reposition and Soluticn 

f Manganese Oxides

;OHN D. HEM

-T EMISTRY OF MANGANESE IN NATURAL WATER

EOLOGICAL SURVEY WATER-SUPPLY PAPER 1667-B

V ITED STATES GOVERNMENT PRINTING OFFICE, WASHINGTON : 1964
UNITED STATES DEPARTMENT OF THE INTERIOR 

STEWART L. UDALL, Secretary

GEOLOGICAL SURVEY 

Thomas B. Nolan, Director

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

Washington, D.C., 20402
CONTENTS

Page

ostract- ______-_-_---__--___--__-_-______--_________-_-__________ Bl

atural occurrence of manganese__________________________________ 1

Manganese in soil and plants____________________________________ 2

Manganese in river water.._____________________________________ 3

Streambed manganese oxide deposits______---___-____--____^.____ 5

Manganese in lakes and reservoirs.______________________________ 7

Desert varnish______________________________________________ 10

Deep-sea nodules._____________________________________________ 11

Separate precipitation of iron and manganese. ______________________ 11

Manganese in ground \vater.__________-_-_--_____-________-_.___ 13

iboratory studies of manganese deposition..__________--__--_-____..__ 14

Nature of sands utilized._______________________________________ 15

Cation-exchange adsorption of manganese______---_-_-___-__-__. 15

Oxidation equilibrium experimentS-___--_----_-__--_-.-_ -_____._. 18

Effects of sand on oxidation rate____________________________ 22

Adsorption column experiments____--__-_---__-________ 33

^plication of laboratory and theoretical data to natural conditions,_ ______ 35

Coprecipitation of iron and manganese__________________________ 38

.immary and conclusions_-_-___----__-_---_-----_----_--_---_-_____ 40

iterature cited_________________-__________-__--_-_-_--____._-___ 41

ILLUSTRATIONS

Page

FIGURE 1. Relationships among dissolved concentrations and adsorbed 

quantities of calcium and manganese in adsorption column 

containing sand 2_.._______--__-_-____---__-_-__-___ B18

2-4. Rate of loss under different conditions of manganese from

solutions in contact with sand 2___________________ 23, 24, 26

5. Rate of loss of manganese from solution at three different

pH levels...--_-__-..-___________________________ 28

6. Rate of loss of manganese from solutions stirred with 15 g.of

sand 2_-__________..___-____-__-_--_.--__---_-______ 30

7. Effects of pH and dissolved sulfate on rate of oxidation of

manganese__. -___..___--___--__--_-_--_--_______-__ 32

in
17 CONTENTS

TABLES

TABLE 1. Equilibrium concentrations of caclium and1 manganese in 

solution and adsorbed on feldspathic sand-________________

2. Manganese concentration, pH, and Eh of wate~ in contact with 

siMca sand___________________________________________

3. Manganese concentration, pH, and Eh of unbuffered solutions 

on contact with feldspathic sand_________________________

4. Manganese concentrations, pH, and Eh of solutions containing 

bicarbonate ions in contact with feldspathic sand___________
CHEMISTRY OF MANGANESE IN NATURAL WATER

DEPOSITION AND SOLUTION OF|MANGANESE OXIEES

By JOHN D. HEM

ABSTRACT

Manganese is essential for plant growth and some aquatic species contain 

msiderable amounts. Manganese oxide deposits, generally mixed with iron 

side, occur in streambeds and in Lakes in some areas. Manganese oxide is 

Q important constituent of desert varnish.

Feldspathic sand removes manganese from dilute solutions by catio^ exhange, and above pH 7.0 the sand acts as a catalyst in the oxidation of mangaous ions by aerated water. The catalytic effect was strongest in sand vhose 

xchange positions had been saturated with manganese and which had been 

tioroughly dried. Manganese was a little more strongly absorbed than calcium. 

Manganese oxide coprecipitated with ferric hydroxide in aerated solutions 

.bove pH 6.5. Mixed deposits of iron and manganese oxides can form in fresh 

rater environments at a pH and redox potential where manganese oxide would 

iot be precipitated alone if the ferrous iron activity in solution is maintained 

it a very low value.

NATURAL OCCURRENCE OF MANGANESE

Manganese oxides may accumulate on the surfaces of rocks in the 

beds of certain lakes and streams, and are a major part of the dark 

;;urface coating known as desert varnish which occurs on the surfaces 

f rocks in some arid regions. Manganese also is a widespread though 

isually rather minor constituent of soils. Whether manganese is presr nt in soils as oxide, as a constituent of silicate or other minerals, or 

i s held by adsorption at the surfaces of clay minerals probably depends 

i x n numerous factors that vary from one soil to another. Manganese 

oxide accumulations below the surface of rice-paddy soils are menfioned by Kawaguchi and Matsuo (1955) and have been observed and 

i-tudied in many places by other soil scientists. Nodules of manganese 

dioxide admixed with iron oxide occur on the floors of ocean basins.

These accumulations of manganese oxide are of interest ir the 

chemistry of manganese in water because they represent deposition 

of a substance which was once carried in solution. The purpose of

Bl
B2 CHEMISTRY OF MANGANESE IN NATURAL WATER

this discussion is to identify some of the more likely chemical facto' 

which caused deposition to take place. Some conclusions as to wh*1 

the controlling factors maybe have been drawn from published liter? 

ture and others have been based upon experiments performed by tr 

writer and his coworkers.

