Chemical Equilibria 

and Rates of Manganese 

'Oxidation

r-i JOHN D. HEM

CHEMISTRY OF MANGANESE IN NATURAL WATER

GEOLOGICAL SURVEY WATER-SUPPLY PAPER 1667-A

UNITED STATES GOVERNMENT PRINTING OFFICE, WASHINGTON : 1963
UNITED STATES DEPARTMENT OF THE INTERIOR 

STEWART L. UDALL, Secretary

GEOLOGICAL SURVEY 

Thomas B. Nolan, Director
CONTENTS

[The letters in parentheses designate separately published chapters]

(A) Chemical equilibria and rates of manganese oxidation, by John D. Hem.

(B) Deposition and solution of manganese oxides, by John D. Hem.

(C) Intracellular and extracellular concentration of manganese and oil <?r elements by aquatic organisms, by Eugene T. Oborn.

(D) Reduction and complexing of manganese by gallic acids, by John D. Hem.

o

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 I-intlng Office 

Washington 25, D.C.
CONTENTS

Page

Abstract-_______________________________________________________ Al

Aims and organization of study____.________________________________ 1

Occurrence and chemical properties of manganese____________________ 2

Nature and significance of stability-field diagrams-____________________ 3

Chemical-equilibrium computations._____________________________ 5

Location of stability boundaries-________________________________ 10

Laboratory studies of manganese-oxidation equilibria-_________________ 14

Composition of precipitated oxides_____________________________ 16

Manganese equilibria involving bicarbonate___________________________ 19

Fixed bicarbonate activity_--____-___-_-_-_---_-_-______________ 19

Fixed carbon dioxide partial pressure___________________________ 23

Laboratory studies___________________________________________ 25

Manganese equilibria involving sulfate_______________________________ 26

Laboratory studies___________________________________________ 29

Soluble manganese complexes-- _____________________________________ 30

Manganese bicarbonate complex_______________________________ 31

Laboratory studies_________________________________________ 31

Determination of equilibrium constant-__________________ 35

Calculation of free energy of the complex_________________ 38

Effect of the complex on manganese equilibria_______________ 38

Stability field of bicarbonate complex.___________________ 39

Manganese sulfate ion pair_____________________________________ 43

Comparison of strength of complexes____________________________ 43

Combined effects of bicarbonate and sulfate complexes on manganepe

equilibria. _ _________________________________________________ 46

Other inorganic complexes______________________________________ 49

Organic complexes____________________________________________ 51

Rates of chemical reaction..,_______________________________________ 52

Kinetics of mangan ese oxidation and precipitation _________________ 53

Effect of sulfate____________________________ 54

Effect of bicarbonate_____________________________________ 56

Factors that may increase reaction rate______________________ 57

Summary__ _____________________________________________ 59

Coprecipitation of manganese with ferric hydroxie_____________________ 59

Coprecipitation of manganese and other elements_____________________ 61

General conclusions._______________________________________________ 62

Literature cited-__________________________________________________ 62

in
IV CONTENTS

ILLUSTRATIONS

Page 

FIGURES 1-6. Stability fields of manganese species:

1. In the absence of bicarbonate and sulf ate _________ A9

2. In presence of 100 ppm HCOa~ and natural rhodochrosite_ __________________________________ 20

3. In presence of 100 ppm HCOa~ and precipitated

MnCO3 -_-_-_____----_-____________-_- ______ 22

4. At a partial pressure of COz^3.8X10~* atmosphere. 24

5. In presence of 10 ppm HCO3~ and 1,000 ppm SO4 -2 _ 28

6. In presence of 2,000 ppm HCO3 -________________ 40

7. Proportions of manganese complexed at various bicarbonate and sulfate activities ________________ 44

8. Stability fields of manganese species in presence of

2,000 ppm HCO3 - and 2,000 ppm S04 -2 _______ 47

9. Rate of consumption of OH~ during manganese oxidation at different pH levels___________-_____-_ 55

10. Rate of oxidation of manganese as affected by pH

and sulfate_-_______-________________________ 56

11. Rate of oxidation of manganese as affected by pH

and bicarbonate..___________________________ 58

TABLES

Page 

TABLE 1. Free energy of formation (25°C) of crystalline, liquid, and

dissolved substances considered in stability-field diagrams. A7

2. Equilibrium constants utilized in constructing stability-field

diagrams-___________________________________________ 8

3. Oxidation of dissolved Mn+2 _________________ 16

4. Experimental data on manganese bicarbonate con?plex______ 37

5. Coprecipitation of manganese with ferric hydroxi<?e________ 60
CHEMISTRY OF MANGANESE IN NATURAL WATER

CHEMICAL EQUILIBRIA AND RATES OF MANGANESE

OXIDATION

By JOHN D. HEM

ABSTRACT

The relationships between manganese in solution, Eh, pH, and the actrdties of 

bicarbonate and sulfate ions are shown by means of seven stability-field diagrams. 

The behavior of manganese in laboratory experiments is in general agreement 

with predictions of the diagrams. In distilled water, with an Eh near 0.55 volts 

and a pH of 7.0, manganese has a solubility of about 1.0 ppm (parts per million). 

An increase in pH or in Eh tends to lower the solubility.

Bicarbonate species in equilibrium with manganese carbonate may control 

manganese solubility in some systems. Increased bicarbonate activity decreases 

manganese solubility. Relationships between manganese and sulfate activities 

are related to solution of manganese sulfide with oxidation of the sulfur.

Manganese forms soluble complexes with bicarbonate and sulfate. The 

association constant for MnHCOs+ ion was determined to be 63. The solubility 

of manganese in natural water, especially that which contains large amounts of 

bicarbonate or sulfate, is influenced by complexing. In a solution where the total 

dissolved manganese is 1.0 ppm, the activity of sulfate is 250 ppro, and bicarbonate activity is 470 ppm, half the manganese present will be in the complexes.

The rate at which divalent manganese is oxidized and precipitated from aerated 

solutions is sharply increased by increasing the pH. The rate is dirrinished 

when sulfate and bicarbonate ions are present. Experimental data suggest that 

the reaction may be first order with respect to manganese concentration initially, 

but the rate increases as the reaction proceeds, probably because of autocr.talysis.

Manganese coprecipitates with ferric hydroxide when the pH is greater than 6.7.

AIMS AND ORGANIZATION OF STUDY

Although manganese is a minor constituent of natural water, even 

small concentrations of the element constitute an undesirable impurity 

in water to be used for domestic and for many industrial purposes. 

The U.S. Geological Survey is conducting research in the chemical 

behavior of manganese and other hydrosolic metals. These elements 

are readily soluble under some of the chemical conditions common to 

natural water, but under some others, also common to natural water, 

they form oxides and hydroxides of very low solubility.

Al
A2 CHEMISTRY OF MANGANESE IN NATURAL WATER

The research work includes systematic description and evaluation 

of the factors affecting solubility of manganese under conditions of 

chemical equilibrium and study of some factors affecting reaction 

rates. The effects of macrobiota and microbiota are also considered, 

together with factors more directly related to inorganic chemical 

behavior. This report is the first of a series describing the results of 

research on the chemistry of manganese.

The research work described in this report was done in laboratories 

of the U.S. Geological Survey in Denver, Colo. The author was assisted in the laboratory work by M. J. Fishman. Helpful suggestions 

in the course of the study and in review of the manuscript were 

furnished by colleagues of the author in the Geological Survey and 

by Professor R. M. Garrels, Department of Geological Sciences, 

Harvard University.

OCCURRENCE AND CHEMICAL PROPEPTIE8 

OF MANGANESE

The element manganese is number 25 in the periodic table. It is 

classified as a transition element and displays some chemical characteristics, such as multiple valency, that are typical of that group of 

elements. Only one isotope of the element, Mn55 , exists naturally, 

although nuclides of higher and lower mass number, all radioactive, 

have been produced artificially. Manganese is located adjacent to 

iron in the periodic table, and the two elements have some chemical 

resemblances. The two elements are often considered together, particularly in the technology of water supply and treatment, because 

they cause similar problems and both may need to be removed. Because the chemistry of manganese is different from that of iron in 

several important respects, manganese is much more difficult to remove 

from water than iron.

It is generally agreed that the amount of manganese present in the 

rocks of the earth's crust averages about 0.10 percent. This contrasts 

with a concentration of nearly 5.0 percent for iron. T' x us, the abundance data suggest that manganese should be much less common in 

solution in natural water than is iron, and, in general, tHs expectation 

is realized. In many places the rarity or absence of manganesecontaining minerals may prevent saturation of the natural water. In 

such places the solubility and rates of solution of the minerals are not 

the only controls on the concentration of manganese in the water. 

However, manganese-bearing minerals are common in sedimentary 

rocks and soils. Manganese in igneous rocks commorly occurs as a 

minor component in many minerals, such as the pyrorenes, olivines, 

and amphiboles, but independent manganese minerals are scarce.

Combined or ionic manganese can occur at oxidation levels from 

+ 1 to +7. In natural environments, however, some of these oxida-
EQUILIBRIA AND OXIDATION RATE A3

tion states are not likely to occur. In the presence of water, under 

conditions to be expected at or near the surface of the earth, only the 

+2, +3, +4, and +6 states are stable. Species of manganese in the 

+ 7 state (permanganate, MnO4~) are commonly utilized in tl e laboratory, and solutions of permanganate have sufficient stability to be 

used in oxidation-reduction titrations. The oxidizing power of permanganate, however, is great enough to release oxygen froir water, 

although ordinarily the reaction is very slow.

The principal common minerals of manganese are the oxides and 

hydroxy-oxides, such as the minerals pyrolusite (MnO2) and psilomelane 

[BaMn+2Mn+48Oi6 (OH) 4 ]. Mixtures of pyrolusite with manganite 

[MnO(OH)], braunite [(Mn,Si) 2O3], hausmanite [Mn3O4], and psilomelane are common in nearsurface environments. Mixturer whose 

exact mineral composition may not be known are termed 'Vad," a 

generic name for earthy mixtures. The adsorptive capacity of manganese oxy-hydroxide precipitates has been credited for the complex 

assemblage of impurities often found in earthy masses of manganese 

minerals. Less common minerals are the more reduced pyrochroite 

[Mn(OH) 2] and manganosite (MnO).

As is common for the transition elements, some of the compCTinds of 

manganese listed above appear to be mixtures of two formal oridation 

states or valences.

Rhodochrosite, the manganese carbonate, occurs in some ore deposits 

and is an impurity in some carbonate rocks. The sulfide alabandite 

(MnS) is relatively scarce.

Data on the speed of oxidation-reduction reactions involving the 

+2, +3, and +4 forms of manganese are very sparse, and nothing 

useful was found in published literature. The present study shows 

that these reactions are slower than similar ones involving iron, which 

Stumm and Lee (1961) found were generally rapid at the pH levels of 

most natural water. Nevertheless, the oxidation of Mn+2 is rapid 

enough under favorable conditions, especially when increased by the 

catalytic action of micro-organisms, to reach equilibrium in natural 

systems. Therefore, an equilibrium treatment of manganese systems 

likely to be found in natural water is logical, particularly for the conditions relating to solution and solubility, and has been chosen as a 

logical approach here. However, it should be recognized that the 

representation of solid phases has been simplified to the point where 

that part of the treatment may not approach natural conditions 

closely.

NATURE AND SIGNIFICANCE OF STABILITY-FIELD

DIAGRAMS

The stability-field, or Eh-pH, diagram is a graphical means of 

showing the conditions of redox potential and hydrogen-ion cor centra-
A4 CHEMISTRY OF MANGANESE IN NATURAL WATER

tion that are compatible with specified dissolved ions and solid species 

of an element at chemical equilibrium in the presence of water. The 

Eh-pH graph apparently was first utilized in Europe by Pourbaix 

(1949). In its original form, the diagram was used principally in 

studies of the factors involved in corrosion of metals. Stability-field 

diagrams, however, have proved to be useful for representing in clear 

and concise form the important features of the chemical behavior of 

many of the elements under the conditions that prevail in natural 

systems. The literature contains numerous examples of the use of 

these diagrams in studies of the origin of ore deposits1 and in other 

geologic applications, and the diagrams have been used also in studies 

relating to oceanography and limnology. Garrels (1960) has described 

methods of preparing and using Eh-pH diagrams. The principal 

features of the chemistry of iron have been shown by such diagrams 

(Hem and Cropper, 1959); and because manganese is strongly affected 

in its behavior in natural environments, both by oxidation and reduction and by hydrogen-ion concentration, stability-field diagrams have 

been used in this discussion.

In preparing the diagrams, certain basic assumptions are made and 

should be kept in mind when making interpretations. The fundamental equations used are those dealing with a state of chemical 

equilibrium among the dissolved and solid-phase species. If all the 

chemical reactions involved are relatively rapid and thus reach equilibrium 

quickly, no uncertainty is introduced by assuming th°s existence of 

equilibrium. If reactions are slow, forms of an element may exist 

under conditions that are chemically unfavorable for long periods of 

time. Where very long time intervals are involved, as in some 

geologic systems, a slow reaction may tend to reach equilibrium, but 

the reaction time in most natural-water systems near the earth's 

surface is short in the geologic sense.

However, even though it is uncertain that all the manganese 

species considered could be logically expected to reach chemical 

equilibrium in natural water, the stability-field diagram is still useful 

to indicate limiting conditions. Laboratory experiments described 

in this report and field observations can be used as aidr to determine 

the practical usefulness of theoretical calculations.

The diagrams are prepared for a pressure of 1 atmosphere and a 

temperature of 25 °C and are strictly applicable only under these 

conditions, although pressures of a few atmospheres and temperatures 

of 10° or 15° above or below 25°C do not introduce serious error. 

The natural systems that are of interest in the study of the chemistry 

of manganese in natural water generally will involve more complex 

compounds than are dealt with here. The aqueous pha^e will usually
EQUILIBRIA AND OXIDATION RATE A5

contain other constituents in addition to the manganese species. 

The systems are affected to a lesser degree by departures from 

standard conditions of temperature and pressure. In general, the 

effects of minor deviations in composition of the substances involved 

in natural systems from the composition assumed in the diagrams 

cannot be evaluated. Standard computations of chemical thermodynamics can provide results applicable to any temperature or 

pressure that is likely to be found.

The manganese concentrations indicated in that part of the 

stability-field diagram where dissolved species predominate are the 

effective, or thermodynamic, concentrations (activities) rather than 

the concentrations that would be measured by chemical anslysis. In 

extremely dilute solutions, these thermodynamic concentrations are 

almost equal to the actual concentrations; but as the amount of 

dissolved material increases, the freedom of charged ions to move in 

the solution is decreased. As a result the effectiveness of a given 

concentration of ions to participate in reactions decreases as total 

concentration increases.

The measured concentrations reported in chemical analysis may be 

converted to thermodynamic concentrations by means of a computed 

activity coefficient. Techniques for performing this computation are 

given in the literature (Klotz, 1950), and adaptations for natural 

waters were described by the writer (Hem, 1961). For an ion with 

a dual charge, such as Mn+2, the activity will be 85 to 90 percent of 

the measured concentration when the dissolved-solids concertration 

of the water in which the ion occurs is about 50 ppm. In water 

containing from 3,000 to 5,000 ppm (parts per million) of dissolved 

solids, the activities may be less than half of the measured concentrations.

When using stability-field diagrams to interpret water analyses, the 

analytical data must be converted to activities. For water that is 

very highly mineralized, accurate conversion to activities may not be 

possible.

CHEMICAL-EQUILIBRIUM COMPUTATIONS

The stability-field diagram utilizes the following basic relationships 

of chemical thermodynamics:

(1)

(3) nl "red. x

668870 63 2
A6 CHEMISTRY OF MANGANESE IN NATURAL WATER

Symbols used in the equations are:

A.F°= standard free-energy change in chemical reaction, in 

kilogram-calories

-R=universal gas constant (1.987 calories per degree mole or

8.314 joules per degree mole) 

T= temperature, in degrees Kelvin 

£n=logarithm to base "e"

n number of electrons appearing in the redox equation 

/=Faraday constant (23,060 calories per volt or 96,484

coulombs per gram equivalent)

-E°=standard potential, in volts 

Eh= redox potential

 ^_= activities of oxidized and reduced species as expressed in mass 

a red. action equations.

