.______

3,3SO_______

Manganese concentration, pH, and Eh of 50-ppm and 200-ppm gallic acid solutions 

for the initial pH indicated

50 ppm gallic acid

Initial pH, 4.1

Mn 

(ppm)

5.9 

13 

9. 7

PH

4. 10 

5.33

7. 25 

7. 80

Eh (volts)

Observed

0. 49 

.49 

.46 

. 54

Calculated

0.71 

.46 

.46

7.7

Mn 

(ppm)

3. 2 

8. 0 

4. 8

pH

7.70 

6.78 

7.38 

7.45

Eh (volts)

Observed

0. 35 

.44 

.47 

.53

Calculated

0.54 

.46 

.46

200 ppm gallic acid

3.8

Mn 

(ppm)

12 

46 

39

pH

3.80 

4.38 

7. 01 

7.70

Eh (volts)

Observed

0.54 

.51

. 48 

.52

Calculated

0. 81 

.48 

.40

7.8

Mn 

(ppm)

6. 3

24 

17

PH

7. 80 

7. 14 

7.45 

7.90

Eh (volts)

Observed

0.32 

.37 

.47 

.50

Calculated

0.49 

.44 

.38

supersaturation. If it were assumed that the difference between 

observed and computed Eh could all be ascribed to effects of complexing of divalent manganese, nearly all the dissolved manganese would 

be assigned to the complexed form.

The amount of complexing to be expected in these solutions was 

computed by using equations analogous to the ones us<?-d for digallic 

acid. The stability constant of 102 - 16 for the 1:1 complex (Kachhawaha and Bhattacharaya, 1962) and the first dissociation constant 

for gallic acid of 10~4 - 40 (Lange, 1956) were used. The computations 

gave a maximum of some 10 percent of total manganese complexed 

in the solutions containing 50 ppm gallic acid and about 20 percent 

complexed in the solutions containing 200 ppm gallic acid.

One might assume that the 310-hour observation for the 200-ppm 

gallic acid solution started at pH 7.8 represents a condition less complicated by oxidation and related alterations that probably affected 

all the observations at 3,380 hours. If this is assumed then the 

proportion of manganese complexed calculated from the pH and Eh
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D25

data is still close to 90 percent of the total dissolved manganese. 

The computation based on the reported stability constant indicates 

that only 13 percent of the manganese is complexed.

Although these computations are inexact and do not seem sufficient 

for calculating a new stability constant for the manganese-gallate 

complex, they suggest that the complex is somewhat more stab!e than 

the published data would indicate; a value nearer 103 would agree more 

closely with results obtained in this work.

EFFECTS OF GALLIC ACID COMPARED WITH THOSE OF

DIGALLIC ACID

Comparison of the results of the experiments with gallic acid and 

with digallic acid suggests that gallic acid is more effective than digallic acid as a reducing agent, both with respect to the rate of reaction 

and to the amount of manganese reduced per unit weight of organic 

acid. However, digallic acid evidently forms a complex with divalent 

manganese that is more resistant to oxidation than the one formed 

with gallic acid.

The amount of manganese brought into solution after 18 hours, in 

solutions containing 50 ppm gallic acid, was greater than amounts in 

solution after a similar length of time in solutions at similar pH containing 100 ppm or even 1,000 ppm of digallic acid. Upon standing, 

a loss of manganese was observed in solutions containing gallic acid. 

This effect was generally absent or less pronounced in solutions containing digallic acid. The more rapid increase in dissolved marganese 

in early stages of the experiment, and the attainment of highe~ manganese content per milligram of organic acid originally present when 

gallic acid was the reducing agent, demonstrate the superiority of 

gallic acid as a reducing agent. However, the apparent loss of 

manganese from solution by oxidation and the behavior of gallic acid 

and manganese in spectrophotometric experiments show that the 

complexing action of the digallic acid is stronger than that of gallic 

acid.

IMPORTANCE OF REDUCING AND COMPLEXING IN 

NATURAL SYSTEMS

As this paper and its three predecessors dealing with the chemistry 

of manganese in natural water have pointed out, organic factors of 

various types are frequently important. Except for the water from 

some thermal springs and water which has been polluted, surface or 

ground water containing more than 0.10 ppm manganese mry well 

have been affected by these organic factors. For example, the move-
D26 CHEMISTRY OF MANGANESE IN NATURAL WATER

ment of ground water into and out of sediments rich in organic debris 

provides an excellent mechanism for producing a low Eh, reducing 

manganese oxides, and retaining the dissolved mangarese in solution 

by complex formation.