MANGANESE IN SOIL AND PLANTS

Manganese is an essential minor element to plart growth. Becaus 

of its agricultural importance, a considerable amount of informatio 

on the manganese content of soils exists. ShawTbi (1952, p. 140 

observed that soils in the United States commonly contain from 0.00 

to somewhat over 1 percent manganese, but that some tropical soi 

contain as much as 15 percent. Thorne and Peterson (1954, p. 260 

noted that manganese deficiency most often occur? in soils at neutrc 

or slightly alkaline pH.

Earlier study by the writer (Hem, 1963) suggests that mangane? 

oxides are appreciably soluble in aerated water of pH as high ? 

9.0. This solubility suggests that manganese migH be readily leache 

from many soils. Iron is sometimes not available to plants when th 

soil pH becomes too high. Manganese, however, vould remain aval 

able well above the pH where iron deficiencies would begin to occu

Manganese is known to be adsorbed strongly in cation-exchan^ 

processes, but little information on the strength of adsorption c 

manganese compared with that of other ions seems to be availab?' 

Because the Mn*2 ion is smaller than the calciurr ion, however, an 

because both are divalent, manganese probably is held more strong! 

than calcium, and is therefore somewhat enriched in the adsorbe 

phase and impoverished in the solution phase.

The role of manganese in plant metabolism is not fully understood 

Tanner and others (1960) reported a chlorophyll synthesis in whic 

manganese is involved, evidently in a cyclic series of oxidations an 

reductions. Whether the main function of the element in plants is re 

lated to this synthesis or to other processes, the manganese con ten 

of plants is generally appreciable. Erkarna (1947) reported a cor 

tent of 13 to 1,280 mg of manganese per kilogram of plant leaves, pre 

sumably on a dry basis. Data of earlier investigators are in approxi 

mate agreement, and seem to suggest that the submerged fresh-wate 

plants are the highest in manganese content of the types examinee 

Oborn (1960, p. 208) reported that the iron content of aquatic plant 

was much greater than that of land plants. The behavior of mangar 

ese is similar to that of iron in this respect, as indicated by publishe* 

literature and work in progress (E. T. Oborn, written communicatior 

1963). Cannon (1960) reported an average (obtained from publishe
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B3

data on a wide variety of plant types) of 0.674 percent iron and 0.4815 

percent manganese in the ash of plant parts.

Oborn (1960, p. 204) and other investigators point out that the 

amount of iron and other metals present in certain plant parts is 

greater than in other parts. The roots of some varieties may contain 

relatively large amounts. However, there also is a tendency for some 

accumulation in the leaves. Besides the differences from one part of 

the plant to another, there are seasonal changes in metal content of 

the leaves and other plant parts. Some species of plants are particularly effective accumulators of certain elements and contain much 

larger amounts than would seem to be required for metabolic purposes. 

Cannon (1960) and Levering (1959) have reviewed literature relating 

to some of these plants. Bloss and Steiner (1960) found a positive 

correlation between the concentration of manganese in soil and underlying rock and the manganese content of leaves of the chestnut oak 

in Tennessee. They reported that the average manganese content of 

the ashes of leaves of trees in mineralized areas was 3.8 percent, and 

in nonmineralized areas was 2.1 percent. These values are considerably above the averages quoted by Cannon (1960) and suggest that the 

chestnut oak is a rather effective manganese accumulator.

Ljunggren (1951) found that spruce needles in the Vermland district of Sweden gave from 1.3 to 7.6 percent manganese in ash. The 

highest contents of manganese; were found in the needles of treer growing in abandoned river channels, where manganese had been deposited 

and was available to the tree roots in large amounts.

Evidently a large amount of manganese can be removed frona soils 

by forest vegetation; the amount returned annually to the soil by leaf 

shedding is also important. Levanidov (1957) estimated that 5 kg 

of manganese was returned annually to each hectare of the soil in birch 

forests in the south Ural region of Russia. About half the manganese in the leaves was readily soluble in water. If the area had an 

annual runoff of 10 cm and all the readily soluble manganese was carried in this runoff the water would have an average concentration 

of 2.5 ppm (parts per million) of manganese. Actually, most of the 

manganese would probably be adsorbed and retained by the soil and 

recirculated to the trees. However, if major runoff events occur during the season of the year when the ground in forested drainage basins 

is covered with freshly fallen leaves, rather large tonnages of nanganese may be carried off in flood waters.

MANGANESE IN RIVER WATER

Konovalov (1959) reported an annual mean manganese content of 

0.56 ppm for the Rioni River (whose basin contains extensive manganese deposits), 0.13 ppm for the Kura, 0.06 ppm in the Dvina and
B4 CHEMISTRY OF MANGANESE IN NATURAL WATER

Dnestr, and 0.04 in the Ob. These were the highest values reported 

for other rivers of the USSR the average amounts of manganese were 

less, ranging down to a few micrograms per liter. Data on manganese content of major rivers in other parts of the world are being 

obtained through a program, of the International Union of Geodesy 

and Geophysics. Preliminary results (W. H. Durim, written communication, 1963) show that manganese is always present, although 

concentrations greater than a tenth of a part per million are unusual.

Detailed records of the chemical composition of the water of many 

rivers in the United States have been published by tie U.S. Geological 

Survey. However, except for a few rivers affected by industrial 

wastes and mine drainage, these records contain no information on 

manganese concentration. The amounts present in water that is subject mainly to natural processes of manganese solution and deposition 

cannot be ascertained from these data.