If the system is more oxidized than the hydrogen half-cell (partial 

pressure of H2 <O atmosphere), the potential of the 1 alf-cell will be 

taken as positive. This is the conventional means of representing 

such values in geochemical literature. The chemical reactions that 

the equations are used to evaluate are written here PS reductions 

that is, the representation of electrons in the equatior is on the lefthand side.

The sign conventions used in Eh calculations are chosen arbitrarily. 

If chemical reactions are written as oxidations and if systems more 

oxidized than the hydrogen half-cell are given a positive sign, the negative sign in equation (2) must be changed to positive.

Table 1 contains published values for the standard free energies of 

formation of manganese species and other species of importance in 

this discussion. The most recent compilation of such data for manganese species is that of Mah (1960). Supplementary data are taken 

mostly from the compilation of Latimer (1952). Several values in 

the table were calculated from information published in other sources, 

such as the compilation of stability constants of Bjerrvm, Schwarzenbach, and Sillen (1958). Free-energy values may be obtained from 

data of this kind from the appropriate chemical equation and the 

relationships

and

AF°=-^AF° products 2AF° reactants

when the standard free energies of formation of all the products and 

reactants, except one, are known. Some writers refer tc the net change 

in standard free energy occurring in a chemical reaction as "standard 

reaction potential".
EQUILIBRIA AND OXIDATION RATE A7

TABLE 1. Free energy of formation (25° C) of gaseous (g), crystalline (c), liquid 

(I), and dissolved (aq) substances considered in stability-field diagrams

Species

Mn+2 (aq). ______________

MnOH+ (aq)___-_________

Mn(OH) 2 (c)____. ________

HMnO2- (aq)__---_---___

MnCO3 (c) (rhodochrosite) _ MnCOs (c) (precipitated) __ 

MnS (c) _ _-______

MnO (c)___ __ ___ _ ___

MnSiO3 (c)___ _ ________

Mn2Si04 (c)__-___________

MnHCO3+ (aq).___ ______ .

MnS04 (aq) ___ ____.___.

Mn3O4 (c)__-______--_-__

Mn+3 (aq)___ _ ______ __

Mn(OH)» (c)___ __________

Mn2 O3 (c)_______________

MnO2 (c)_-__--___-_-___-

MnOr2 (aq)__-_-_-_-_-_-

H2 0 (1)________._________

OH- (aq) ----_-_--_---

H2CO3 (aq)___---_.___-__

HCO3- (aq) __.___ __

COa-' (aq) 

H2S(aq)-____-_-______-__

HS- (aq)_ ----__ -__

HSOr (aq)_-_-_-___-___-

S04 -2 (aq)---------------

C02 (g)-_ ....._.___.___.

S(c)_________--____--__.

AJ?° (kg-cal)

-54.4

1 -96. 6

- 146. 9

-120.9

-195. 4 

-194.7 

-50. 55

-86. 75

-289. 0

-381. 95

- 197. 1

-234. 8

-305. 85

-19.6

-181.

-209. 85

-111.3

-120.4

-56. 69

-37. 595

-149.00

-140. 31

-126. 22

-6. 54

3.01

- 179. 94

-177.34

-94. 26

0.0

Source of data

Latimer (1952, p. 235).

(1952). 

Latimer (1952, p. 235).

Do.

Mah (1960), Latimer (1952, p. 235). 

Garrels, Thompson, and Sieve" (1960). 

Mah (1960).

Do.

Do.

Do.

bach, and Silicon (1958, p. 82). 

Mah (1960).

Latimer (1952, p. 235).

Do.

Mah (1960).

Do.

Latimer (1952, p. 235).

Latimer (1952, p. 39).

Do.

Latimer (1952, p. 128).

Do.

Do.

Latimer (1952, p. 72).

Do.

Do.

Do.

Latimer (1952, p. 128).

Latimer (1952, p. 72).

i Determined at 30 °O.

The standard free energy of formation of the complex ion MnHCO3+ 

was determined experimentally in this research study.

Table 2 is a compilation of some of the equilibrium constants that 

are utilized extensively in this report. Values shown in table 2 as 

"calculated" were computed from free-energy values in table 1. 

The values ascribed to other authors were taken directly from their 

quoted values. Probably, because of rounding off, some of these 

differ slightly from the results obtained by calculation from the free 

energies in table 1.

The distribution of manganese species in stability-field diagrams in 

this report is given only for conditions under which water is stable. 

The redox potential at which water is oxidized with the liberation 

of gaseous oxygen is 1.23 v (volts) at a pH of 0. At a pH of 14, 

the Eh for this reaction is just over 0.4 v. Water may be reduced 

to liberate gaseous hydrogen at Eh values of less than 0 v f,t a pH 

of 0 and at Eh values of less than 0.82 v at a pH of 14. Although 

it is usually assumed that pH-Eh conditions beyond the water-
A8 CHEMISTRY OF MANGANESE IN NATURAL WATER 

TABLE 2. Equilibrium constants utilized in constructing stability-field diagrams

Reaction Equilibrium 

constant

Source of data

Mn(OH), 

Mn(OHh (c)^B

4.9X104... 

4.2X101"..

8.7X10-»»- 

63.

MnOOj (c) (rhodochrosite) +2HjO^Mn(OH)a (c)+2H+

MnCOi c (ppt)+2HjO^Mn(OH)i

1.9X10*.. ...

4.14X10' . 

4.16X10-7- 

4.84X10-"-. 

1.1X10-7 

1.26X10-*.. 

7.4X10-"-.

2.3X10-J8-- 

4.2X104....

Oalcu'ated.

OhabRrek, Courtney, and 

Martell (1952).

Oalcu'ated.

Detemined in this study.

Jamer, J. O. (1947) in 

Bjerrum, Schwarzenbach, 

and Sillta (1958) p. 82.

CalcuHted.

Latiirnr (1952, p. 135). D-.

Laturer (1952, p. 71).

Laturer (1952, p. 74).

Calculated.

stability boundaries are not likely to occur in nature, p.n Eh considerably above or below these limits is necessary in order to cause water 

to decompose at a readily observable rate. Consequently, the area 

outside the water-stability boundaries may have some practical significance, especially in laboratory work. For present purposes, 

however, only the area within which water is stable wiT be considered.

Figure 1 shows the fields of stability for 4 solid phases of manganese and tW fields of dominance of 3 dissolved species of manganese. 

Interfering anions and cations are absent. This is about the simplest 

representation that can be made for a manganese system that includes 

enough variables to approach actual conditions in laboratory solutions. The four solids having fields of stability in the figure are 

MnO2 , Mn2O3, Mn3O4 , and Mn(OH)2 . The symbol (c) on the diagram indicates crystalline (solid) material. Free-energy values given 

in table 1 show that MnO would be changed to Mn(OH)2 and that 

Mn(OH) 3 would dehydrate even in contact with water; consequently, 

such forms do not appear on the diagram. The compcrmd Mn(OH) 4, 

has been mentioned in the literature but proof of its existence is 

lacking.

The solid lines represent the positions of the boundaries of the 

stability fields of solids, when the effective concentration of dissolved 

manganese is 0.01 ppm (1.8X10"7 molal). The dissolved species 

considered are Mn+2, MnOH+, and the anionic form HMnO2~. All 

these are varieties of divalent manganese. Within the stability limits of water, Mn+3 (aq) is less abundant than Mn+2 F,nd is not considered. The line along which Mn+3 =Mn+2 is above 1.23 v at a pH 

of 0. Also, with one minor exception, the highly oxidized anionic 

species are stable only outside this area. Therefore Mn47 (as in 

permanganate) is not considered. Near a pH of 14 and an Eh of 

+0.40, the stability boundary between the manganate ion, in which
EQUILIBRIA AND OXIDATION RATE A9

1.20

1.00

0.80

o 0.60

UJ

0.40

f= 0.20

z

UJ

I2 o.oo

0.20

x

-

-0.40

-0.60

-0.80

Water oxidized

6 8 

PH

10 12

FIGTJRE 1. Stability fields of manganese species in aqueous solution, free of bicarbonate and sulfate ions. 

Total dissolved-manganese activity ranges from 0.01 to 100 ppm.

the manganese is hexavalent, and MnO2 (c) intersects the waterstability boundary. However, the area of stability of manganate 

is very small and lies in a region that never exists in nature. Hence, 

the manganate species have been omitted from the diagrams. The 

area in which the aqueous hydroxide complex is the dominant dissolved species is separated from the Mn+2 (aq) region by a dotted 

line. Dashed lines represent the positions of the stability-field boundaries of the solids when dissolved Mn+2 activity is 0.10, 1.0, 10,
A10 CHEMISTRY OF MANGANESE IN NATURAL WATER

and 100 ppm. These dashed lines are, in effect, solubility "contours" 

for manganese in the system described as functions of pH and Eh.

Activities of dissolved manganese and of anions are given in parts 

per million, because this unit is the most widely used in water chemistry. However, the calculations require that reactant activities be 

expressed in moles. Molar concentrations (moles per liter) are related to molal concentrations (moles per kilogram of solution) in 

the same way as milligrams per liter are related to prrts per million. 

In dilute solutions, such as the ones considered here, the difference 

between molal and molar quantities is insignificant.

LOCATION OF STABILITY BOUNDARIES

The chemical concepts and calculations used in c'etermining the 

positions of the boundaries in figure 1 are similar to the procedures 

described by Garrels (1960) and previously used by the writer (Hem 

and Cropper, 1959, Hem, 1960a) in making similar diagrams for iron. 

The steps in preparing figure 1 are as follows: Boundaries for the 

field of stability for water are drawn first. Within the water-stability 

area, species whose stability fields are separated by vertical boundaries 

(no reduction or oxidation involved) are those of Mn+2 and represent 

a simple starting point.

The only solid Mn+2 variety is Mn(OH) 2 . Thr hydroxide is 

amphoteric. As pH decreases, the equilibrium

Mn(OH) 2 (c)+H+^±MnOH++H2O

shifts to the right. The equilibrium constant (table 2) is 4.9X10*. 

We may write

[MnOH+][H20] [Mn(OH) 2 ][H+r4' yX1U -

The activity terms, which are expressed in moles per l: ter, or in moles 

per kilogram of dilute solutions, are those enclosed in brackets. If 

the activity of dissolved manganese, here almost all in the MnOH+ 

form, is taken as 1.8X10"7 molal (0.01 ppm) and the activities of 

water and of the solid Mn(OH) 2 are taken as unity (by definition), 

the equation may be solved for [H+].

~ 12

4.9X104

The negative log (pH) of this value is 11.4, henco the boundary 

between Mn(OH) 2 (c) and MnOH+ is drawn vertically at this pH 

value.
EQUILIBRIA AND OXIDATION RATE All

At higher pH the equilibrium

Mn(OH) 2 (c)^HMnO2 -+H+

increases in importance. The equilibrium constant for this reaction 

(table 2) is 8.7XKT20. This leads to the equation

[HMnOr] [H+] 20 [Mn(OH)2] 8'7X1° '

Here again the activity of Mn(OH) 2 (c) is taken as unity, and the 

activity of the manganite ion as 1.8X10"7 molal. Thus

Q 7 vi n-2o

rH+i= =4 8xio-13 I-I--I- I - r\ ^ s 4 r\ _ 7 ^I«O /\ J. \J «

1.8X10 7

and the equivalent pH is 12.3. This represents the high pH boundary 

of the Mn(OH) 2 (c) stability field.

The MnOH+ complex ion is further dissociated into Mn+2 and OH" 

as the pH decreases. The dissociation may be written

MnOH+-fH+^Mn+2 -fH2O

for which the equilibrium constant is 4.2X1010 (table 2). This leads 

to the relationship

[Mn+2][H2 0] [MnOH+l[H+]~~4^X1U '

The activity of water may again be assigned a value of 1. At the 

boundary between the Mn+2 and the MnOH+ fields, the two species 

will be at equal activity. The equation then simplifies to

which is equivalent to pH 10.6. It should be noted that this boundary 

is not dependent on manganese concentration. To prevent it from 

being confused with the solid lines separating solid from saueous 

species in figure 1, the boundary is represented by a dotted line.

The other field boundaries on figure 1 are influenced by both pH 

and Eb. The solids considered have already been enumerated. The 

equilibrium between manganous hydroxide and the mixec1 oxide 

Mn3O4 is (written as a reduction):

Mn3O4 (c)+2e+2H+-f 2H2O^3Mn(OH) 2 (c).
A12 CHEMISTRY OF MANGANESE IN NATURAL WATER

The mixed oxide can be considered as containing 1 unit of Mn+2 and 

2 of Mn+3. The two Mn+3 atoms are reduced to the Mn+2 state, 

using two unit negative charges, as represented by "e" in the equilibrium. The net standard free-energy change in thir equilibrium, 

&F° net, can be determined from data in table 1 :

2 o (1) 

in which the standard free energy of H+ is zero.

AF°net=3(-146.9)-(-305.85)-2(-56.69) 

= -440.7+419.23 = -21.47 kg-cal.

The standard redox potential E° for the equilibrium can now be 

determined from the relation

21A7

nj 2(23.06) 

E°= 0.466 v.

The value for the oxidation potential of the system wten the participating substances are not at unit activity can be computed from the 

Nernst equation. For the above reaction, the Nernst equation with 

numerical values entered for constants and using base 10 logs takes 

the form

W, Q 1flfl , 0-0592 [Mn3Q4] [H2Q]2 [H+]£ Eh=0.466+-^- log [Mn(OH)J8 

The activities of solids and of water are again unity, and the equation 

may then be simplified to

Eh=0.466-| '-= log [H]2 , or 

Eh=0.466-0.0592 pH.

The vertical boundaries already drawn in figure 1 determine the 

lateral extent of the Mn(OH)2 field. At a pH of 12.3, the Eh becomes 

 0.26 v, and at a pH of 11.4 the Eh value is 0.22 v. A line connecting these points is then drawn to define the bourdary between 

Mn(OH)2 (c) and Mn304 (c).

Similar calculations are based on the equilibria 3lln2O3 (c)+2e 

+2H+;=±2Mn3O4 (c)+H20, where #°=0.842 v, and 2Mn02 (c)+2e 

+2H+^±Mn2O3 (c)+H20, where #°=0.953 v, to establ^h the boundary between Mn304 (c) and Mn2O3 (c) and between Mn2O3 (c) and 

MnO2 (c).
EQUILIBRIA AND OXIDATION RATE A13

Boundaries between dissolved and solid species involve the following 

equilibria:

Mn3O4 (c)+2e+5H+^3MnOH++H2O, #°=0.881 v 

Mn3O4 (c)+2e+2H2O^±3HMnO2-+H+, E°=- 1.226 v 

Mn3O4 (c)+2e+8H+^3Mn+2+4H2O, E°= 1.824 v 

Mn2O3 (c)+2e+6H+^2Mn+2+3H2O, £°=1.496 v

and

MnO2 (c)+2e+4H+^Mn+2+2H2O, E°= 1.225 v.

Values for the standard potentials are computed from net standard 

free-energy change. The value for Eh for the first of these equilibria 

then is

Eh=0.881+«og 5 [MnOH+]3

The activity of dissolved Mn species along the boundary is 1.8X10"7 

molal. Hence,

Eh=0.881+0.0888(6.74)-0.148 pH.

A similar procedure is followed for each of the equilibria.

If the activities of dissolved manganese species are increased, the 

stability fields of solids become larger. The dashed lines in figure 1 

represent the positions of the solid-aqueous species boundaries when 

dissolved manganese activity is increased successively to 0.10, 1.0, 

10, and 100 ppm. These dashed lines were plotted from calculations 

similar to those made for the 0.01 ppm boundaries, but the larger 

values for manganese activity were used. Boundaries between solids 

are not affected by the concentrations of dissolved manganese.

In general, no more than one form of dissolved manganese needs 

to be considered in the fields of dominance portrayed. In figure 1, 

however, near the MnOH+-Mn+2 boundary, a more exact position of 

the dashed line representing 0.10 ppm total dissolved manganese can 

be made if presence of both Mn+2 and MnOH+ is assumed. The 

equilibria that need to be considered in calculating H+ for this dissolved manganese activity are

Mn(OH) 2 (c)+H+^MnOH++H2O and MnOH++H+^Mn+2+H2O.

Equilibrium constants from table 2, assuming water and solids to 

have unit activity, lead to the equations

[MnOH+] 1X1Q2 and [Mn+*] _

[H+] -5axiu and [MnOH+][H+]~4-2X1U

668870 63 3
A14 CHEMISTRY OF MANGANESE IN NATURAL WATER

Also, because the total manganese in solution is fixed at 0.10 ppm. 