Published d.ata on composition of river water g: ve few direct 

indications of the influence of organic matter in the occurrence of 

manganese in natural water. One might expect, for example, a positive correlation between manganese content and the amount of organic coloring matter present in river waters, at least in regions where 

manganese is readily available in the soil. Although cnly a few data 

exist for testing this hypothesis, because so few waters have been 

analyzed for manganese, no strong correlation is indicated. In 

pools adjacent to the channel of the Cacapon River in West Virginia, Slack (1964) found that increased manganese concentration 

was due to leaching of leaves which had fallen or ^ere otherwise 

brought into the pools, and the amount of manganese as well as some 

other elements was greatest in the most highly colored water. The 

leaves themselves were probably the source of manganese at this location. The data on composition of river waters of the world collected 

by Durum, Heidel, and Tison (1960) show no well-defined relation between color intensity of water and manganese content for the whole 

mass of data, but for some individual streams there is an indication 

of higher manganese content at times when the water is more highly 

colored.

Manganese determinations for samples of streams in the Ohio River 

basin show that frequently the higher manganese concentrations occur 

in the more highly colored waters, but there are many exceptions to 

this general trend (U.S. Geol. Survey, 1960), especially in streams 

where large amounts of manganese are brought in by disposal of 

industrial waste or by mine drainage.

The principal importance of processes involving organic materials 

in the occurrence of manganese in natural water is probably in connection with the reduction of manganese oxide. Absorption of 

manganese from soil by growing plants is no doubt a major factor in 

making the element available for solution in runoff. Tv e reduction of 

oxide by organic matter in soil moisture or in the bottom mud of lakes 

is probably also a common process. It is not difficult to visualize 

mechanisms by which water might be induced to move from organicrich environments through zones containing oxidized manganese, as for 

example in collector wells near streams. The manganese content 

of water pumped from some wells might thus tend to increase with 

time.

The aim of the work which has been described was to present in a 

systematic way a summary of the chemical factors which control the
REDUCTION AND COMPLEXING OF MANGANESE BY GALLIC ACIDS D27

behavior of manganese in relatively pure natural water at and near the 

land surface. The behavior of manganese in hydrothermal environments and in brines, or factors which might bring about accumulation 

of manganese minerals, may be related to some of the topics which 

have been discussed. These topics, however, represent aree.s which 

were not directly investigated.

Most of the theoretical concepts discussed in this series of papers 

were tested in the laboratory. Field investigations of manganese 

behavior in ground water, giving particular attention to mineral 

sources of the element in rocks and soils, and studies of arer.s where 

solution and deposition of manganese occur in beds of rivers ar d creeks 

would be desirable as a further test of the concepts. Colleagues in 

the Geological Survey and others who might be able to do so are urged 

to undertake these studies.

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Bjerrum, Jannik, Schwarzenbach, Gerold and Sillen, L. G., 1957, Stability constants, part 1, Organic ligands: Spec. Pub. no. 6, Chem. Soc. (London), 1957, 

105 p.

Black, A. P. and Christman, R. F., 1963a, Characteristics of colored surface waters: 

Am. Water Works Assoc. Jour., v. 55. p. 753-770.

 1963b, Chemical characteristics of fulvic acids: Am. Water Works Assoc. 

Jour., v. 55, p. 897-912.

Durum, W. H. and Haffty, Joseph, 1963, Implications of the minor element content of some major streams of the world: Geochim. et Cosmochim. Acta, 

v. 27, p. 1-11.

Durum, W. H., Heidel, S. G. and Tison, L. J., 1960, World-wide runoff of dissolved solids: Pub. 51 Internat. Assoc. of Sci. Hydrology, p. 618-628.

Feitknecht, W., Brunner, P. and Oswald, H. R., 162, Uber den Einfluss der 

Feuchtigkeit auf die Oxydation von Manganhydroxid durch molecularen 

Sauerstoff: Zeitschr. anorg. u. allg. Chem., v. 316, p. 154-160.

Frondel, Clifford, 1953, New manganese oxides hydrohausmannite and woodruffite: Am. Mineralogist, v. 38, p. 761-769.

Hem, J. D., 1960, Complexes of ferrous iron with tannic acid: U.S. Geol. Survey 

Water-Supply Paper 1459-D, p. 75-94.

 1963, Chemical equilibria and rates of manganese oxidation: U.S. Geol.

Survey Water-Supply Paper 1667-A, 64 p. 

Hutchinson, G. E., 1957, Treatise on Limnology: New York, John Wiley & Sons,

Geography, Physics and Chemistry, v. 1, 1015 p. 

Ingols, R. S. and Wilroy, R. D., 1963, Mechanism of manganese solution in lake

waters: Am. Water Works Assoc. Jour., v. 55, p. 282-290.

Kachhawaha, M. S. and Bhattacharya, A. K., 1962, Gallate complex of manganese (II): Indian Acad. Sci. Proc., v. 32a, p. 376-382. 

Lamar, W. L. and Goerlitz, D. F., 1963, Characterization of carboxylic acids in

unpolluted streams by gas chromatography: Am. Water Works ASIBOC. Jour.,

v. 55, p. 797-802. 

Lange, N. A., 1956, Handbook of chemistry: Sandusky, Ohio, Handb. Publishers

Inc., 1969 p. 

Lieth, Helmut, 1963, The role of vegetation in the carbon dioxide content of the

atmosphere: Jour, of Geophys. Research, v. 68, p. 3887-3898.