In a few areas where the manganese content of stream water was 

high enough to be objectionable when the water was used for a public 

or industrial supply, some analyses have been mad°>. For example, 

Wiebe (1930) found from 0.04 to 0.13 ppm of manganese in the water 

of the Mississippi River at Fairport, Iowa, in a series of 10 samples 

taken between May 24 and October 19. More extensive future studies 

of minor constituents of water will be helpful in understanding more 

fully the natural processes by which manganese is brought into solution in river water.

The important role of vegetation in making manganese more readily 

available for solution in water of lakes and rivers seems to be well 

established by published information. Whether the dissolved manganese in river water is present as an organic compler, or in some other 

form, seems to be open to question. Theoretically, in aerated water 

whose pH is near 7.0, the maximum redox potential (as measured with 

a bright platinum electrode) that is likely to be reached will permit 

from a few tenths to more than one part per million of divalent manganese to remain in solution as Mn+2 at equilibrium. Consequently, it 

is not really necessary to postulate organic complexing, or stabilization 

of colloidal suspensions of oxide by organic solutes, in order to explain 

the presence of manganese in concentrations that probably are usually 

found in river water. In this respect the behavior of manganese is 

different from that of iron, for which solubility in aerated river water 

is normally very low.

Complexes of divalent manganese and organic anions are generally 

somewhat less stable than corresponding ferrous complexes. However, 

although references to the possible importance of organic complexes 

or chelates in stabilizing manganese against oxidation and precipita-
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B5

tion in natural water are common in published literature, few actual 

data on manganese complexes with natural organic materials have 

ever been obtained. SerdoboPskii and Sinyagina (1953) noted that 

manganese complexes with citric, acetic, oxalic, and humic acids did 

not occur in soils whose pH was below 7.8, but that almost all manganese was complexed at pH 10.2. The writer has observed that, at 

a pH above 9.5, rather stable manganese oxide colloids are readily 

formed in laboratory solutions, and the apparent complexing action at 

high pH may be related to organic stabilization of colloidal oxides. 

In any event, a pH of 10.2 is higher than normally occurs in natural 

water, although it may be reached in some soils.

The literature contains a few implications that partly oxidized 

manganese species in natural water may be important. Ingols and 

Wilroy (1962), for example, mentioned the possibility of organic 

complexes in which the manganese was present as Mn+3 in river water. 

Almost no data on stability of trivalent manganese complexes have 

been published. Nightingale (1959) used a triethanolamine complex 

of trivalent manganese for spectrophotometric determination of the 

element, but the complex required a pH near 13 to be stable. In the 

present study, the manganese oxide precipitated at pH above 9.0 contained some manganese that was not oxidized to the +4 state. Possibly, colloidal oxide suspensions formed at very high pH consist of 

manganese oxide or hydroxide in which not all the manganese i* at the 

+4 state. However, there is little reason to postulate forms of manganese other than the divalent state in explaining the presence of dissolved species. Until positive evidence of their existence is obtained, 

trivalent dissolved species need not be seriously considered.

STREAMBED MANGANESE OXIDE DEPOSITS

In certain areas of the world, deposits of manganese oxide, containing an admixture of iron oxide, form on rocks in the beds of streams. 

Such deposits occur, for example, in northern Maine (F. C. Canney, 

oral communication, 1962). in Summit County, Colorado (Theobald, Lakin and Hawkins, 1963), and in the Vermland district of 

Sweden (Ljunggren, 1953), and probably in other areas also. The 

processes by which manganese oxide is deposited in such areas are not 

fully understood. Possible mechanisms have been proposed by others, 

and further suggestions based on results of laboratory study will be 

presented later in this report. Deposition of manganese oxide from 

a dilute solution of manganous ions requires high pH and strongly 

oxidizing redox potential. The exact Eh and pH values required can 

be computed or predicted from graphic representations, such ?,s those 

of the writer (Hem, 1963), if dissolved manganese activity is known.

711-572 64 2
B6 CHEMISTRY OF MANGANESE IN NATURAL WATER

The deposits in the Swedish streams have been rather closely investigated by Ljunggren (1953,1955), and his data are more complete 

than any others known to the writer. Ljunggren's field studies were 

made during the summers of two successive years. On his first visit 

to the area, a considerable number of water samples were taken for 

determination of dissolved manganese, and pH measurements were 

made in the streams of interest. At that time, the streams were at a 

low stage of flow because there had been no rain for several weeks. 

On his second visit to the area, the streams were at a high stage of 

flow as a result of wet weather. No dissolved marganese data were 

obtained on this second visit, but pH measurements were made of the 

water in the streams.

Dissolved manganese concentrations at low flow v^ere generally less 

than 0.10 ppm, although one stream contained about 0.3 ppm. The 

pH of streams where manganese oxide had been deposited generally 

ranged from 6.5 to 7.0. When the streams were at a high stage of 

flow, they had lower pH's, some less than 6.0. Direct precipitation 

of manganese at pH 6.0 from a solution containing 0.10 ppm of manganese would require a redox potential of 0.70 volt?. which is higher 

than is normally attainable by the solution of oxygen from the air. 

Ljunggren reported that the oxide deposits seemed to be associated 

with a species of moss, Marsupella aquatica, which grew on the rocks. 