[Mn+2]+[MnOH+] = 1.8X10- 6 (or 1.8X1Q-6 molal).

These three equations can be solved simultaneously to give a value 

of 1.9X10-"for[H+].

Calculations for activity of dissolved manganese=0.01 ppm in the 

field of dominance of MnOH+ can be made by assuming that Mn+2 

is virtually absent. When total manganese activity in solution is 

1.0 ppm or more, the MnOH+ activity is too small to affect the total 

appreciably.

Comparison of figure 1 with a similar diagram for a system involving 

water and ferric and ferrous hydroxides (Hem and Cropper, 1959) 

indicates some interesting points of difference. If the pH of a system 

is held near 7.0 and the redox potential is increased, the activity of 

dissolved iron is decreased to less than 0.01 ppm when the Eh exceeds 

0.22 v. To decrease the solubility of manganese to thv same extent, 

however, requires an Eh of nearly 0.60 v at a pH of 7.0. Literature 

on water conditioning and practical experience of water-plant operators in the removal of iron and manganese have shown that removal 

of manganese by aeration is difficult to accomplish and, in any event, 

requires a high pH. In some types of water, the results of aeration 

and removal of iron will include a decrease of pH, which could halt 

the precipitation of iron and bring the system to equilibrium at a point 

where troublesome quantities of iron and probably all the original 

manganese remain in solution.

In water treatment it is essential to maintain conditions that are 

favorable for rapid-reaction rates. The effect of pH or rates of manganese oxidation will be considered later in this discussion.

LABORATORY STUDIES OF MANGANESE-OXIDATION

EQUILIBRIA

The system presented in figure 1 is only approximately applicable 

to natural water, because other ionic and solid specie? besides those 

considered are likely to be present in any natural system. However, 

the system can be duplicated fairly well in the laboratory and, if the 

reactions involved are rapid enough, laboratory measurements of 

Eh, pH, and manganese content should agree with values that would 

be estimated from figure 1.

A series of six experiments (table 3) was performed to ascertain 

whether dilute solutions of manganous ions followed the behavior 

predicted by the theoretical relationships shown in figure 1. A stock 

solution of manganous ions was prepared by dissolving electrolytically 

refined metallic manganese in hydrochloric acid and by then driving
EQUILIBRIA AND OXIDATION RATE A15

off the excess acid by evaporating the solution to dryness. Portions 

of the resulting manganous chloride stock solution were diluted to a 

manganous-ion concentration of 10 ppm. The portions wer?. then 

raised to several selected pH values, ranging from 7.0 to 9.5, by adding 

sodium hydroxide. The solutions were then allowed to remain in 

contact with air, and samples were taken from time to time for 

determination of dissolved manganese by the standard penodate 

oxidation procedure. Before the manganese concentration vas determined, the pH and potential of a bright platinum electrode, 

assumed to be the Eh of the sample, were measured, and the aliquot 

was filtered through a plastic-membrane filter whose pore diameter 

was 0.45 ju. Results are given in table 3.

All these solutions tended to reach a constant pH near 7.0 after 

standing for a day or more. All the manganese was retained in solution except where the initial pH was 9.0 or greater. In most of the 

solutions the measured potential had a steady upward trend, ard after 

5 days it was between 0.54 and 0.56 v. The Mn+2 in theee solutions was oxidized to Mn+3 or Mn"1"4, in part, and these forms precipitated as Mn2O3 or MnO2 , thereby lowering the pH. When the pH 

became unfavorable for further oxidation, the reactions stopped.

In solution 2, measured and computed potentials remained close 

together for the entire series of observations (table 3). In solutions 

3 to 6, a decided lack of equilibrium was observed at the start of the 

test. After 30 minutes, equilibrium had returned to solution 3, 

and after a day to solutions 4 and 5. Solution 4, however, v^as not 

in equilibrium at the end of 5 days, but solution 6 was. When 

measured and computed Eh values agree to ± 0.05 v or less, equilibrium is assumed to have been reached. Previous experience bar shown 

that laboratory measurements of potentials in systems of the kind 

considered here cannot be reproduced much more closely thr.n this. 

In most of the solutions at equilibrium, the Eh values agreed within 

±0.05 v.

Figure 1 suggests that at a pH of 7.0 manganese should not precipitate from these solutions unless the Eh exceeds 0.47 but ttat the 

solutions brought to a higher pH should lose manganese at lower Eh 

values. The hydrogen ions produced in the oxidation and precipitation neutralized the solutions and returned them to a pH of 7. Where 

oxidation had taken place and some manganese precipitated, one 

would expect a condition of equilibrium to occur after a time, as the 

system adjusted itself to the new pH and Mn+2 levels. This condition is shown by solutions 5 and 6 after 1 day and 5 days, respectively. 

The solutions that were nearer neutrality originally had measured 

potentials that are near the level generally expected in aerated water, 

when measured with the bright platinum electrode. To the extent
A16 CHEMISTRY OF MANGANESE IN NATURAL WATER

that measurements with this electrode system are valid indications 

of actual redox potentials, the data in table 3 indicate that the equilibrium approach to the study of the chemistry of manganese is a 

reasonably valid one.

TABLE 3. Oxidation of dissolved Mn+2

Solution

1. __--.__-.-- __ -

2_____-_______-_-

3-_-._-_. --------

4

5.--. _ ___-__._-

6.___. ._ __ 

Time (hours)

0 

.5 

24 

120

0 

. 5 

24 

120

0 

.5 

24 

120

0 

. 5 

24 

120

0 

. 5 

24 

120

0 

. 5 

24 

120

Manganese concentration (ppm)

10 

10 

9. 9 

9.9

10 

10 

9. 9 

9.8

10 

9. 8 

9.9 

9.8

10 

9.5 

9. 9 

9.9

10 

8.3 

8.3 

8.0

10 

.95 

.85 

. 92

pH

7.0 

6.8 

6. 8 

7.0

7.5 

7.3 

6. 9 

6.9

8. 0 

7.5 

6. 7 

6.8

8.5 

7.7 

7.0 

6. 9

9.0

7.8 

7. 1 

7.0

9. 5 

8.3 

6. 9 

7.0

Eh (volts)

Measured

0.49 

.49 

.58 

. 56

.41 

.45 

.52 

. 54

.39 

.43 

. 51 

. 54

.34 

.43 

.48 

. 54

.26 

.43 

.48 

.56

.23 

.42 

.47 

. 54

Estimated

0.48 

. 51 

. 51

.48

.40 

.43 

. 49 

.49

.28 

.40 

.53 

. 51

. 16 

. 35 

.48 

. 49

. 05 

.33 

.46 

.49

.05 

.29 

. 55 

. 54

COMPOSITION OF PRECIPITATED OXIDES

The data in table 3 show that very little manganese was precipitated from manganese solutions whose pH remained b°low 8.5. The 

behavior of solutions, however, varied with pH. An attempt was 

made to determine the composition of the oxides precipitated at 

constant pH.

A series of aliquots containing 10 mg of Mn+2 each were diluted to 

about 100 ml, and the pH was increased rapidly to a predetermined 

value, ranging from 8.5 to 10.0, by adding 0.01 N sodium hydroxide. 

The pH was maintained at the selected value until precipitation 

appeared to be almost complete; then the precipitates were collected 

by filtration in sintered glass crucibles and were dried for 1 hour at 

180° C. This drying time was sufficient to give a constant weight,
EQUILIBRIA ATSTD OXIDATION RATE A17

although it could have brought about further oxidation and perhaps 

dehydration of the precipitate.

The weights of oxide obtained and the apparent percentage of 

manganese in the precipitates for the indicated pH values were as 

follows:

Weight

of

oxide

(milli- Percent 

pH grams') Mn

10.0_____ ________________________________________ 15.5 64.5

9.5-____--____-_-__- __--_-_____-______-____-_-__ 15.1 66.2

9.0-__-------_-----_---__----------_---_-_------ 17.2 58.1

8.5_-_-------_----------------------------_----- 17.5 57.1

The manganese percentages of the pure oxides and hydrox'des for 

which stability fields are shown in figure 1 are:

Oxide Percent Mn

Mn(OH)2 -_ _____________________________________ 61.7

Mn3O4 _ ---------------------------------------- 72.0

Mn.O3 - ---------------------------------------- 69.6

__________-_-__ _---____-____-____-_-_-__ 63.2

Evidently the precipitates obtained in this experiment do not have 

the composition of any of the pure oxides. Further investigation was 

made by preparing a second series of aliquots containing 1C mg of 

Mn+2. The oxides were precipitated as before, using pH vrlues of 

9.5, 9.0, and 8.5. The precipitates were collected on filter paper and 

transferred at once to a titration beaker without allowing the precipitate to become dry. Water was added and the suspension was 

titrated, using a bright platinum electrode and a calomel electrode to 

indicate the end point potentiometrically. The pH during th^ titration was maintained near 8.5. A potassium permanganate solution 

capable of oxidizing 1.66 mg of Mn+3 to Mn"1"4 per ml was usec1 .

The initial potential in the first of the three titrations, that of the 

oxide precipitated at a pH of 9.5, was 0.42 v, which is in the Mn2O3 

stability field. The titration curve had a single inflection point, at 

about 0.60 v, and the results indicated that about 85 percent of the 

manganese oxide was in the Mn+3 state. The remaining 15 percent 

probably was in the Mn*4 state.

In the other two titrations, the initial potential was in the MnO2 

stability field. However, the oxide precipitated at a pH of 9.0 gave 

a fairly well defined inflection point, again at about 0.60 v, which 

indicated that about 30 percent of the oxide was of the Mn+3 variety. 

The oxide that precipitated at a pH of 8.5 apparently contained only
A18 CHEMISTRY OF MANGANESE IN NATURAL WATER

Mn4"4, because the titration curve displayed no inflection point, even 

though the potential rose to 0.74 v.

The results of the titrations are interpreted as showing that no 

Mn+2 was present in the oxides either as Mn(OH)2 or as a part of 

Mn3O4. If an appreciable amount of Mn+2 had been present, the 

titration curves should have had two inflection points, one representing oxidation of Mn+2 to Mn+3 and one representing oxidation of 

Mn+3 to Mn44.

The titration results are not fully conclusive, because some oxidation could have taken place during the process of collecting and filtering the precipitates. The pH value of the water in which two of the 

precipitates were titrated was from 0.5 to 1.0 unit lovrer than the pH 

value at the time the precipitate was formed, and air was not excluded. However, if Mn+2 is absent from the titrated oxides, it 

doubtless is absent also from the dried and weighed oxides.

No combination of MnO2 and Mn2O3 could give a proportion of 

manganese as low as was obtained for the precipitate^1 formed at pH 

8.5 and 9.0. These precipitates must therefore have contained some 

water. The precipitates formed at pH 9.5 and 10.0 could be mixtures 

of Mn2O3 and MnO2 or, as seems more likely, could b^ mostly Mn2O3 

with some water. The apparent presence of water suggests that the 

solid present here was partly Mn(OH)3, which would be 51 percent 

manganese. If the stability of the hydroxide is similar to that of the 

oxide, a combination of Mn2O3 and Mn(OH)3 might o°.cur.

The findings on hydration of the oxides represent the composition 

of material dried at 180° C. Precipitates that are freshly formed, in 

contact with water, may have different compositions. As noted by 

Garrels (1960, p. 114) and by others, ferric hydroxide precipitates 

lose water and decrease in solubility with aging. The behavior of 

hydrated Mn+3 oxides may be similar.

The tendency for manganese to form partly hydrated oxide ores of 

mixed-valence species is well known and agrees with the behavior 

noted here. Because the solids formed are not the pure simple oxides 

that were assumed in preparing figure 1, the laboratory results deviate 

from those predicted from theory. Of particular interest is the tendency for Mn44 oxide to be precipitated in preference to Mn+3 oxide 

at a pH of 8.5 and an evident tendency for the Mn+4 form to be more 

and more predominant at a lower pH.

The behavior of manganese in these experiments is in general 

accord with theoretical calculations used in preparing figure 1. It 

should be noted that the amount of carbonate-specie^ activity introduced by air is too small to affect the calculations significantly.
EQUILIBRIA AND OXIDATION RATE A19

MANGANESE EQUILIBRIA INVOLVING BICARBONATE

The anions present in natural water affect the behavior of manganese in various ways. For example, compounds of low solubility 

may form when manganese combines with carbonate, and complex 

ions formed by manganese and certain anions may increase the 

solubility of manganese. Relationships among dissolved and solid 

species of manganese are likely to be different from those shown in 

figure 1, if bicarbonate species are present in the system, and such 

species are present in practically all natural water becauss of the 

widespread availability of carbon dioxide.

FIXED BICARBONATE ACTIVITY

Figure 2 was prepared from base data in tables 1 and 2. The 

presence of a constant total activity of bicarbonate species equivalent 

to 100 ppm in HCO3~ was assumed. This represents a system that is 

sealed from contact with air or from other sources that coulc1 supplement the amount of carbon dioxide initially present. Many of the 

relationships in figure 2 are the same as in figure 1. However, the 

equilibrium

MnCO3 (c)+H+^Mn+2+HCO3-

affects a considerable part of the area of figure 2. The equilibrium 

constant for this equation depends on the form to which the solid is 

assigned. If the form is rhodochrosite, the constant is 0.31 [A.F0 for 

rhodochrosite from Mah (I960)]. If it is precipitated MnCO3, the 

constant is 1.0, by using the &F° value of Garrels, Thompson, and 

Siever (1960). For figure 2, the form was assumed to be rhodochrosite. The pH at which Mn+2 activity is 1.8X10"7 molal is tt^n computed to be 9.0.

Garrels, Thompson, and Siever attribute the difference between 

A.F0 values for rhodochrosite and for precipitated MnCO3 to impurities 

in the natural rhodochrosite.

Free-energy data show that the boundary between rhodcchrosite 

and Mn(OH) 2 (c) is at a pH of 11.7. Thus, the carbonate field occupies virtually all the MnOH+ field of figure 1 and infringes on the 

Mn(OH)2 (c) area. The boundary between Mn3O4 (c) and MnC03 (c) 

shown in figure 2 was computed from the relationship

Mn304 (c)+2e+3HCO3-+5H+^3MnC03 (c)+4H2O,

where #°=1.87.

The activity of the specific bicarbonate species present in the system 

is a function of pH and can be calculated from the two dissociation 

constants for H2CO3 given in table 2 and the amount specified to be
A20 CHEMISTRY OP MANGANESE IN NATURAL WATER

1.20

12 14

FIGUBE 2. Stability fields of manganese species in aqueous solution free of sulfate. Total dissolved-manganese activity ranges from 0.01 to 100 ppm. Bicarbonate-species activity is 100 ppn as HCOs". Symbo ] 

(c) indicates crystalline form.

present (100 ppm as HC03~). As the activity of b; carbonate decreases with increasing pH above 8.2, the MnC03 (c) Mn304 (c) 

boundary curves somewhat.

The activity of bicarbonate ions must be known to compute the 

positions of the dashed vertical lines in figure 2 that represent the 

MnCO3 (c)-Mn+2 boundary at Mn+2 activities of 0.10, 1.0, 10, and 

100 ppm. An exact treatment would involve considering simultaneously the equilibria
EQUILIBRIA AND OXIDATION RATE A21

MnCO3 (c)+H+^±Mn+2 +HCO3- 

H2CO3

and possibly others. The dissociation constants for carboric acid 

however, show that when pH ranges from 7 to 9 the bicarbonate ion 

predominates over the carbonate ion or the undissociated carbonic 

acid by a factor of at least 10 to 1. In this pH range, the dissociation 

of water is also a minor factor. Consequently, the boundares that 

define manganese activities of 0.01, 0.10, and 1.0 ppm can be located 

with satisfactory accuracy by assuming that all the bicarbonate 

activity is present as HCO3~. These boundaries are at pH 9.0, 8.0 

and 7.0.

Boundaries for activities of 10 and 100 ppm of dissolved Mn+2 must 

be located by considering both HCO3~ and H2CO3 species. A t a pH 

of 6.4, as indicated by the dissociation constant hi table 2, 1 alf the 

total bicarbonate species will be present as HCO3~ and half as undissociated H2CO3 (aq). At a pH of 5.4, the HCO3~ form is only about 

10 percent of the total. As a result of decreased availability of 

bicarbonate ions, the solubility of MnCO3 (c) increases at an increasing rate as the pH becomes less than 6.4 and the spacing between 

the manganese-solubility lines decreases.