Whether the oxide was deposited as a result of influence on the system 

by the moss or because of microbiological growths associated with the 

moss was uncertain. The water was probably poorly buffered and the 

depletion of dissolved carbon dioxide in the vicinity of the vegetation 

during photosynthesis may have raised the pH of surrounding water 

to a level where the existing Eh would have been high enough to cause 

MnO2 to be brought down even from the very dilute solution. The 

deposits reported in Maine and Colorado had no obvious algae or plant 

association, but the possibility of microbiologic influence cannot be 

ruled out.

Several questions regarding these streambed deposits of manganese 

oxide remain unanswered. If they occur as a result of photosynthesis, 

the deposition should show seasonal and posssibly diurnal cyclic effects, and perhaps a good deal of re-solution of th°! deposits during 

the nongrowing season. It is also possible that deposition occurs 

mostly during certain periods when the stream water contains more 

manganese or has a higher Eh or pH than usual. Analysis of a 

series of samples of the stream water collected at different seasons 

might help solve these questions.

All the streambed deposits of manganese oxide which have been 

analyzed contain considerable amounts of feme iron, and apparently
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B7

the iron oxide can be predominant. This fact, which would seem to 

have an important bearing on the mechanism by which the precipitates 

are formed, is considered later in this paper (p. B38).

Ljunggren (1955) reported that the streambed deposits contained 

many minor constituents, apparently because the negatively charged 

surface of the oxide tends to adsorb cations from the solution in contact with it. Among the more abundant minor constituents were 

barium and strontium. The oxide was somewhat radioactive, probably owing to the presence of radium. Probably these ions were retained at the oxide surface by adsorption. Cobalt, nickel, and vanadium were also present but, inasmuch as these ions are nearly the same 

size as the managanese and iron ions in the oxide lattices, they may be 

incorporated in the structure of the oxides rather than being renoved 

from solution by adsorption.

The adsorption of barium by precipitates of manganese dioxide 

was briefly studied in the laboratory by the writer. The results suggest that the percentage of barium in solution that is adsorbed is a 

function of pH and that below pH 7.0 the adsorption is weak. The 

degree to which the oxides are enriched in barium may be- an approximate index of the pH at the site of precipitation of the oxides. The 

fact that the deposits studied by Ljunggren contained a considerable 

amount of barium may be additional evidence supporting the idea that 

the moss associated with the deposits raised the pH of water adjacent 

to the plants and thus promoted the deposition of manganese.

MANGANESE IN LAKES AND RESERVOIRS

The highest concentrations of dissolved manganese in surface runoff 

are likely to occur in forested areas where the soil and underlying 

rocks are manganese bearing. In recently glaciated regions, drainage 

patterns are generally poorly formed and there are many lakes. 

Where there are lakes in regions favorable for manganese to be present 

in surface water, many interesting phenomena have been reported.

Because stable thermal stratification generally occurs in lakes during summer months, the supply of oxygen in the near-surface water 

derived from the atmosphere may not be circulated into the deeper 

water. In a lake supporting appreciable amounts of biological 

growths, the deeper water has an oxygen demand because organic 

debris is attacked by oxidizing organisms at and near the bottom. During the period of stratification the dissolved oxygen of the deeper water 

may thus be depleted and the environment made relatively reducing 

and favorable for increased solution of iron and manganese. Increases 

in iron and manganese content with depth in lakes during the summer 

stratification period are well documented. Hutchinson (1957, p. 809),
B8 CHEMISTRY OF MANGANESE IN NATURAL WATER

for example, reports observations on Linsley Pond, Conn.;, showing 

manganese in excess of 3 ppm at moderate depth, and states that manganese concentrations can be expected to increase at conditions less 

reducing than are required for solution of iron. In this particular 

set of observations, the iron in solution exceeds the manganese content 

near the bottom of the lake.

Mixing of the lake water, which generally take7 place seasonally, 

raises the redox potential of the deep water and cr,n generally be expected to cause iron and probably manganese to precipitate. Because 

outflow from a lake is generally of the relatively oxidizing water at the 

surface, depleted in manganese by biochemical action, the lake can be 

considered a site for accumulation of the manganese brought in by the 

surface runoff. This managanese contribution includes dissolved 

Mn+2 species that would be detected in a chemical ar alysis of the water 

and also some contributions that might not normally be detected. For 

example, some manganese might enter the lake in suspension in the 

surface water, either adsorbed by sediments or contained in organic 

debris. There might also be an important amount of manganese 

brought in by ground-water inflow. Although the general long-term 

trend will favor manganese accumulation, the changing conditions 

within the body of water will sometimes favor higl concentrations ir 

solution and at other times low concentrations.

Where accumulation of manganese is rapid, a condition usually accompanied by rapid iron accumulation, the lake deposits may constitute a source of ore. Ljunggren (1953, 1955) described bog ores 

consisting of iron and manganese in the Lakes Tisjon and Grycken 

in Dalecarlia, Sweden. Some of the lake-bottom areas were particularly favorable for manganese deposition, and MnO2 contents as high 

as 30 percent were reported for some of the deposited oxides. The 

deposits richest in manganese occurred in shallow water at depths 

between 1 and 2 meters, where the lake beds were permeable gravel 

and it was suggested that ground water especially rich in manganese 

may have entered the lake in these areas. The Eh and pH of the lake 

water probably would have favored oxidation and precipitation o: 

manganese from this inflow. Ljunggren believed micro-organisms 

were involved in the oxidation process, and the deposits containe^" 

considerable amounts of organic matter.