If the precipitated form of manganese carbonate rather than the 

natural rhodochrosite is present, the stability field of the solid is 

decreased in size considerably. The boundary between dissolved 

Mn+2 and the solid at 0.01 ppm activity of manganese would be 

shifted to the right a half pH unit to 9.5, and the boundary between 

the solid manganese carbonate and the solid Mn(OH)2 would be at a 

pH of 11.4. The positions of the boundaries for such a system are 

shown in figure 3. Computations are the same as for those for f -*ure 2, 

except that when MnCO3 (c) is involved the free-energy value used is 

the same as that for precipitated material, 194.7 kilogram-calorie 

per mole.

Natural water frequently contains activities of 100 ppm or more of 

bicarbonate. Many carbonate rocks contain some manganese. The 

equilibria involving carbonate are likely, therefore, to have much 

practical importance. The behavior of manganese hi natural water 

is indicated approximately by figures 2 and 3, although in nature the 

systems are more complex than the simple ones represented here. 

However, the effects of carbonates may help to explain why natural 

water seldom contains more than 1 ppm of manganese. Ffgure 1 

indicates manganese solubility to be considerably greater than 1.0 

ppm in most of the Eh-pH area common to natural water for

668870 63 4
A22 CHEMISTRY OF MANGANESE IN NATURAL W^TER

(O

1.20

1.00

0.80

0.60

- 0.40

UJ

j 0.20

<

ty 0.00

g -0.20

O 

UJ

cc

-0.40

-0.60

-0.80

o, »%.

JOO ( *ff^

^>xl£:'%x X

\x$\\

Water oxidized

-Water reduced

6 8 

PH

10 12 14

FIGUBE 3. Stability fields of manganese species in aqueous solution. Conditions are the same as in figure 2 

except that precipitated MnCO jis substituted for rhodochrosite. Symbol (c) indicates crystalline form.

example, from pH 6.0 to 8.5 and from Eh +0.45 to 0.20. In a 

similar area in figures 2 and 3, the solubility exceeds 1.0 ppm only 

below pH 7 or pH 7.5, respectively. A higher bicarbonate activity 

than 100 ppm would, of course, decrease the solubility of manganese 

still further.

In the systems represented in figures 2 and 3, the supply of bicarbonate or related species is fixed at a single and constant value. 

Such a fixed value might be attained in an aquifer that does not 

contain a gas phase or any carbonate minerals. However, any change 

of pH in a natural system containing carbonate rocks will alter the
EQUILIBRIA AND OXIDATION RATE A23

amount of dissolved bicarbonate, so that the model is not entirely 

congruent with all natural conditions.

FIXED CARBON DIOXIDE PARTIAL PRESSURE

Another type of system involving the same solids and ions p.s those 

in figure 3 is one in which a gas phase is in contact with an r.queous 

phase and the partial pressure of carbon dioxide is at a fixec1 value. 

In this system, the amounts of bicarbonate species available in 

solution are a function of the pH and the partial pressure of CO2 

and are related to the form of solid manganese carbonate r,nd the 

amount of dissolved manganese present. Figure 4 represents a 

system exposed to laboratory air. The solid MnCO3 is the freshly 

precipitated variety. The literature indicates that pure air contains 

0.03 percent CO2 by volume. The carbon dioxide content of air 

obtained from the laboratory compressed-air line was found to be 

0.047 percent by volume. The atmospheric pressure is affected by 

the fact that the laboratory is west of Denver, Colo., at an altitude 

of about 5,600 feet above sea level. Therefore, when the barometric 

pressure is in the general fair-weather range, the partial pressure of 

carbon dioxide in the laboratory air is approximately 3.8X10"4 

atmosphere.

Accordingly, the value for pCO2 (partial pressure of CO2) used in 

preparing figure 4 was 3.8X10"4 . To apply this value in the calculations, the fundamental equation is

C02 (g)+H20 (1)^±H2C03 (aq).

A value for the equilibrium constant can be obtained by using freeenergy data in table 1 and the relationship

-RT lnK=L

For this equilibrium, the value of K is 3.7X10"2. Hence, 

[H2CO3 ]=3.7X10-2 XpCO2 ,

and the activity of H2CO3 (aq) available in the solution would be 

1.4X10-*molal.

The chemical equilibria for the first and second dissociation of 

H2CO3 must be satisfied in this system. Hence,

[HCO.-][H+] 

[H2C03 ] "

[HOOT] -4-*4Xlu 
A24 CHEMISTRY OF MANGANESE IN NATURAL WATER

- 0.60 -

FIGURE 4. Stability fields of manganese species in aqueous solution in the pH 

range from 4.0 to 10.0. The solution contains no sulfate. Dissolved-manganese 

activity ranges from 0.01 to 100 ppm at a constant partial pressure of CC "=3.8X10-* 

atmosphere. Symbol (c) indicates crystalline torm.
EQUILIBRIA AND OXIDATION RATE A25

Also, the equilibrium equation for solution of precipitated MnCO3 

applies

[Mn+2 ][HCOr] _

and the ionic product of water is involved

For electrical neutrality, the actual concentrations of negative and 

positive ions present must be equivalent, and we may write

2 C/Mn+2 -|- CH+ 2 G cos ~" ~\~ CHCOS ~r C/ OH~ >

where C represents molar concentration.

For all practical purposes, the activity coefficients may be assumed 

to be unity in the dilute solution under consideration, and activities 

may be substituted for concentrations in the preceding equation.

The treatment can be simplified considerably in locating the 

MnCO3 (c)-Mn+2 boundary, where [Mn+2] = 1.8Xl(r7 . THs is in 

the pH range where the concentrations of CO3~2 and OH~ can be 

neglected. Hence,

_., 4.16X1Q- 7X[H2CO3]^6.0X10- 12^

rnro -i- [H+] - [H+] [MUU3 J-^Mn+ip^xK)-/

Hence 6.0X1Q- 12

1.8X10~7~

[H]+2 =1.08X10-18 

IH+] = 1.04X10-9 

pH=9.0.

As the pH of this system decreases, the solubility of MnCO3 (c) 

increases rapidly. However, as the pH increases, the dissolved 

carbonate in the system also increases, and, as a result, the etability 

field for MnCO3 (c) occupies a large area, which includes tl e fields 

of Mn(OH)2 (c), HMnO2~, and most of Mn3O4 (c). The significance 

of the effects at a high pH is not great in natural water because 

water open to the atmosphere rarely would have a pH greater than 

9.0. The pH range shown in figure 4 is from 4.0 to 10.0.

LABORATORY STUDIES

A series of 10 experimental solutions was prepared to check the 

position of the boundaries in figures 2 and 3. The solutions each
A26 CHEMISTRY OF MANGANESE IN NATURAL WATER

contained 10 ppm Mn+2 and 100 ppm HCO3~. The pF of each solution was adjusted with a potassium hydroxide solution to a predetermined value ranging from 6.5 to 11.0, and the solutions were 

transferred to glass-stoppered bottles. The bottle1* were filled 

completely so that no air space remained. The solutions were allowed 

to stand undisturbed for 13 days, after which the pH and the Eh 

were measured. To remove precipitated material from them, the 

solutions then were filtered through plastic-membrare filters with 

a pore-diameter of 0.45 juOnly the solutions at a pH of 8.0 or less had Eh anc1 pH values in 

the stability region for MnCO3 (c) (precipitated) in figure 3. Some 

light-colored precipitate identified as manganese carbonate was 

filtered out of the solutions held at pH 7.5 and pH 8.0. In the solutions held at higher pH, a brown or black precipitate of oxide was 

obtained. This is to be expected, because the pH and Eh of these 

solutions plot in the oxide-stability region.

Although figure 3 indicates that manganese carbonate solubility 

is less than 10 ppm in terms of activity at pH 7.0, no precipitate 

appeared in the solution at that pH. The activity of Mn+2 in the 

solution was also less than 10 ppm, owing to the amounts of other 

ions present; hence, this solution approached equilibrium closely 

and precipitation should not have been expected.

In an attempt to obtain carbonate precipitates r,t higher pH, 

another set of solutions containing manganese and bicarbonate was 

prepared. To maintain a low redox potential, 4,000 ppm of hydroxylamine hydrochloride was added to each solution. Manganese 

carbonate was precipitated from the solutions having a pH greater 

than 8.0. However, the large amount of hydroxylamine in these 

solutions probably influenced the behavior of the manganese in 

other ways besides altering the Eh. Consequently, the results cannot 

be used to verify the positions of solubility lines on figure 3.

MANGANESE EQUILIBRIA INVOLVING 8TTLFATE

Manganous sulfate is a readily soluble compound. However, if a 

solution containing manganese and sulfate ions were subjected to a 

strongly reducing environment, in which sulfate ions would be 

reduced to sulfide, a precipitate of relatively insoluble manganese 

sulfide could be formed. The behavior of manganese in such systems 

is affected, therefore, by redox reactions involving sulfur.

Equilibria involving sulfur species have been discussed by Garrels 

and Naeser (1958). Sulfur can exist at oxidation state^ ranging from 

 2 to +6; and, because of the covalent character of pulfur bonding, 

some species apparently have fractional oxidation numbers. Thermo-
EQUILIBRIA AND OXIDATION RATE A27

dynamic calculations, however, show that at equilibrium in water at 

25 °C and 1 atmosphere pressure the stable species are S~2, S~*, S°, 

and S+8 and that the situation is not complicated by the intermediate 

oxidation states, such as thiosulfate or sulfite.

Occasionally, some of the partially oxidized species of sulfur are 

reported in published water analyses. Generally the analyses represent thermal springs to which the calculations for 25 °C and 1 atmosphere pressure do not apply. Metastability, however, is characteristic 

of sulfur species in nonthermal environments. Hence, the assumption 

that equilibrium generally exists for sulfur species in natural water 

is less justified than for many other dissolved constituent Some 

reactions involving sulfur are very slow when catalysts are absent. 

The reduction of sulfate, for example, although of fairly common 

occurrence in ground water, does not seem to take place unless 

anaerobic bacteria of a suitable type and a supply of organic food for 

the bacteria are present.

The stability-field diagram is, nevertheless, a valuable indicator 

of limiting conditions for systems that contain sulfur. Figure 5 

represents a system containing activities of 10 ppm dissoVed-COg 

species as bicarbonate and 1,000 ppm dissolved-sulfur species as 

sulfate. Solids represented besides the oxides and hydroxide of 

manganese in figure 1 include the rhodochrosite form of manganese 

carbonate and crystalline manganese sulfide (MnS). The fre0, energy 

of formation of the polysulfide MnS2 corresponding to pyrite is not 

known; hence, this variety was not considered.

Figure 5 shows that an area of stability exists for MnS (c) at a 

high pH and under strongly reducing conditions. However, it is 

essential that the activity of sulfur species be very much greater 

than the activity of bicarbonate species. When the activities of 

both sulfur and bicarbonate species are equal to 100 ppm, the MnS (c) 

field disappears.

Boundaries for the carbonate, oxide, and hydroxide species in 

figure 5 were calculated from equations used for previous illustrations. 

The positions of the boundaries are somewhat altered from those of 

figures 2 and 3 because of the low bicarbonate activity. Boundaries 

between the MnS (c) field and the fields of the other solids were 

determined from two equilibria

Mn(OH)2 (c)+SO4-2 +10H++8e^±MnS (c)+6H20, 

where £r°=0.36 v, and

MnCO3 (c)+SOr2 +8H++8e^MnS (c)+CO3-2+4H2O,
A28 CHEMISTRY OF MANGANESE IN NATURAL W-dTER

«:.

-0.80

0/<> r*'*-.

10 -*' "'% Water oxidized

' ^0.10 %. 0C 

vs*. r\ m °nt *b,^

12 14

FIGURE 5. Stability fields of manganese species in aqueous solution. Total dissolved manganese activity 

ranges from 0.01 to 100 ppm; bicarbonate species activity is 10 ppm as HCOj-; and sulfur species activity 

is 1,000 ppm as SO*-2. Symbol (c) indicates crystalline form.

where #°=0.17 v. As the pH decreases, MnS is disscVed to yield 

Mn+2, as in the equilibrium

MnS 

where K=4.13, or in the equilibrium

MnS (c)+2H+^±Mn+2+H2S (aq), 

where #=4.14X107.
EQUILIBRIA AND OXIDATION RATE A29

At high pH, the sulfide may go into solution to yield im.nganite 

ions, as in the equilibrium

MnS (c)+2H2O^±HMnO2 -+2H++HS-,

where.K:=1.8X10-34 .

Sloping boundaries between MnS (c) and the fields of dominance 

of Mn+2 and HMnO2~ are defined by the equilibria

Mn+2+SOr2 +8H++8e^±MnS (c)+4H2O, 

where E°=0.25 v, and

HMnO2-+SOr2 +llH++8e^MnS (c)+6H2O,

where #°=0.50 v.

At a high pH, the MnS (c) boundary is affected by the relation

MnS (c)+2H2O^±Mn(OH)2 (c)+HS-+H+

where #=2.OX 10~15 .

Hence, when the pH value is greater than about 13.0 and when 

the activity of the sulfur species is 1,000 ppm, the MnS (c) field is 

replaced by the Mn(OH)2 (c) field.

LABORATORY STUDIES

An experiment was performed in the laboratory to evaluate the 

behavior of manganese sulfide. A few hundred milligrams of manganous sulfide was prepared by bubbling H2S gas through manganese 

chloride solution. The prepared material was put into a bottle, 

and 1 liter of de-ionized distilled water was added. The mixture was 

continuously agitated by air bubbled into the bottle from the laboratory compressed-air line. From time to time, samples of the solution 

were withdrawn, and their pH, Eh, and manganese content were 

determined. Results are given as follows:

Dissolved

manganese Eh 

Time (hours) (ppw) pH (volts)

3. ______________________________ 4.2 ____ ____

24________________________________ 8.2 7.4 0.16

48___-----_-____-_--_--_______.__- 10 7.4 .23

72._______________________________ 10 7.2 .08

144____-___-______-___-_-________- 11 7.2 .09

312_______________________,,______ 14 7.2 .36

480_--_-___._______---__-__._.__-- 13 7.0 .39

Although the Eh throughout the experiment was favorr.ble for 

^ulfide oxidation, the amount of manganese dissolved did not increase

668870 63 5
A30 CHEMISTRY OF MANGANESE IN NATURAL TTATER

continuously but seemed to reach a fairly stable value. The fluctuations of Eh suggest that if an equilibrium was reached, it was one that 

did not necessarily involve redox potential.

The simplest explanation of the results seems to be that a two-step 

reaction occurred in which the manganous sulfide war converted to 

carbonate that then partly dissolved. The first reaction that took 

place was oxidation of the manganous sulfide

MnS(c)+4H2O^±Mn+2 4-SOr2 4-8H+4-8e.

However, the air bubbled through the solution contained some 

carbon dioxide that, as it dissolved, tended to precipitate manganese 

as the carbonate

Mn+2 4-H2CO3 (aq)^MnCO3 (c)+2H+.

Thus, manganous sulfide dissolved and manganous carbonate precipitated. The first reaction practically ceased after a protective 

coating of carbonate formed on the sulfide particles. The ultimate 

equilibrium involving only pH, dissolved manganese, and bicarbonate 

activity would be

MnCO3 (c)-f-H+^Mn+2+HC03 -.

If the partial pressure of CO2 in the air remained constant, a constant 

[H+] would be attained. The ionic strength of the final solution is not 

known and temperature was not considered. However, the temperature did not vary greatly and was very near 25 °C throughout the 

experiment. If it is assumed that the concentrations are low enough 

so that the activity coefficients of the dissolved ions are equal to unity, 

the observed pH- and dissolved-Mn+2 values should agr^e closely with 

the manganese-solubility lines in figure 4, which represents a system 

containing MnCO3 (c) (precipitated) exposed to air.

Figure 4 suggests that a solubility of 10 ppm Mn+2 ou^ht to occur at 

a pH of 7.5. The 24- and 48-hour values represent a very close 

approach to these conditions and are evidence that an equilibrium may 

be reached. Later observations in the experiment seem to indicate 

conditions further from equilibrium, but the relatively small deviations may result from changes in composition of the air or contamination during the rather long period that air was continuously bubbled 

through the solution.

SOLUBLE MANGANESE COMPLEXES

The solubility of manganese can be considerably affected by formation of complexes or ion-pairs in solutions where certain anions com-
EQUILIBRIA AND OXIDATION RATE A31

mon in natural water are present. Trie effects of these require consideration in studies of chemistry of manganese.

MANGANESE BICARBONATE COMPLEX

Bjerrum, Schwarzenbach and SillSn (1958) refer to a 1942 publication by Nasanen reporting the existence of a complex cation with the 

formula MnHCO3+. The equilibrium constant for the reaction by 

which this complex ion is formed

Mn+2+HCO3 -^MnHCO3+

is given as 3.0XIO3 . The equilibrium constant was obtained by 

experiments conducted in a rather strong sodium chloride solution. 

The published value for the equilibrium constant implies a strong 

complexing action. Direct application of this value to relatively 

dilute solutions, such as those considered in figure 2, suggests that the 

 complex often is the dominant manganese species present. For 

example, in a solution with a bicarbonate activity of 3.3X10"3 molal, 

the ratio [MnHCO3+]: [Mn+2] will be about 10, as long as the total 

jnanagnese activity remains small compared to the bicarbonate activity. 

An activity of 3.3X10"3 molal of the bicarbonate amounts to a little 

more than 200 ppm and is a value commonly attained in natural water.

LABORATORY STUDIES

Because of the possible importance of the bicarbonate complex 

in the chemistry of manganese in natural water, some laboratory work 

was done to check the equilibrium constant. The value reported in 

the literature (3.0XIO3 ) was obtained under conditions different from 

those of the investigations for this report.

In the first experiment, about 10 g of analytical-grade manganese 

carbonate powder was placed in a glass reaction jar containing 1,000 

ml of de-ionized distilled water. The jar was fitted with a large rubber 

stopper. Through holes in the stopper a glass electrode, a calomel 

electrode, and a thermometer were introduced into the solution. 

Carbon dioxide gas was bubbled into the solution through a glass 

tube reaching to near the bottom of the j ar. A fifth hole in the stopper 

served as a vent for excess carbon dioxide. The electrodes were 

connected to a pH-meter, and the gas flow was started. The pH 

slowly decreased as the carbon dioxide in solution in the water caused 

some of the manganese carbonate to dissolve. After about 4 hours, 

the pH reached a value of 5.37 and remained at that value during an 

additional 12 hours that COo ** was bubbled through othe solution. The

apparent chemical equilibrium of the reaction was tested by making 

periodic determinations of the manganese content of the solution after
A32 CHEMISTRY OF MANGANESE IN NATURAL

filtration through a plastic-membrane filter whose pore diameter 

was 0.45jti. The manganese content appeared to remain constant 

near 80 ppm for the last 8 hours of the experiment.

The reaction vessel was exposed to minor temperature fluctuations in 

the laboratory, but the equilibrium pH was reached r.t a temperature 

between 24° and 25° C. Hence, no temperature correction to standard conditions (25° C) was made. The pressure on the system was 

about 0.85 atmosphere, because the altitude of the laboratory is 5,600 

feet. This departure from standard conditions affects the carbon 

dioxide equilibrium but is not otherwise significant.

In the experiment, the following chemical relationships were at 

equilibrium:

C02 (g)+H2O^H2C03 (aq) (1)

H2CO3 (aq)^H++HCOr (2)

MnCO3 (c)+H+^Mn+2+HCOr (3)

Mn+2 -fHCOr^±MnHCO3+ (4)

The value of K for equilibrium (1), calculated previously for the 

experiment in which air was bubbled into a solution containing MnS 

(c) is 3.7X10~2 . The molal activity of H2CO3 , calculated for the 

pressure at the altitude of the laboratory, is 3.1X10"2.

The activity of bicarbonate can then be calculated for the equilibrium pH of 5.37 from equilibrium (2), using the_enuilibrium constant in table 2

-i 4.16X1Q-7X3.1X1Q-2 3 ]= 4.3X10-8 

=3.0X10~3 molal.

This value can be used with equilibria (3) and (4) to compute the 

activity for Mn+2 and MnHCO3+. The equilibrium constant for 

equilibrium (3), however, must be computed first, ar<d to do so one 

must select a value of the standard free energy of formation of precipitated MnCO3 (c).

A series of experiments that are virtually the same as the one done 

here has been reported by Garrels, Thompson, and Siever (1960). 

The equilibrium pH for their precipitated MnCO3 was 5.35. The 

value of 5.37 obtained here is considered a reasonable check, and the 

value reported by Garrels, Thompson, and Siever for the free energy 

of formation of MnCO3 is evidently nearly applicable to the material 

used here. The equilibrium constant determined from equilibrium 

(3) is 1.0, and the free-energy value is given in table 1.

To complete the solution of equilibria (3) and (4), activities for the 

dissolved manganese species Mn+2 and MnHCO3+ are needed. The
EQUILIBRIA AND OXIDATION RATE A33

analytical value for manganese is a concentration rather than an 

activity and also is a total of all dissolved manganese species. Activity coefficients for the two ions can be computed from concentrations 

with the Debye-Huckel limiting law, using the method of Klotz 

(1950). The basic equation is

, *

log 7=

where

7= activity coefficient sought

z= ionic charge

fj.= ionic strength of solution

A=& constant relating to solvent (0.5085 for water at 25 C0) 

B a constant relating to solvent (0.3281 X 108 for water at 

25 °C)

a = effective diameter of ion of charge z.

Ionic strength is computed from the equation

where mf represents the concentration of each different ion, in moles 

per liter and zt is the charge of that ion.

In the experimental solution, the total concentrations of tl e ionic 

species that contribute to ionic strength are divided among urivalent 

and divalent forms. If no complexing occurred, all the 80-ppmdetermined manganese could be assigned to Mn+2. The ionic strength 

computed from this and from the bicarbonate activity, which probably 

does not differ greatly from the concentration, would be

3.0X10-3 . 1.5X10-3 X4 

"= 2 + 2 

=4.5X10~3 .

If all the manganese were complexed, the ionic strength would be 

2.25 XlO~3 .

Thus the upper and lower limits of the value for fj. can be determined. 

However, the exact value that should be used cannot be determined 

until the proportion of the complexed to uncomplexed species of dissolved manganese is known.

The experimental data provide a means for checking the validity 

of the published value of K, 3.0 XlO3 for equilibrium (4). The
A34 CHEMISTRY OF MANGANESE IN NATURAL VATER

activity of bicarbonate is already available, from equilibria (1) and 

(2), and is 3.0XlO~3. Therefore,

[MnHC°3+U.OX10-3 X3.0X103 =9.

[Mn+2 ]

This ratio of activities is somewhat larger than the ratio of actual 

concentrations because the activity coefficients for the two species 

are not the same. However, for present purposes, the difference is 

not important. Complexing would be extensive and th e ionic strength 

of the solution is therefore about 2.5X10"3 . Substitution of this 

value for n in the Debye-Hiickel equation gives actmty coefficients 

of 0.82 for Mn+2 and 0.95 for HC<V and MnHCO3+. The value of 

a in the Debye-Hiickel equation for the complex ion iir assumed to be 

the same as that given for HCO3~ by Klotz (1950).

The total molal concentration of manganese present, 1.5X10"3 

includes the two ionic species. Therefore,

OMn=<?Mn+2 +£MnHC03+==1.5X10-3 , (5)

where C represents measured molal concentrations. The concentrations of the two species are converted to activities by means of the 

activity coefficients:

-[Mn+2] [MnHC03+]

- + 

From this relationship and from the ratio of the activities of complexed 

to uncomplexed manganese in equilibrium (4) the two activities can 

be evaluated. Simultaneous solution of the two equilibria gives 

1.4X10"4 for [Mn+2]. Substitution of this value in equilibrium (3), 

together with the previously computed bicarbonate activity and the 

equilibrium pH of 5.37 gives an equilibrium constart of 9.9X10"2, 

compared with the value of 1.0 computed from free-erergy data.

Obviously, complexing in the system falls far shor* of the degree 

predicted by the reported constant for equilibrium (4). If, as a 

limiting factor, no complexing is assumed and the measured manganese concentration corrected to activity is used in equilibrium (3), 

the value for K obtained is 0.84 a near approach to 1.0, the computed 

value. However, to assign any manganese to the complex form would 

lower the value for the computed K even further. Evidently, the 

amount of complex present is small, but a more refined experimental 

procedure is needed to evaluate the system.

At least three possible explanations could be offered for the apparent 

lack of complexing: (1) the complexing reaction is slow and had not
EQUILIBRIA AND OXIDATION RATE A35

reached equilibrium when the pH reached a constant value (this 

explanation is rather unlikely); (2) the complex form is less stable 

than the published data indicate; and (3) in filtering the solution 

before determining Mn+2, some dissolved manganese was precipitated. 

Any or all these factors might influence the results.

Several approaches can be used to learn more about the importance 

of the complex. If the solutions containing bicarbonate and the 

manganese bicarbonate complex are titrated with standard acid, as 

in the alkalinity titration made in a routine water analysis, the 

breakdown of the complex can be represented as

MnHCO3++H+^H2CO3 (aq)+Mn+2.

The equilibrium constant for this reaction is 7.2X102 if the standard 

free energy of formation of the complex is taken as 199.5 kg-cal, 

the value that is obtained from the published data on the complex. 

From this equation it can be shown that the pH at which the activities 

of the complexed and uncomplexed varieties are equal is about 4.3, 

where the total of bicarbonate and dissolved CO2 species present is 

3.4X1CT2 molal. The equilibrium solution under study sho^ild have 

about this activity of bicarbonate and related species, mainly in the 

form H2CO3 (aq). The end point of the titration would be well below 

a pH of 4.

Thus, a solution containing the complex would have a buffering 

capacity below a pH of 4.5, the usual end point of the alkalinity 

titration if the free energy of the complex were as low as indicated by 

the published data. The behavior of solutions during alkalinity 

titration should give evidence regarding the complex, but the system 

in equilibrium at 1 atmosphere of pressure of CO2 is not easy to 

evaluate in this way.

Accordingly, another set of solutions was made up to permit a 

longer time for complexing to take place, to provide more favorable 

systems for titration, and to avoid filtration of the sample used for 

manganese determination.

DETERMINATION OF EQUILIBRIUM CONSTANT

A system containing water, solid manganese carbonate, and carbon 

dioxide gas at 1 atmosphere pressure was allowed to come to equilibrium (constant pH). Aliquots containing solution and sclid were 

poured off into glass-stoppered bottles. The bottles were filled completely so that no gas space remained, and the aliquots were allowed to 

stand until the solution cleared. Portions were then taken from the 

bottles at predetermined times, and manganese and alkalirity were 

determined.
A36 CHEMISTRY OF MANGANESE IN NATURAL WATER

The equilibria involved in these solutions were Nos (2), (3), and 

(4), and probably

MnHC03++H+^±Mn+2+H2C03 (aq).

Because no gas phase is present, however, equilibrium (1) is not 

applicable. The only constituent available in excess is now the solid 

MnCO3 (c) and the system ought to reach a pH that ie controlled by 

equilibrium (3), with the balance between HCO3~ and H2C03 (aq) 

adjusting itself accordingly. The amounts of manganese complexed 

can be determined from analyses of the solution and from equilibrium 

computations.

After 3 days, the pH of the clear solution was determined to be 5.48, 

and an aliquot was titrated with standardized sulfuric acid. The 

bicarbonate content so determined was 235 ppm. Manganese content was determined to be 93 ppm. The value for bicarbonate includes both HCO3 " and MnHCO3+, and the value for manganese 

includes both Mn+2 and MnHCO3+. From the Debye-Hiickel limiting law the activities of the four principal ionic species present can be 

estimated. For this calculation, the Mn+2 and HCO3~ ions were 

assumed to dominate the system. Thus the total ioni°, strength can 

be estimated closely from the concentrations for total rianganese and 

bicarbonate species. The value obtained for ionic strength is 0.0054. 

The activity coefficient (7) for Mn+2 is computed to b<»< 0.76 and for 

the bicarbonate and also the complex is computed to be 0.93.

The amount of manganese complexed can be evaluated by a method 

of successive approximation. Assume first that no complexing has 

occurred. Activity data then can be substituted in eouilibrium (3), 

as follows:

[HCOrl]Mn+2j_

[H+]

[H+]=3.31X10- 6

[HC03 -]=3.85X0.93X10-3=3.58X10-3 

[Mn+2 ]=1.69X10-3 X0.76=1.28X10- 3 ;

hence,

7r_3.58XlO-8 Xl.28XlO-8 _1 

3.31X10-6 "

The equilibrium constant, K, however, should be l.C^as computed 

from free-energy data. The difference is assignable to complexing 

that has affected the activities of Mn+2 and HC03~.
EQUILIBRIA AND OXIDATION RATE A37

The activity value for manganese (uncomplexed) then is lowered 

by a suitable amount, and the concentration represented by this 

activity, calculated from the relation

n ,._[Mn+»]

An equivalent concentration of bicarbonate is assigned to this concentration of manganese, and its activity is computed. The activity 

is subtracted from [HCO3~] in equilibrium (3), and a new value for 

K is calculated. The procedure is repeated, if necessary, until the 

value of K obtained is 1.0. The amount of complexed manganese in 

terms of concentration is thus calculated to be 3.3X10"4 molar or 18 

ppm. The ratio of activities of complexed to uncomplexed manganese 

in this solution is 0.24. From equilibrium (4), the equilibrium 

constant Complex for the formation of the complex is computed to 

be 75.

The analysis and computation was repeated three times, after the 

solutions had stood for longer periods. The results are given in 

table 4.

TABLE 4. Experimental data on manganese bicarbonate complex

Test

1 ...

2......

la 

2a -

Reaction

time 

(hours)

49

150

216

360

Concentration (ppm)

C'Mn

93

78

83

64

CHCOJ

234

190

220

141

Activities (moles per liter)

[Mn+q

1.28X10-'

1.11X10-'

1.16X10-'

0.91X10-'

[HC03-]

3.58X10-'

2.90X10-'

3.36X10-'

2.17X10-'

[H+J

3.31X10-«

2.75X10-0

3.16X10-0

1.58XKH

Calculated K

(equilibrium 3)

1.37

1.17

1.23

1.25

Concentration of

complexed 

manganese

(moles

per liter)

3.3X10-4

1.4X10-4

2.1X10-4

1.7X10-4

[MnHC03+]

[Mn+s]

0.24

.11

JC

.17

K complex *

75

40

51

85

i Mean value for ^rcomPi«=63.

Tests 1 and la represented solutions with an excess of solid MnCO3 

present and no gas phase. Tests 2 and 2a were made on tl e same 

two sample bottles, but because some solution was taken out in the 

earlier runs, a gas phase was present. The bottles were kept stoppered 

while standing. The loss of bicarbonate and increase in pH was 

related in part to equilibration of the dissolved CO2 with a portion 

of previously dissolved CO2 that had migrated to the gas phase. 

At the same time, the amount of manganese present readjusted 

itself to the new pH and bicarbonate levels, with precipitation of 

manganese carbonate. Because the calculated data agree fairly well 

for all four tests, it seems likely that equilibrium was reached.

When the usual end point of a pH of 4.5 was approached during 

the titration of alkalinity in the samples, their behavior was unlike
A38 CHEMISTRY OF MANGANESE IN NATURAL WATER

that usually observed. The pH decreased as the titration proceeded 

and exhibited the usual inflection point near a pH of 4.5. However, 

the pH did not remain stable when the titration wa^ stopped but 

rose slowly. More acid was added; the pH again dropped to 4.5 

but then rose slowly again. Several repetitions of this occurred until 

the effect disappeared and a stable end point was reached.

The best explanation of this behavior seems to be that the bicarbonate present as the complex MnHCO3+ ion reacts more slowly 

with H+ than the HCO3~ ion does. Hence, a residue of complexed 

bicarbonate is left after the uncomplexed ions have been titrated. 

As these complex ions react slowly with H+, they cause the pH to 

rise, but are ultimately all converted to Mn+2 and H2CO3 (aq) at a 

pH of 4.5. When a stable pH of 4.5 was reached, the solution retained 

no buffering capacity.

The behavior of the solutions during the alkalinity titration is a 

qualitative indication that complexing of part of the manganese takes 

place. No quantitative conclusions are justified from this evidence, 

however.