Deposits of manganese oxide on pebbles in lakes of Nova Scotia 

were mentioned by Kindle (1936) who attributed them to the actior 

of algae. Krotov (1950) has discussed iron and manganese deposition 

in lakes in other parts of northern Europe. The topic is discussed 

in a general way by Krauskopf (1957) who has summarized processes 

which may lead to separation of iron and manganese in nature.
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B9

Small, .amounts of manganese in public water supplies are highly 

>bjectionable. The limit of 0.05 ppm for manganese given in th^ suggested U.S. Public Health Service Standards (1961) for drinking 

vater may be too high for some industrial water. Griffin (1960) 

suggests that as much as 0.02 ppm manganese may be tolerated in the 

iverage domestic supply. Because the oxidation of manganese is a 

rather slow reaction under conditions generally reached in water treatnent processes, the reaction may not be completed in the time the water 

;an be held in the treatment facilities, and water containing small 

imounts of manganese can give rise to deposits of oxide in waterlines 

ind other parts of the distribution system after the water has been 

filtered and chlorinated. The deposits of manganese oxide m^y be 

lislodged or partly redissolved at later times to give sudden increases 

in manganese content in water from the mains at some places.

Public water-supply reservoirs on streams, like natural lake?, can 

be expected to serve as accumulating basins for manganese. For reasons that have already been outlined, some of the accumulated manganese obviously could be released to solution at times; as this release 

occurs, occasionally water in the deeper parts of the reservoir or sometimes even the whole lake may contain objectionable concentrptions 

of manganese. Intakes for water-supply systems are usually located 

at a considerable depth below the reservoir surface, to avoid wide temperature fluctuations, and for other reasons. This placement of ir takes 

increases the probability of obtaining manganese-bearing water from 

"-line to time.

Although it seems evident that a few hundredths of a part per mil- 

'ion, or more, of manganese might very often be found in water near 

^he bottom of a reservoir, when the concentration of manganese of 

the inflow and of near-surface water in the reservoir is much smaller, 

this fact is not widely recognized.

Published literature indicates that the owners of water-supply 

"eservoirs did not expect problems of excessive manganese in the water. 

"Vhen such problems did occur, various theories were suggested to 

explain the presence of manganese. Many of these speculations were

 erroneous. Because the inflow to the reservoir did not have a high 

manganese content at the time the water withdrawn did, it w?-s assumed, for example, that the manganese originated within the reser- 

^oir. It was sometimes assumed the problem would disappear after 

a reservoir had been used for a time. Unfortunately, there is usually 

no sound reason to expect this to occur. On the contrary, the general 

' endency will be for manganese accumulations to increase with time. 

Although sources of manganese probably differ from place to place

- vhere the problem is found, the general principles which govern the 

occurrence of manganese in water-supply reservoirs include:
BIO CHEMISTRY OF MANGANESE IN NATURAL WATER

1. Biological growths remove manganese from aerated water and 

precipitate it to the reservoir bottom.

2. Biochemical action in bottom mud and adjacent water may cause 

reducing conditions which periodically will return manganese 

to solution.

3. In regions where manganese oxides are presert in considerable 

amounts as constituents of soil and of outcropping rock, the 

solution of manganese from submerged surfaces within the 

reservoir may be important. However, the reservoir will tend to 

accumulate any manganese brought to it in solution as the result 

of leaching of soil and organic debris. The rate of accumulation 

would be rapid in such an area.

DESERT VARNISH

Another form of manganese and iron-oxide deposits on rock surfaces which has received a good deal of attention over the years is the 

dark-colored coating called desert varnish which occurs on exposed 

rock surfaces and on surficial gravels in arid regions. There has been 

considerable debate as to the role of plants and other organic factors. 

The earlier investigators seemed generally to favor the view that the 

deposits are brought about largely, if not entirely, by primitive plants 

such as lichen. Some of the more recent investigators, however, have 

minimized the importance of the biochemical factor. Engel and 

Sharp (1958) in an extensive study of desert varnish in the Mohave 

Desert of California concluded that the deposits on the gravel of 

"desert pavement" came from underlying soil aid were deposited 

from solution, and, that such coatings can be formed fairly rapidly (in 

one place \vithin 25 years). Usually, however, the deposition process 

is much slower. Deposits on larger areas of bare rock were attributed 

to weathering solutions from rainfall, runoff, and dow, and to precipitation of airborne particulate matter. Although Engel and Sharp 

do not entirely dismiss the idea that lichen can play a part in the formation of desert varnish, they apparently do not believe that organic 

processes are a major factor. Deteriorated varnish coatings were 

noted by Engel and Sharp. In most instances, coatings that were 

being destroyed were flaked off as a result of weathering of underlying 

rock. Apparently the coating represents a residue resistant to weathering and probably is mechanically removed where water is more 

abundant.

Hunt (1961) has expressed the opinion that desert varnish deposits 

have not been actively forming in most of the Western United States 

during the past 2,000 years. This conclusion is basei on archeological
DEPOSITION AND SOLUTION OF MANGANESE OXIDES Bll

evidence for the most part, and it is suggested that desert varnish 

forms readily only when plentiful supplies of moisture are available. 