CALCULATION OF FREE ENERGY OF THE COMPLEX

A value for the standard free energy of formation of the complex 

can be calculated from the equilibrium constant for the reaction

Mn+2+HC03-^MnHC03+

and from data in tables 1 and 4. The net change in standard free 

energy in the reaction is related to the equilibrium constant by the 

equation

On substituting numerical values for R and T and converting to 

base 10 logarithms, this becomes AF°= log ^(1.364). The free 

energy of formation of the complex is

log K=kF°MnSCC.,+

For the range of values of the equilibrium constant, the standard 

free energy of formation of the complex ranges frcm 196.9 to 

 197.3 kg-cal.

EFFECT OF THE COMPUEX ON MANGANESE EQU1TJBRIA

Experimental evidence obtained in this work regarding bicarbonate 

complexing may be summarized as follows:

1. A complex cation assumed to have the formula MnHCO3+ increases 

the solubility of manganese in solutions containing an excess of 

dissolved carbonate species.
EQUILIBRIA AND OXIDATION RATE A39

2. Reactions involving the bicarbonate part of the MnHCO3+ complex are slower than reactions of uncomplexed bicarbonate.

3. The equilibrium constant for the reaction

is 63, and the standard free energy of formation of the complex 

is 197.1 (±.2)kg-cal.

4. In natural water containing an activity of 940 ppm of 1 .; carbonate, half of any manganese present would be complexed. Thus 

the total solubility of manganese might be double that expected 

if complexing was not considered.

5. Bicarbonate complexing increases the solubility of manganese in 

natural water. When unusually large amounts of bicarbonate 

are present, the increase in solubility may be substantial.

Experiments reported by Listova (1959) indicate a somewhat 

greater solubility for manganese carbonate than the presert work 

indicates. The solid formed in Listova's experiments was described 

as rhodochrosite but was in a finely divided state. Its behavior 

was more like the precipitated MnCO3 used in this study tl an like 

the natural crystalline rhodochrosite described in the work of Garrels, Thompson, and Siever (1960). Data published by Listova are 

insufficient for calculation of free energies.

STABILITY FIEID OF BICARBONATE COMPLEX

Figure 6 represents a system similar to that of figure 2, except 

that activity of carbonate species is assumed to be 2,000 ppm in 

terms of HCO3~. Under these circumstances, most of the dissolved 

manganese is complexed when the pH is favorable. The equilibrium 

constant indicates that the ratio [MnHCO3+] : [Mn+2] actually has a 

value of a little more than 2, so that complexing is about two- thirds 

complete. The total solubility of manganese in the systeir is the 

sum of the solubilities of complexed and uncomplexed ions.

The stability boundaries that are calculated from equilibria in 

which bicarbonate-species activities are not involved are the same in 

both figure 2 and figure 6. Boundaries between fields of the oxides 

are identical in both diagrams.

The principal effect of increased activity of bicarbonate species 

is to enlarge the MnCO3 (c) (rhodochrosite) field and to add a new 

field in which the complex MnHCO3+ is the dominant form. The 

MnCO3 (c) field is enlarged at the expense of the area occupied in 

figure 5 by the solids Mn2O3 , Mn3O4 , and Mn(OH)2 .
A40 CHEMISTRY OF MANGANESE IN NATURAL WATER

</)

o

>

UJ

J

<

H 

Z 

Ul

o

o.

X

o 

o

Ul

a:

-0.

-0.80

FIGURE 6. Stability fields of manganese species in aqueous solution. Total dissolved manganese activity 

ranges from 0.01 to 100 ppm. Bicarbonate species activity is 2,000 ppm as HOOs". Symbol (c) indicates 

crystalline form.

No Mn(OH)2 (c) stability field exists in figure 6, owing to the 

expansion of the MnCO3 (c) field. In the HMn02~ area, the activity 

of dissolved manganese is governed by the equilibrium

MnC03 (c)+2H2O^HMn02 -+3H++C(V2

for which the value of K is 6.2X1CT46 and by the redox equilibrium 

involving HMnO2~ and Mn304 (c) used in previous diagrams.
EQUILIBRIA AND OXIDATION RATE A41

Some of the boundaries of the MnCO3 (c) field were computed 

from the same equations used in preparing figure 2, but large" values 

for HCO3~ activity were used in the computations. The MnCO3 (c) 

and Mn2O3 (c) fields did not touch each other in figure 2, however, 

and for their common boundary in figure 6 in the equilibrium applied 

was

Mn2O3 (c)+2HCO3-+4H++2e^±2MnCO3 (c)+3H2O,

where E°= 1.53 v.

The stability field of MnHCO3+ is controlled by a new series of 

equilibria. The boundary along which the complex breaks down to 

form Mn+2 and HCO3~ or H2CO3 (aq) is a function of pF. The 

equilibria governing concentrations along this boundary are:

Mn+2+ HCO3-^±MnHCO3+ 

MnHCO3+ +H+^Mn+2+H2CO3 (aq) 

H2CO3

Equilibrium constants for the reactions are given in table 2. The 

mathematical relationships that apply are:

[MnHC03+] _63X1Q1 [Mn+2 ][HCO3 -]~°'dA

H2C03]

[Mn+2][H2C03] 

]MnHCO3+][H+

=4.

=4.2X104 .

Also, because the total activity of the bicarbonate species is fixed 

and the activity of CO3~2 and OH~ is negligible,

[HCO3-] + [H2CO3]=3.28X10-2

and, because the activities of complexed and uncomplexed manganese 

are equal along the boundary,

[Mn+2] = [MnHCO3+].

These five equations can be solved simultaneously to give the pH 

at which the left boundary for the stability field of the complex 

should be drawn. The value obtained for the pH at the Mn+2- 

MnHCO3+ boundary is 6.35. This boundary is represented by a 

dotted line on the diagram and is independent of the total manganese 

concentration, provided that a large excess of uncomplexed bicarbonate remains present.
A42 CHEMISTRY OF MANGANESE IN NATURAL VATEB

The vertical segment of the boundary between TfnC03 (c) and 

MnHCO3+ is established from four equilibria:

MnCO3 (c)+H+^±Mn+2+HC03- 

where K=Q.3l

Mn+2+HCO3-^±MnHCO3+ 

where K=Q3

H2C03

and the known conditions at the boundary

[Mn+2] + [MnHCO3+l = 1 .8 X 1 0~7.

Actually, because the boundary is in the pH range where HCO3~ is 

strongly dominant, it is sufficiently accurate to set the activity of 

HCO3~ equal to 3.28X 10~2 molal. The number of equations required 

to find the equilibrium pH is thus decreased. The value of pH for 

the boundary was found to be 8.2. The boundary between the 

MnHCO3+ complex and MnO2 (c) was determined from the redox 

equilibrium

MnO2 (c) +HCOr+2e+4H+^MnHCO3++2H2O, 

where

E°= 1.28v,

and the first dissociation of carbonic acid. The amount of HCO3~ 

available decreases with pH, and the slope of the boundary therefore 

changes below a pH of 7.0.

Some Mn+2 is in equilibrium with the complex and the total 

solubility of manganese can be computed from this equilibrium and 

the Nernst equation if the activity of the complex is taken as 

2[Mn]-[Mn+2].

The stability-field diagrams show some of the effects of changes in 

bicarbonate activity and complexing. In figure 2, the equilibrium 

activity of Mn+2 in the presence of rhodochrositc and 100-ppm 

HCO3~ species is 0.10 ppm at a pH of 8.0. In figure 5, the activity 

of Mn+2 is about 1.0 ppm at this pH, owing to the decrease of HCO3~ 

species to 10 ppm. In figure 6, a system containing 2,000 ppm HCO3~ 

species activity is considered. If there were no complexing in this 

solution, manganese activity would be expected to b<? less than 0.01
EQUILIBRIA AND OXIDATION RATE A43

ppm at a pH of 8.0. Actually, the manganese activity at a pH of 

8.0 is considerably greater than 0.01 ppm, as shown in figure 6. 

Because activity coefficients would decrease as concentration ir creased, 

the manganese concentration that would be measured in a system 

corresponding to figure 6 would be nearly double the activity.

The solubility of the oxides is increased a little by complexing, as 

shown by the "bulge" in the oxide boundaries in the stabiTty field 

of the complex.

In figure 6, as well as in the other stability-field diagrams, it is 

assumed that anions containing C+4, such as carbonate and bicarbonate, are stable in strongly reducing systems. Actually, organic 

forms of carbon theoretically would be the stable varieties in most 

of the reduced-sulfur region. The redox reactions of carbon are not 

readily reversible and have been omitted from consideration here.

MANGANESE STTLFATE ION FAIR

Bjerrum, Schwarzenbach, and Sillen (1958) report an equilibrium 

constant of 1.9X102 for the reaction

Mn+2+SOr2^±MnSO4 (aq).

This suggests that sulf ate complexed with manganese is stable enough 

to be important in natural water. In a solution where activity of 

sulfate ions is 2,000 ppm (2.1X10"2 molal), about 80 percent of the 

manganese would be complexed.

[Mn+2][SO4 -

A water that had a concentration of 2,000 ppm of sulfate would have 

an ionic strength of considerably more than 0.06, and the activity 

of sulfate would be substantially less than 2,000 ppm under these 

conditions.

COMPARISON OF STRENGTH OF COMPLEXES

Figure 7 compares the strength of the bicarbonate and sulfate 

complexes and indicates the degree of complexing to be expected 

when both HCO3~ and SO4~2 are present. The diagonal dashed line 

on the diagram is based on the equilibrium

MnSO4 (aq)+HCO3-^MnHCO3+ +SO4-2
A44 CHEMISTRY OF MANGANESE IN NATURAL WATER

BICARBONATE ACTIVITY, IN PARTS PER MILLION
EQUILIBRIA AND OXIDATION RATE A45

for which the equilibrium constant is 0.32. When the activities of 

the complexes are equal,

[SO4-2]=0.32[HCO3-].

The diagonal line based on this equation separates the area in which 

the sulfate complex would be the dominant form from the area where 

the bicarbonate complex would be dominant. Lines parallel to this 

line could be plotted for any desired value of the ratio [MnE"CO3+] : 

[MnSO4 (aq)].

If it is assumed that the total activity of manganese in solution is 

the sum of the activities of Mn+2, MnHCO3+, and MnSO4 (aq) and 

that the amount of manganese involved is small enough that the 

bicarbonate or sulfate combined with it does not significantly affect 

the supply of the anions, the ratios of the total activity of manganese 

species in solution to the total of the two complexes can be readily 

calculated. The results of these calculations are shown in figure 7 

as contours that permit estimation of the proportion of manganese 

present as a complex when the activities of bicarbonate and sulfate 

are known.

The positions of the contours were calculated in two steps. First, 

parallel lines were drawn on the diagram for the following value? of the 

ratio

[MnHCO3+]:[MnSO4 (aq)]:

1/100, 1/10, 1/2, 1/1, 2/1, 10/1 and 100/1. The intersection of these 

lines with those for the activity of HCO3~ or SO4~2 that would give 

any selected percentage of complexing was then computed. For 

example, when the ratio

+1 [MnHCO 

[MnSO4 (aq)]

i

and the manganese is 90 percent complexed, the three species are 

present in the ratio [Mn+2] = 10, [MnHCO3+]=45, and [MnSC 4 (aq)] 

=45. In the equation

[MnH003+] _R,rHm -, 

-63[HCO° 1

the ratio on the left can now be determined as 4.5, and the value for 

[HCO3~] as 7.14X10"2 molal, or 4,360 ppm. Similar calculations 

for each of the other values for the ratio of the complexes gives a 

series of points to which a curve is fitted to represent 90 percent 

complexing.
A46 CHEMISTRY OF MANGANESE IN NATURAL \TATER

Figure 7 is based on activities and is directly applicable only to 

solutions with a rather small manganese concentration. However, 

the usual manganese concentration in natural water is less than 1.0 

ppm; and when suitable activity corrections are calculated, the graph 

will give useful indications of the state of manganese dissolved in 

natural water. Obviously, in many of these solutions, 25 percent 

or more of the dissolved manganese occurs as a complex. The complexed manganese can be expected to behave differently in chemical 

equilibria than does uncomplexed manganese. The complexes probably will also react at rates different from those c* uncomplexed 

manganese ions.

COMBINED EFFECTS OF BICARBONATE AND SULFATE COMPLEXES 

ON MANGANESE EQUILIBRIA

Figure 8 represents a system in which sulfur-species activity is 

2,000 ppm in terms of sulfate and bicarbonate-species activity is 

2,000 ppm as HCO3~ (same as in figure 6). The increased availability 

of sulfate is sufficient to complex manganese in a considerable part of 

the diagram. The stability field for MnHCO3+ is confined to the 

reducing region, because the sulfate complex would predominate when 

SO4~2 and HCO3~ activities are 2,000 ppm.

The boundaries of the MnSO4 (aq) field were determined from 

equilibrium constants and standard potentials.

For the Mn+2-MnSO4 (aq) boundary, the data needed are:

MnS04 (aq)+H+^±HS04-+Mn+2 , 

where #=4.6 X10- 1 ;

MnSO4 (aq)+8H++6e^±S (c)+Mn+2+4E,O

where S°=0.335 v; and similar redox equilibria for reduction to H2S 

and HS~.

For the MnSO4 (aq)-MnO2 (c) boundary, below a pH of 6.4,

MnO2 (c)+SOr2+4H++2e^MnSO4 (aq)+2H2O,

where E°= 1.291 v. This activity of MnSO4 (aq) in tt<?, area of dominance of the complex is 80 percent of the total manganese activity, 

and therefore the value for [MnSOJ used in computing the boundary 

is taken as 0.01 ppmX0.80 or 1.46X10' 7 molar.

Both the sulfate and the bicarbonate complexes affect the area of 

figure 8 that coincides with the area occupied by the bicarbonate 

complex in figure 6. The total activity of manganese in this area is

S[Mn] = [Mn+2]+[MnHCO3+]+[MnSOJ.
EQUILIBRIA AND OXIDATION RATE A47

-0.80

FIGUBE 8. Stability fields of manganese species in aqueous solution. Manganese activity ranges from 

0.01 to 100 ppm; bicarbonate-species activity is 2,000 ppm as HCOj-; and sulfur-species activity is 2,000 

ppm as SOr3. Symbol (c) indicates crystalline form and (aq) aqueous solution.

The relative amounts of the three species are in the ratio 1 to 2.1 to 4, 

where all the 2,000 ppm of bicarbonate species are present as HCO3~. 

To compute the position of the MnCO3 (c)-MnSO4 (aq) boundary 

for 0.01 ppm total dissolved manganese activity, a good approximation can be obtained from the equilibrium

MnCO3 

where K 55.

(aq)+HC(V,
A48 CHEMISTRY OF MANGANESE IN NATURAL

The total manganese activity is 1.8X10" 7 molal, and the fraction of 

the total activity due to MnSO4 (aq) is

[MnSOJ 4 _______

[MnSO4] + [Mn+2] + [MnHCO3+] ~7 . 

Multiplying

or

-7 molal.

The computation gives a pH of 8.5 for the boundary. When the 

total manganese activity is increased to 1.8X 10~6 mold, the pH is 7.5. 

Further increases in total manganese activity require a change in 

the fundamental equations. When manganese activity is 1.0 ppm, 

the pH evidently will be near the edge of the MnHCO3+ field, so that 

activities of the uncomplexed Mn+2 and the MnHCO3+ will be about 

the same. The ratio of species will then be 1 to 1 to 4 and the molal

1 8X10~S X4 MnSO4 (aq) activity will be 7; - Also, the bicarbonate

activity will be only about half the total at pH 8.5, or about 1.6X10"2 

molal. Substituting these values in the equilibrium expression gives 

a pH of 6.7 for the 1 .0-ppm-manganese activity line.

The positions of the 10- and 100-ppm lines were computed from the 

equilibrium

MnCO3 (c)+2H++S<V2^MnSO4 (aq)+H2CO3 (aq),

where K=1.2SX108, and the assumption that the Mn£O4 (aq) activity 

is four-fifths of the total activity of dissolved manganese species in 

this region.

From the equilibrium

HMnO2 -+SO4 - 2 +3H+^±MnSO4 (aq)+2F2O,

for which the value of .K"is 4.1X1036, it follows that the MnSO4 (aq) 

form will not predominate above a pH of 11.6 v^hen SO4~2 and 

CO3~2 activities are at the levels postulated.

Equilibria that were used to locate the boundaries of the MnS (c) 

field are the same as those in figure 5, except for the vertical segment 

of the MnCO3 (c)-MnS (c) boundary. This is calculated from the 

equilibrium

MnC03 (c)+HS-^±MnS (c)+CO3 -2+H+, 

for which the value of .K"is 3.5X10~12.
EQUILIBRIA AND OXIDATION RATE A49

The dashed lines in the HMnO2~ field are the same as those in that 

area in figure 6, except where they are affected by equilibria irvolving 

MnS (c).

When sulfate ions are reduced, free sulfur is a possible stable 

variety. Therefore, the equilibrium

S (c)+Mn+2+H2CO3 (aq)+2e^±MnCO3 (c)+H2S (aq),

where E°= 0.239 v, was considered in locating the dashed lines of 

dissolved manganese species below a pH of 6.0. This relation is 

independent of pH and gives rise to short horizontal segments in the 

lines for dissolved-manganese activities of 10 and 100 ppm.

The effects of sulfate complexing are plainly indicated in f gure 8. 

In the area where MnSO4 (aq) is dominant, the lines representing 

dissolved-manganese activity are each displaced about 0.3 pH unit 

to the right, as compared with their positions in the MnHCO3~ and 

Mn+2 areas.

Mah (1960) gives the free energy of formation of two manganous 

silicates, MnSiO3 and Mn2SiO4 . No computations have been 

attempted with these, however. Figure 8 shows about as many 

variables as can usefully be shown on a single diagram, and the 

silicates of manganese in nature are generally not simple enough to be 

represented by the compounds for which thermodynamic data are 

available.

OTHER INORGANIC COMPLEXES

The chloride complexes with divalent manganese evidertly are 

rather weak, although relatively little information seems to have been 

obtained about them. Kraus and Moore (1953) found that anionic 

complexes of manganese and chloride, with a coordination number 

of 6, such as of the type MnCl3~ and MnCl4~2 , are not formed in 

detectable amounts unless the chloride concentration exceeds about 4 

molal. According to Bjerrum, Schwarzenbach, and Sillen (1958) 

J. C. James reported in 1947 that the value of K for the equilibrium

Mn+2+Cl-^±MnCl+

is only about 1. For half the manganese to be complexed, the chloride activity would thus need to be about 1 molal, or some 35,000 

ppm. This chloride value is attained only in brines. This discussion 

is centered on the more dilute kinds of natural water, and therefore 

the chloride complexes are not given further consideration here.

Phosphate ions can form strong complexes with manganese. Bjerrum, Schwarzenbach, and Sill6n (1958) quote Jones, Mork, and
A50 CHEMISTRY OF MANGANESE IN NATURAL WATER

Davies (1949) who report a constant of 3.7X103 for the equilibrium

P3(V3 +Mn+2;^MnP309-.

The same publication quotes Jones and Monk (1950), who found a 

constant of 5.5X10 5 for the equilibrium

Also, Van Wazer and Campanella (1950) reported a constant of 

3X10 5 for the equilibrium

Apparently, only a few parts per million of polyphospb ate in solution 

should serve to complex most of any small amount of manganese 

that may be present. Complexing of manganese by orthophosphate 

and pyrophosphate is reported to be either weak or to be nonexistent 

by several investigators quoted by Bjerrum, Schwarzenbach, and 

Sillen.

Natural water usually does not contain phosphate in amounts 

sufficient to complex manganese. In treatment processes, however, 

polyphosphates are rather widely used to delay or to prevent precipitation of calcium carbonate or iron or manganese oxide or hydroxide. 

Polyphosphate anions are not stable in water and slowly change to 

orthophosphate; hence, the effects of adding polyphosphate are only 

temporary.

Rather large amounts of phosphate are required for complete complexing of calcium and magnesium, which are major dissolved constituents in water. A considerable effect on the behavior of calcium 

can be observed from an amount of polyphosphate thr.t is not nearly 

large enough to complex all the calcium. The retarding of calcium 

carbonate precipitation that results from small amounts of polyphosphate is termed the "threshold effect." It is generally attributed to 

dispersion of the small aggregates of CaCO3 as they form in the solution, so that the particles do not grow large enough to precipitate.

Manganese and iron are minor constituents in most waters. Illig 

(1957, 1960) stated that the amounts of polyphospl ate needed to 

control the precipitation of manganese are considerably in excess of 

the amount of manganese in solution, which suggests that a complexing effect is the principal one involved. The complex is destroyed 

both by oxidation of manganese and by degradation of the polyphosphate to form orthophosphate.

Phosphate generally constitutes a considerable par* of household 

detergents. Water polluted by domestic or industrial sewage can be 

expected to contain phosphates from this and other sources.
EQUILIBRIA AND OXIDATION RATE A51 

ORGANIC COMPLEXES

Soluble organic complexes containing iron were studied by Shapiro 

(1958). The writer (Hem, 1960b) found that ferrous iron formed a 

soluble complex with tannic acid and that the complexed iron is more 

slowly oxidized than the uncomplexed form. This feature of the 

chemistry of iron is at least locally important because the us;ial conditions to which surface water is exposed result in almost immediate 

oxidation of uncomplexed ferrous iron to the nearly insoluble ferric 

iron. The complexing action permits ferrous iron to be carred considerable distances in solution in streams or in soil and subsoil moisture.

Enough evidence of the behavior of manganese already has been 

presented to show that small amounts of manganese can be stable in 

a soluble form in most surface water, either as Mn+2 or as bicarbonate 

or sulfate complexes. Media that could transport manganese in the 

reduced soluble form are commonly present and organic complexing 

is not essential for such transport. Nevertheless, published literature 

ascribes certain features of manganese behavior to organic complexes, 

and some work was done to see if a tannic acid-manganese complex 

could be formed.

A solution containing 10 ppm Mn+2 and 500 ppm tannic rcid was 

prepared, together with a similar tannic acid solution containing no 

manganese. Aliquots of these solutions were raised to selected levels 

of pH by adding sodium hydroxide solution, and their percentage 

transmittance was measured by use of a Beckman model-B spectrophotometer over the wave-length range 400 to 650 m/u. The resulting 

curves of transmittance plotted against wavelength were similar for 

the two solutions below a pH of 5. At a higher pH the solutions 

had a deeper color, but the shapes of the transmittance wave-length 

curves remained about the same; apparently, no distinctively colored 

complex was formed.

At a pH of 7.0, the manganese tannic acid mixture had a pronounced 

greenish color. This solution grew darker slowly on standing, although the rate of change became almost imperceptible after an hour 

or so. The corresponding tannic acid solution at a pH of 7.0 (no 

manganese present) also had a greenish color, but at first it was a 

paler shade than the solution containing manganese. As time- passed, 

the color of the tannic acid solution deepened. Apparently, most of 

the change of color in the solutions was a result of changes in the 

tannic acid as the pH was raised. Manganese seemed to accelerate 

these changes.

A solution of tannic acid and manganese raised to a pH of 8.5 

contained a small amount of finely divided black material that may 

have been a colloidal dispersion of a tannic acid-Mn*4 complex similar
A52 CHEMISTRY OF MANGANESE IN NATURAL WATER

to the black material formed by tannic acid in combination with 

ferric iron.

The results of these experiments suggest that manganese could be 

present as a constituent of organic coloring matter in surface waters, 

although no well-defined complex of tannic acid and divalent manganese is indicated. Colloidal suspensions of manganese oxides, however, may be stabilized by organic matter at pH levels near neutrality 

or in the alkaline range.

The behavior of mixtures of tannic acid and Mn+2 was studied 

further by mixing tannic acid solutions of concentrations from 5 to 

500 ppm with manganese stock solution, adjusting the pH, and 

observing changes in concentration of manganese while the solutions 

were standing in contact with air over a period of 3 veeks. Some 

of the solutions retained some manganese that passed through a 

plastic-membrane filter whose (pore diameter was 0.45ju) at pH as 

high as 10 for the entire 3 weeks. However, some of this manganese 

may have been colloidal particles of oxide.

Because the tannic acid molecule is about six times as heavy as 

the Mn+2 ion, a 1 to 1 complex requires a ratio of 6 ppm of tannic 

accid to 1 ppm of manganese. Thus, complete complexing of 10 

ppm Mn+2 would not occur in solutions containing les^1 than about 

60 ppm tannic acid. Apparent effects on the oxidation of manganese 

in the presence of 5.0 and 50 ppm concentrations of tannic acid were 

noted, however. These may result from the reducing effect of tannic 

acid.

Complexes of manganese with many different orgr.nic ions are 

cited by Bjerrum, Schwarzenbach and Sillen (1957). In general, the 

stability constants for ferrous and manganous complexes with the 

same organic ligand show that the ferrous complexes are stronger. 

The behavior of manganese with tannic acid solutions observed in 

this study, seems to agree with this general observation. Irving and 

Williams (1953) reported that several complexes containing manganese 

were less stable than similar complexes containing 5 otl ?T transition 

metals, including iron.

BATES OF CHEMICAL REACTION

Thermodynamic calculations such as those made in preparing figures 

1-8 require that the evaluated systems be in a state of chemical 

equilibrium. Whether equilibrium actually is attained in a system 

depends on the rates of the various chemical reactions that may take 

place and on the length of time that the reacting substances and 

products remain undisturbed in contact with one another. Reaction 

rates are affected by concentrations of reactants and products, by 

temperature, and by the action of catalysts; but they are not readily
EQUILIBRIA AND OXIDATION RATE A53

predicted on theoretical grounds and must generally be determined 

by experiment. A convenient measure of reaction rate is the "halflife" time, the length of time in which half the reacting substances 

present at the beginning of this period of time are converted to 

products.

Whether a given reaction rate is "fast" or "slow" depends on the 

point of view of the individual. In analytical chemistry the most 

useful reactions have short hah0 lives, preferably not more than a few 

seconds. The reactions of value in industrial chemistry may b? slower, 

but for most processes a half-life time greater than an hour or so 

would be excessive.

Although man's manipulation of chemical processes emphasizes 

rapid reactions, in nature the pace can be much more leisurely. In 

geochemical processes, such as ore deposition or metamorphic alteration of rock minerals, reactants and products may remair in undisturbed contact for many thousands of years; hence, equilibrium 

can be reached even in very slow reactions. Natural processes that 

take place in surface-water bodies or in ground water near the land 

surface, however, are subject to fluctuations in temperature and 

pressure. Because the solutions generally are in motion, whereas the 

solids are stationary, or move at a slow rate, the time available for 

equilibrium to be established is ordinarily not long in the geologic 

sense. In surface-water systems, a reaction whose half-life time is a 

few hours will generally reach equilibrium; and in grourd-water 

systems, a reaction whose half-life time is many days, or evn a few 

years, can generally be expected to reach equilibrium.

Some knowledge of the rates of reaction of manganese is essential, 

so that one may know the extent to which thermodynamic calculations 

can be applied. Obviously, such calculations have little practical 

usefulness if equilibrium cannot be expected to be reached in nature. 

Approximate indications of rate will suffice for this purpose. A closer 

determination of reaction rates for oxidation of Mn+2 in aerated water 

and the effects of other ions on the rate is of considerable practical 

value as this reaction is important in water treatment.

KINETICS OF MANGANESE OXIDATION AND PRECIPITATION

No data on the kinetics of oxidation of Mn+2 to Mn+3 ard Mn"1"4 

in aerated solutions resembling natural water could be found in published literature. The general features of similar oxidation reactions 

such as those involving iron may give some indication of what should 

be expected. Stumm and Lee (1961) studied the oxidation of ferrous 

iron in aerated water containing bicarbonate and carbonate ions. 

They concluded that the oxidation reaction is first-order with respect 

to ferrous-iron content and that the rate is strongly influenced by pH.
A54 CHEMISTRY OF MANGANESE IN NATURAL WATER

An increase of 1 pH unit brought about a 100-fold increase in oxidation rate.

Manganese oxidation reactions are likely to be more complicated 

than the ones for ferrous iron, because manganese car be oxidized to 

the Mn+3 or Mn*4 states. The bicarbonate and sulfate complexes 

of reduced manganese can also be expected to affect the reaction rate.

In earlier experiments the oxidation of manganese was accompanied 

by a rise in H+ activity. A series of titrations was m ade in an effort 

to evaluate this effect and to see if an indication of reaction rate as 

affected by pH could be obtained. Distilled-water samples containing 

2.0 mg of Mn+2 were prepared and were raised to a previously selected 

pH by rapid addition of O.QlN KOH from a burette. The solution 

was maintained at a constant pH by additions of the standard base, 

and the burette readings were plotted against elapsed time. The 

results obtained at a pH of 8.0, 8.5, and 9.0 are given in figure 9.

The oxidation and precipitation at a pH of 9.0 occurred rapidly and 

appeared to be complete in about 5 minutes. At lower pH, the 

reaction was much slower and did not appear to reach a definite end 

point. The amount of base consumed was also greater at lower pH. 

No quantitative interpretation of the results in figure 9 is possible. 

Evidently, the products of the reaction change as the pH changes, 

and the effects of carbon dioxide species on the titrations are difficult 

to evaluate. Air was not excluded from the samples, but the base 

used was free of carbonate.

To provide more specific data, a second series of titrations was 

made. In this series, the pH was held constant and samples were 

taken at intervals of a few minutes. The samples were filtered at 

once through a plastic membrane filter having a pore diameter of 

0.45/z, and manganese content of the filtrate was determined. The 

results for solutions held at a pH of 8.5 and 9.0 are shown in figure 10.

Although the individual points are somewhat scattered, they give 

a fairly well defined linear relationship between log of dissolved 

manganese concentration and time and suggest that the oxidations 

at these two pH values are first order with respect to manganese 

concentration. The half-life time at a pH of 9.0 is about 1% minutes; 

and at a pH of 8.5, about 30 minutes.

Precautions were taken to exclude carbonate from the standard 

base used, but the solutions were open to the air during the titration. 

Effects of CO2 introduced from this source probably were minor.

EFFECT OF STJLiFATE

The effect of sulfate on the rate of oxidation was studied in a series 

of titrations in which pH was held at 9.0, 9.2, and 9.5. The solutions 

used were prepared in the same way as in the preceding series, except
0.01

10 20 30 40 50 60 70 80 

TIME, IN MIN

FIGURE 9. Eate at which base KOH, was consumed to maintain constant pH in three solutions, e

species were absent.
A56 CHEMISTRY OF MANGANESE IN NATURAL

ec 0.001

TIME.IN MINUTES 

0 10 20 30 40 50 60 70 80 90 100

S 0.0005}- 

Q- 0.000^

<2 0.0003p°^

z o.ooo2r

-. ^

o o.oooi 

y o.oooos

§0.00004

^ aoooos

w

z 0.00002

PH 9.o with St

FIGXJEE 10. Eate at which manganese was oxidized at a pH of 8.5 and 9.0 in the absence of sulfate, and at a 

pH of 9.0 with a sulfate concentration of 2,000 ppm.

that sodium sulfate was added to give a concentration of 2,000 ppm 

SO4~2. Samples were withdrawn for manganese determination at 

various time intervals. Results obtained at pH 9.0 are shown in 

figure 10.

The relationship between log CMn+t and time is approximately 

linear during the first hour, but the rate appears to increase somewhat 

thereafter. A similar effect was noted in later studies using solutions 

containing bicarbonate. Extrapolation of the initial rate shows that 

when pH is 9 the apparent oxidation half life is about 80 minutes in 

the presence of 2,000 ppm sulfate, compared to 1.5 minutes in the 

absence of sulfate. The presence of sulfate also slowed oxidation 

rates at a pH of 9.2 and 9.5, but no rates of oxidation at these pH 

levels for sulfate-free solutions are available. The effect of sulfate 

is undoubtedly related, at least in part, to the presence of MnS04 (aq) 

complex.

EFFECT OF BICARBONATE

A further series of experiments was made to determine the effects 

of bicarbonate ions on the rates at which manganese is oxidized and 

is precipitated. In this series, solutions containing 2.0 mg of Mn+2 

in 100 ml of distilled water with 100 ppm of HC03~ present were 

prepared. The pH of these solutions was then rrised to values 

near 9.0. The buffering effect of bicarbonate presented the pH 

from changing rapidly as a result of the hydrolysis of oxidized manganese and permitted a somewhat closer control of pH than was
EQUILIBRIA AND OXIDATION RATE A57

possible in the other experiments. During each titration, an aliquot 

was removed every few minutes and filtered through a plasticmembrane filter with a pore diameter of 0.45 /z. The manganese 

concentration of each of the filtered aliquots was determined. Figure 

11 is a graph in which the concentrations of manganese in solution 

are plotted against time elapsed after raising the pH. The pH of 

the three solutions represented in figure 11 was maintained at 9.0, 

9.1 and 9.3.