The chemical environment for deposition of manganese anc1 iron 

oxides on streambed rocks probably is much different from that prevailing in the occasional film of water on the surface of a rock on the 

land surface of a desert. However, the deposits have some points of 

resemblance in their minor-element content. These resemblances are 

probably related to the surface characteristics of ferric hydroxide and 

manganese dioxide. The anlyses of Engel and Sharp for minor constitutents of varnish are more complete than the data for the st~eambed deposits. Both, however, show strong enrichment of barium and 

somewhat similar amounts of cobalt, nickel, and vanadium. Desert 

varnish evidently contains more barium than the streambed deposits 

do. A rather high pH in the thin surface film of water on the rocks 

is certainly not unlikely.

DEEP-SEA NODULES

Manganese oxide is a principal constituent of nodules that occur at 

the bottom of the oceans in many places. These nodules and th^ factors causing them to be formed have been extensively studied. A 

recent report by Goldberg and Arrhenius (1958) suggests that the 

nodules form through accretion of oxidized manganese and iron on 

a suitable surface. The oxidation is attributed to effects of dissolved 

oxygen in the sea water. Redox potentials and pH in deep ocean water 

apparently are generally favorable for manganese precipitation. Willis and Ahrens (1962) have discussed the composition of the noiules. 

The effectiveness of manganese oxide in adsorbing metal cations is 

strongly indicated by these data.

SEPARATE PRECIPITATION OF IRON AND MANGANESF

Although the manganese oxide deposits considered here generally 

contain a considerable proportion of iron, the predominance of manganese in many can be explained only by some mechanism favoring 

the precipitation of manganese, rather than iron, from solution-' containing both elements. It is difficult to postulate such a mechanism 

for the two elements' chemical behavior is somewhat the sam°, A 

solution of the two might be expected to yield a precipitate of both 

iron and manganese oxides when subjected to an increased Eh and 

(or) pH. Because iron is the more readily oxidized, however, the 

iron would tend to be removed first, and if a very high pH and redox 

potential are reached, the manganese will be oxidized and precipitated 

also. Krauskopf (1957) considered possible mechanisms for d°position of relatively pure manganese oxides and concluded that th0* pre-
B12 CHEMISTRY OF MANGANESE IN NATURAL WATER

cipitation of iron first and later precipitation of manganese from 

the same solution was a reasonable possibility. However, he pointed 

out that this theory implies that a considerable volume of deposited 

iron oxide should be somewhere near all or most manganese oxide 

deposits and that this condition does not always seem to exist.

The lakebed and streambed deposits of manganese oxide which 

have been considered here are actually all rather rich in iron, and 

some actually contain more iron than manganese. Processes of coprecipitation, perhaps with some degree of enrichment with manganese, seem adequate to explain them.

In surface streams the pH will normally be near neutrality, and 

Eh will seldom exceed 0.50 v except where unusual biochemical or 

other factors intervene. Stability field diagrams for the two elements (Hem, 1960a, 1963) show that manganese is soluble to the 

extent of more than 100 ppm at pH 7.0 and Eh 0.40 v, whereas iron 

under these conditions can be dissolved only to the extent of about 

0.0001 ppm. This striking difference in solubility suggests that in 

weathering and related processes the two elements can be expected 

to show different behavior, that the change from an environment 

favoring iron deposition to one favoring manganese deposition may 

be a major one, and that a considerable separation in space may be 

involved. The presence of such an effective means for separating the 

elements readily explains the occurrence of some nerrly pure deposits 

of manganese.

It is possible to precipitate manganese from a solution in which 

iron apparently is present in an amount comparable to or in excess 

of the amount of manganese present in solution. Under some conditions the apparent iron concentration may be mostly in an unreactive form. As pointed out elsewhere by the writer (Hem and Cropper, 1959), the iron reported in most analyses of river water must 

usually represent either a colloidal variety of ferric hydroxide or 

iron held in an organic complex. Because of the very low solubility 

of iron in waters having the usual pH and Eh of surface streams, 

the amount in ionic solution must be small. The colloidal particles 

of ferric hydroxide have a negative surficial char^ at neutral pH, 

and the colloid is relatively stable and is not attracted to most mineral surfaces. Neither the colloidal hydroxide no** the organically 

complexed iron would be able to participate actively in chemical 

reactions, and could, therefore, be transported readily in the flowing 

water. Manganese, on the other hand, is likely to be in solution as 

Mn+2, or possibly as the inorganic complex ions MnHCO3+1 or 

MnSO4°, because these forms are favored in the near-neutral oxidizing solutions of river water. Manganese oxide colloids are favored
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B13

v ly by a very high pH. Organic complexes of manganese arc,, in 

^neral, less stable than corresponding ferrous complexes (Hem, 

H3).

In view of these factors, then, manganese should be present in r;ver 

^ater in a form susceptible to chemical reactions that can cause 

"<:.position as oxide even though the apparent content of iron in the 

"ater is comparable to or greater than the manganese, and it is likely 

c be present in a form readily available to aquatic life. The colo : dal iron would not be affected by such reactions, being already 

.ridized. The organic complexes of ferrous iron appear to slow the 

ridation of the iron greatly (Shapiro, 1958; Hem, 1960b), and thus 

c prevent it from being precipitated.

MANGANESE IN GROUND WATER

The chemical behavior of manganese in ground water can be preicted approximately by applying theoretical chemistry. The genr*»l principles and the relationships which are important here are 

' vered in detail in an earlier paper (Hem, 1963) .

In ground water systems the source of manganese for solution may 

': from minerals in which manganese is a relatively minor constituT t. In evaluating such systems, chemical equilibrium involving simJ<H manganese minerals, such as oxide, carbonate, or sulfide, cannot 

v: assumed. Equations and equilibrium diagrams may, therefore, 

 f: be directly applicable unless the ground water is precipitating, 

v is in contact with, one of the simpler manganese minerals at or 

< ar the point at which a water sample may be obtained.

Thermal springs which yield water containing from a few tenths 

c several parts per million of manganese are fairly common. Some 

f these waters are apparently associated with deposits of manga- 

'°,e oxide. Others may possibly be associated with rhodochrosite 

] ^nCO3 ) below the surface, for this mineral is sometimes found in 

^in deposits. A few analyses were available in the U.S. Geological 

>^rvey files in Denver which represent samples where the pH was 

i^asured in the field at the time of collection. These analyses v^ere 

ramined for adherence to the equilibrium :

MnCO3c + H+1?±Mn+2 + HCXV1,

activity coefficient and complexing corrections were used on 

IT reported concentration data. Of a group of five analyses so 

jammed, the total calculated manganese solubility checked fairly 

r->sely with the observed manganese concentration in only two. One 

j- these analyses represented the drainage flowing from an al ^.nmine near Silverton, Colo., and the other was for Pinkerton

711-572 64 3
B14 CHEMISTRY OF MANGANESE IN NATURAL WATER

Springs, a thermal spring in La Plata County, Colo. The other 

waters were all below saturation with respect to rhodochrosite by a 

wide margin.

No very specific conclusions can be drawn from th^se results. More 

information as to mineral species that may exist in the areas from 

which the water samples came would be needed before much use of 

the equilibrium equations with field data would be possible.

Much manganese-bearing ground water probably has been in contact with manganese oxide. Alluvium deposited by streams may contain sand and gravel incrusted with manganese oxide. Ground water circulating through this alluvium would dissolve some manganese. 

In areas where vegetation is plentiful and the surface drainage pattern 

includes many lakes, accumulations of manganese will form, probably 

as oxide, and at a later period these may become available to circulating ground water. This kind of condition seems to be common 

in the part of the United States which was subjected to glaciation in 

Pleistocene time.

Manganese commonly occurs in ground water in alluvium. Gidley 

(1952) described a number of collector-type wells along the Ohio 

River between Wheeling and Huntington, W. Va. Two wells at 

Parkersburg, W. Va., yielded water that contained 1.8 ppm of manganese. These wells obtained their water through horizontal pipes 

extending into gravel below the bed of the river. A Ithough the manganese in the water from these wells evidently was not expected, manganese can be expected in the alluvium, especially in this area. The 

Ohio at Newell, W. Va., contained an average of 0.36 ppm manganese 

during the year ended September 30,1957, and 0.03 ppm for the same 

period at Florence, Ind. The manganese entered as a result of acid 

mine drainage and other pollution along the Ohio r.nd its tributaries 

upstream. The analytical data for stations on the river in this region 

were published by the U.S. Geological Survey (196C), and the results 

for 1957 as well as those for the 5 previous years show that manganese 

is being lost from solution and probably accumulates in the bottom 

sediment, especially in ponded sections.

The sediment along the Mississippi River in some places also yields 

water that is high in both manganese and iron. WeT-known instances 

of the occurrence of manganese in wells adjacent to the river occur 

at Brainerd, Minn., and La Crosse, Wis.

LABORATORY STUDIES OF MANGANESE DEPOSITION

Experiments were designed to determine more closely some of the 

conditions which may control the deposition of manganese oxide on 

the surfaces of certain rock minerals.
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B15 

NATURE OP SANDS UTILIZED

Experiments were performed with two different types of rand 

' hich were available in the laboratory. One of these (sand 1) was 

i. purified silica sand obtained from the J. T. Baker Chemical Co. 

[his material was almost pure SiO2 and had uniformly well-rounded 

i articles. The range of particle diameters was from 0.8 to 0.25 mm 

or 95 percent of the material, and none of the particles had diameters 

oss than 0.1 mm. The cation-exchange capacity of the sand was- too 

mall to be reliably measured.

The second type of sand hereafter called "sand 2" was a Monterey, 

"Hlif., beach sand whose particles were moderately rounded and well 

orted, 85 percent having diameters between 0.8 and 0.25 mm. Fone 

rere larger than 0.8 mm and none smaller than 0.10 mm. The aproximate composition of the sand was determined by means of stainr o; technique and a grain count as:

Percent

Quartz ____________________________________ 45 

Orthoclase feldspar____________________________ 21 

Plagioclase feldspar ______ __ _____ 34

[he cation-exchange capacity of the sand was determined by eqnilirating a weighed part with a 1.0 molar solution of MnCl2 . The sand 

v,s washed with ethyl alcohol to remove adhering manganese chloi de; the adsorbed manganese was removed with 1.0 molar HNO3 , and 

[i termined by the periodate procedure. The exchange capacity deterrined in this way was 0.66 meq (milliequavalent) per 100 g.