Manganese remaining in solution, presumably as Mn+2 or 

MnHC03+, was measured directly in this experiment, and the results 

in figure 11 are comparable to those in figure 10.

Data in figure 11 show that the bicarbonate ions slowed the reactions by which manganese was precipitated at about the rate as the 

sulfate ions did. At a pH of 9.0, for example, the half-life time is 

about an hour when bicarbonate is present. In the absence of bicarbonate and sulfate, the half life was 1% minutes at this pH.

The half life at a pH of 9.0 in the presence of 2,000 ppm SO4~2 was 

about 80 min. The results suggest that bicarbonate ions are somewhat more effective in slowing the oxidation than are sulfate ions.

The relationship of log CMn+» and time was nearly linear for the 

first part of each of the titrations, which suggests that, at least at 

first, the oxidation at constant pH is a first-order reaction whose 

rate depends only on the amount of Mn+2 available. Rate constants 

for this part of the reaction could be calculated, but they would 

not represent an overall reaction rate.

As the reaction proceeds, the rate tends to increase; as stown by 

the changing slopes of the curves in figure 11. This suggests that 

the rate of manganese precipitation may be increased by autocatalysis that is, self-catalysis by the oxide causes further oxidation and 

precipitation of manganese.

The effect of bicarbonate on the oxidation rate does not seem to 

be directly related to complexing, because only a small amount of 

complexing occurs when 100 ppm of bicarbonate is present. To 

obtain comparable rates in duplicate experiments where bicarbonate 

is supposedly absent requires considerable care to exclude atmospheric carbon dioxide from the standard base, and the base must be 

prepared from carbonate-free hydroxide. Small amounts of carbonate 

species introduced in the reagent have a very marked effect on the 

reaction rate.

FACTORS THAT MAY INCRKASE REACTION RATE

Although the laboratory work with HCO3~ and S04~2 ions showed 

that they retarded the rate of oxidation of manganese, otl er substances may increase the rate. Some micro-organisms can greatly
LITER PER MOLES IN CONCENTRATION, MANGANESE 

o.

X 

O,

± 5 

- s 

I I I I I 

7

7

7

111 U*M I 

8QY £0 NI nvaoxvN: 
EQUILIBRIA AND OXIDATION RATE A59

increase the rate of oxidation of iron and others are knom to be 

effective in catalyzing the oxidation of manganese.

Futral and Ingols (1953) reported that 0.5 ppm CuSO4 that had 

been added to a water-supply reservoir speeded the removal of manganese from the water by oxidation. Small amounts of copper also 

catalyze the oxidation of ferrous iron. Dhar and Kishore (1950) 

reported that Mn(OH)2 was oxidized in the presence of otter metal 

cations at a more rapid rate as pH was increased and that most of 

the other metallic hydroxides that were studied increased the rate. 

Some of the other metal hydroxides, however, had no effect and 

AgOH decreased the rate.

SUMMARY

The results of the rate experiments may be summarized ae follows:

1. The rate at which manganese is oxidized and precipitated from 

aerated solutions is affected by pH and by other ions that may 

be present.

2. The rate is increased by raising the pH and is decreased by presence 

of HCOr and SOr2 .

3. The rate at which manganese is precipitated by the oxidation 

effect of air is rapid enough to be of use in water treatment 

only at pH 9.5 or more unless the water is unusually pure or 

if the reaction is catalyzed.

4. Although the rate of oxidation of manganese is slower than that 

of iron, in ordinary ground-water systems the time required 

for circulation is generally long enough so that equilibrium is 

probably generally reached. However, some metastable reduced 

manganese could be transported for considerable distances in 

surface streams.

5. The rate experiments suggest that the chemical equilibrium 

needed for application of the stability-field diagram is likely to 

exist in most natural-water systems where a sufficient supply of 

manganese minerals is present.

COPRECIPITATION OF MANGANESE WITH 

FERRIC HYDROXIDE

In studies of the chemistry of iron, Hem and Skougstad (1960) 

found that when ferric hydroxide precipitated from solutions containing about half a part per million of copper, the copper also was removed 

when the pH of the solution was high enough. The critical level for 

pH was found to be about 5.5. Above this level of pH, appreciable 

amounts of copper were removed from solution by the precipitated 

ferric hydroxide.

Manganese and iron often occur together in natural water. If the 

water is subjected to mild oxidation or if the pH is raised, the r?latively
AGO CHEMISTRY OF MANGANESE IN NATURAL WATER

unstable ferrous iron is likely to be oxidized and precipitated. Manganese should remain in solution while these processes are going on, 

however, unless rather strong oxidizing agents are present or the pH 

is very high, but there is a possibility of coprecipitation of manganese 

with the iron. This possibility was investigated in simple solutions 

containing ferrous iron that was being oxidized and from which Fe 

(OH) 3 was precipitating. A series of solutions each containing 10 

ppm Fe+2 and 10 ppm Mn+2 was prepared and their pH was raised 

to a selected level ranging from 6.0 to 8.0 in half-unit steps. The iron 

began to precipitate from all the solutions almost at once. The pH 

decreased slowly in each solution. After 3 days the pF, Eh, and the 

Fe+2 and Mn+2 contents of each solution were measured. This was 

repeated 10 days and 14 days after the start of the experiment. The 

results are given in table 5. A comparison of the results of the coprecipitation experiment with similar solutions containing no iron can 

be obtained by comparing tables 3 and 5.

Solutions 1-3 in table 3 and 3-5 in table 5 shoud be closely comparable. The solutions initially brought to a pH of 7.0 or 7.5 lost very 

little manganese in either experiment and attained about the same pH 

at the end of 3 to 5 days. The Eh of the samples containing manganese and iron, however, was considerably less at 3 days than that of 

the ones which contained only manganese after 5 days. The Eh 

values probably were affected by iron species.

TABLE 5. Coprecipitation of manganese with ferric hydroxide

Solution

!___-______ ... __ . __ __

2__ ______________________

3__ ___ __ ___ __-

4

5 __ --__ ___

Time 

(hours)

0

72

240 

408

0

72

240 

408

0

72

240 

408

0

72

240 

408

0

72

240 

408

Mir1* concentration (ppm)

10

10

10 

10

10

10

10 

10

10

9.8

9. 5 

10

10

9. 8

9.2 

9.2

10

6. 2

5.8 

5.0

Fe+s concentration (ppm)

10

3.2

10

.04

10

. 00

10

.00

10

.00

pH

6.0

6. 2

4.7 

4.7

6.5

6.5

6.4 

6. 1

7.0

6.7

6.6 

6.2

7. 5

6. 8

6.7 

6. 3

8.0

6.9

6. 8 

6.4

Eh (volts)

0.24

. 47 

. 46

. 18

.40 

. 41

.20

. 39 

.41

.20

. 39 

.41

.20

.39 

. 41
EQUILIBRIA AND OXIDATION RATE A61

The samples initially brought to a pH of 8.0 provide a comparison 

of manganese behavior in the presence of iron with that when iron 

is absent. When iron was present, about 40 percent of the manganese 

was removed from solution after 3 days and 50 percent, after 17 days. 

When manganese was present alone, no important amount was lost 

from solution over a 5-day period.

The amount of manganese lost from solution at a pH of 8.0 probably resulted from some manganese being incorporated in tbe ferric 

hydroxide crystal lattice. The amount lost appears to be too large 

to be attributed to adsorption on the small amount of precipitate that 

was present. The loss between the 3-day and 17-day observations, 

however, could have been related to adsorption.

At a pH of 8.0 some direct oxidation of manganese could 1 ave occurred but ths solution probably did not remain at that pH long enough 

for much of this to take place.

Some minor amounts of adsorption of manganese may have taken 

place at a pH of 6.7 or more.

In the work of the writer (Hem and Skougstad, 1960) with coprecipitation of copper, the critical pH above which copper ions were removed 

from solution by ferric hydroxide precipitates was near 5.5. Cupric 

ions are somewhat smaller than Mn+2 ions; hence, stronger adsorption 

effects would be expected for copper than for manganese. Thh agrees 

with the observation that manganese is not adsorbed until the pH 

reaches 6.7. The strength of negative charges on ferric hydroxide 

particles increases with increasing pH from the isoelectric point, which 

is reported by Krause (1928) to be 5.2.

Krauskopf (1957) has discussed the geochemical factors that may 

govern the separation of manganese from iron in sedimentary deposits.

COPRECIPITATION OF MANGANESE AND

OTHER ELEMENTS

The literature contains many papers on the coprecipita tion of 

manganese with other elements. Hikime (1954) reported that as 

much as 2 mg Mn+2 could be carried down on 100 mg of A1(OH)3 . 

The solutions used in these studies were relatively high in manganese 

content.

The precipitation of other elements with manganese orides is 

evidently a common occurrence, as suggested by the abundant 

impurities in some types of manganese ores. The tendency for 

manganese oxide to adsorb cations is demonstrated by the composition 

of manganese-bearing nodules recovered from the floor of the Pacific 

Ocean. Goldberg (1954) found that the nickel and copper contents 

of manganese nodules was proportional to the manganese content.
A62 CHEMISTRY OF MANGANESE IN NATURAL TFATER

GENERAL CONCLUSIONS

The general features of the chemical behavior of manganese in 

natural water and in the laboratory are in accord with chemicalequilibrium calculations. Reasonable agreement between laboratory 

data and calculated data was obtained in dilute solut: ons of Mn+2 in 

distilled water. The oxides produced, however, wero not pure and 

some appeared to be hydrated.

In the presence of bicarbonate and sulfate anions and in the con 

centration range common in natural water, manganese displays a 

pronounced tendency to form complexes. These complexes can be 

an important factor in manganese chemistry. Tv e equilibrium 

constant for the complexing reaction

Mn+2 +HCO3 -^±MnHCO3+

was found to be 63. An uncharged ion pair, MnSO4 (aq), which is 

reported in published literature to be somewhat more stable than the 

bicarbonate complex, is formed with sulfate. In the presence of an 

activity of 250 ppm of sulfate and 470 ppm of bicarbonate, and a 

total dissolved manganese activity not exceeding about 1.0 ppm, 

half of the manganese would be present in the form of complexes. 

Manganese may form strong complexes with phosphate and with 

some organic ions.

In general, the effects of Eh, pH, and commonly occurring anions 

in natural water are less restrictive on the solubility of manganese 

than they are on the solubility of iron. Hence, amounts of manganese 

from a few-hundredths ppm to 1.0 ppm may often be retained in 

solution at equilibrium, either in ground water or in surface water. 

However, at a pH above 8, oxidation by air may lower the concentration 

of dissolved manganese to less than 0.01 ppm.

Rates of manganese oxidation are increased by raising the pH and 

are decreased by the presence of SO4~2 or HCO3~. The oxidation and 

the precipitation reactions are complex and may yield mixed products. 

Experimental studies indicated that the reactions are first order with 

respect to manganese, but that the rates are increased by autocatalytic 

effects of the oxides produced.

LITERATURE CITED

Bjerrum, Jannik, Schwarzenbach, Gerold and Sill£n, L. G., 1958, Stability constants of metal-ion complexes, pt. 2, Inorganic Ligands: Chem. Soc. 

(London) Spec. Pub. no. 7, 131 p.

Chaberek, S., Jr., Courtney, R. C., and Martell, A. E., 1952, /3 hydroxyethyliminodiacetic acid, pt. 2 of Stability of metal chelates: Am. Chem. Soc. Jour. 

v. 74 p. 5057-5060.
EQUILIBRIA AND OXIDATION RATE A63

Dhar, N. R., and Kishore, N., 1950, Mn(OH) 2 in the presence of other metallic 

hydroxides, pt. 3 of Oxidation of metallic hydroxides: India Nat. Acad. 

Sci. Proc., v. 19A, p. 89-113.

Futral, F. L., and Ingols, R. S., 1953, Copper catalysis for manganese oxidation: 

Am. Water Works Assoc. Jour., v. 45, p. 804-806.

Garrels, R. M., 1960, Mineral equilibria at low temperature and pressure: New 

York, Harper and Bros., 254 p.

Garrels, R. M., and Naeser, C. R., 1958, Equilibrium distribution of dissolved 

sulfur species in water at 25° C and one atmosphere total pressure: Geochim. 

et Cosmochim. Acta, v. 15, p. 113-130.

Garrels, R. M., Thompson, M. E., and Siever, R., 1960, Stability of some carbonates at 25° C and one atmosphere total pressure: Am. Jour. Sci., v. 258, 

p. 402-418.

Goldberg, E. D., 1954, Chemical scavengers of the sea, [pt.] 1 of Marine geochemistry: Jour. Geol., v. 63, p. 249-265.

Hem, J. D., 1961, Calculation and use of ion activities in natural water cl ?mistry: 

U.S. Geol. Survey Water-Supply Paper 1535-C, PC1-C17.

 1960a, Restraints on dissolved ferrous iron imposed by bicarbonate, redox 

potential, and pH: U.S. Geol. Survey Water-Supply Paper 1459-B.

1960b, Complexes of ferrous iron with tannic acid: U.S. Geol. Survey

Water-Supply Paper 1459-D.

Hem, J. D., and Cropper, W. H., 1959, A survey of ferrous-ferric chemical equilibria and redox potentials: U.S. Geol. Survey Water-Supply Paper 1459-A.

Hem, J. D. and Skougstad, M. W., 1960, Coprecipitation effects in solutions 

containing ferrous, ferric, and cupric ions: U.S. Geol. Survey Water-Supply 

Paper 1459-E.

Hikime, Seiichiro, 1954, Coprecipitation of manganese with aluminum hydroxide: 

Japan Analyst, v. 3, p. 239-242.

Illig., G. L., Jr., 1957, Glassy phosphates in water treatment: Am. Water Works 

Assoc. Jour., v. 49, p. 805-816.

 1960, Use of sodium hexametaphosphate in manganese stabilization: 

Am. Water Works Assoc. Jour., v. 53, p. 867-874.

Irving, H. M., and Williams, R. J. P., 1953, Stability of transition metal complexes: Jour. Chem. Soc. (London) v. 1953, p. 3192-3210.

Klotz, I. M., 1950, Chemical thermodynamics: New York, Prentice-Hall, p. 

329-332.

Krause, Alfons, 1928, Transformation of ortho-ferric oxide hydrate into metaferric oxide hydrate: Zeitschr. Anorg. u. Allg. chemie, v. 176, p. 398-402.

Kraus, K. A., and Moore, G. E., 1953, Anion exchange studies. VI. The bivalent transition elements Mn to Zn in HC1: Am. Chem. Soc. Jour. v. 75, pp. 

1460-1462.

Krauskopf, K. B., 1957, Separation of manganese from iron in sedimentary processes: Geochim. et Cosmochim. Acta, v. 12, p. 61-84.

La timer, W. M., 1952, Oxidation potentials, 2d cd.: New York, Prentice-Hall, 

392 pp.

Listova, L. P., 1959, The precipitation of rhodochrosite from solutions of manganese chloride, sulfate, and bicarbonate: Doklady. Akad. Nauk. SSSR, v. 

124, no. 1, p. 183-186.

Mah, A. D., 1960, Thermodynamic properties of manganese and its compounds: 

U.S. Bur. Mines, Report of Investigations, 5,600, 34 p.

Pourbaix, M. J. N., 1949, Thermodynamics of dilute aqueous solutions: London, 

Edward Arnold and Co., 136 p.
A64 CHEMISTRY OF MANGANESE IN NATURAL WATER

Shapiro, Joseph, 1958, Yellow acid-cation complexes in lake Trater: Science, v. 

127, p. 702.

Stumm, Werner and Lee, G. F., 1961, Oxygenation of ferrous iron: Ind. and 

Eng. Chemistry v. 53, p. 143-146.

Van Wazer, J. R., and Campanella, D. A., 1950, Complex ion formation in polyphosphate solutions, [pt.] 4 of Structure and properties of the condensed 

phosphates: Am. Chem. Soc. Jour. v. 72, p. 655.

o

f