CATION-EXCHANGE ADSORPTION OP MANGANESE

The extent to which manganese behavior in natural water may 

H'. influenced by cation-exchange reactions has a considerable potential 

i?nificance, and some experiments were made to determine the 

t^ength of adsorption of manganese on the sand surfaces as compared 

c the strength of adsorption of calcium.

A convenient-sized exchange column was prepared of sand 2 

h rough which solutions could be passed under conditions of unsp.tu- 

£ ted flow. About 200 g of sand was placed in a glass tube 60 cm longrd 2 cm in diameter. The upper end of the tube was enlarged 

o a diameter of 4 cm over a length of about 10 cm. This enlarged 

x rtion served as a reservoir for solution being passed through the 

clumn. The lower end was closed by means of a stopcock and the 

and was retained in the tube by a fritted glass disc just above the 

topcock.

The volume of the 200 g of sand in the column when dry was 128 

.:", A volume of 47 ml of water was required to saturate the 200 g of
B16 CHEMISTRY OF MANGANESE IN NATURAL WATER -

sand, and no change in volume of sand was observed. The void space 

in the column, therefore, was 47 ml or 37 percent of the sand volume, 

and the sand-water mixture was 23.5 percent water by weight.

The column was supported in a vertical position and solutions were 

passed through it by gravity. This arrangement permitted unsaturated or intermittently saturated flow. When solutions were passed 

through the column, a flow rate of about 200 ml p?r hour was generally attained after the water had flushed out mo?* of the air from 

the sand column; a head of about 5 cm of water was maintained above 

the top of the sand surface. Slower rates were obtainable by partly 

closing the stopcock at the bottom of the column; however, the rate 

of water movement was not closely controlled.

Before it was put in the column, the sand was saturated with sodium 

by agitating it in a 1.0 molar solution of sodium chloride. Excess 

salt was removed by washing with distilled water, and the sand was 

dried on a steam bath before filling the column. This procedure was 

intended to remove any adsorbed divalent catior*? from the sand 

surfaces.

A solution of known calcium and manganese content was prepared 

from stock solutions of calcium chloride and manganous chloride and 

was poured into the column. All the effluent was collected in successive 100 ml parts, and the dissolved calcium and manganese contents in each part were determined. When the concentrations in the 

effluent became the same as those of the solution being put into 

the column, the flow was stopped and a new solution containing 

different amounts of calcium and manganese was made up. The 

new solution was used to repeat the experiment. Generally, between 

2 and 3 liters of solution had to be passed through the column before 

equilibrium was reached. Solutions used had initial calcium concentrations near 100 ppm and near 10 ppm and O.rMn mole ratios 

ranging from 250 to 1.3. One solution containing- 1.0 ppm Ca and 

10 ppm Mn also was used. All the solutions had a pH of about 

6.5. No loss of manganese by oxidation would be anticipated in these 

solutions and none was observed.

From the calcium and manganese content of the effluent, the 

amounts of these ions absorbed in the column w^re calculated, in 

milligrams and in milliequivalents. The results arc* given in table 1.

Amounts of calcium and manganese adsorbed are generally between 

0.3 and 0.4 meq per 100 g of sand, which is about yhat would be expected, considering the rather dilute nature of these solutions. When 

the two ions are present in near-equivalent amount? (solution 5) the 

ratio of adsorbed calcium to adsorbed manganese is simewhat less than 

the mole ratio of dissolved calcium to dissolved manganese. This
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B17

relationship indicates that the sand adsorbs manganese somewhat more 

than calcium.

PABLE 1. Equilibrium concentrations of calcium and manganese in solution 

and adsorbed on feldspathic sand

Mu- 

*ion

.....

.....

Mn 

(ppm)

0.50 

1 

10 

10

10 

1

.65

CM. 

(mol per 1)

9.1XWH 

1.8XUH 1. 8X10-* 

1. 8X10-*

1.8XMH 1.8X10-*

1.2X10-'

Ca 

(ppm)

92 

92 

93 

1.0

9.2 

9.2

9.2

Cc. 

(mol per 1)

2.3XKH 

2. 3XUH 2.3X10-*

2 HV1O-S

2.3X10-* 

2.3XKH 2.3X10-*

Cc.:CMa

250 

130 

13 

.14

1.30 

13

19

XMO

(mg)

0.95 

1.55 

4.67 

12.60

8.80 

4.66

4.10

XM*

(meq)

0.035 

.056 

.170 

.459

.320 

.170

.149

-STc.

(mg)

13.7 

13.6 

12.8 

3.3

7.1 

9.4

9.3

Xc* 

(meq)

0.685 

.680 

.640 

.165

.355 

.470

.465

A"c«:^Mn

20 

12 

3.8 

.36

1.1

2.8

3.1

P

ll.fS

[2.4S

k

1.1

1 1

The equilibrium reached in the exchange-column experiment can 

:T written

CaX+Mn+2=Ca+2+MnX

rhere X represents the cation-exchange complex. The mass-law 

Equilibrium expression:

~+*

cannot be strictly applied to the results, however. Apparent equilibrum constants for the seven sets of data in table 1, based on nrolar 

concentrations in solution and moles of ions adsorbed, range from 

D.39 to 12.9.

The results of the experiments, however, do fit the empirical eiuati on of Rothmund and Kornf eld :

\ -here C terms are molar concentrations of dissolved ions and X terms 

a re equivalent quantities of adsorbed ions. The p and k terms- are 

empirical constants.

In figure 1 where the ratio (7Ca : Oun is plotted as abscissa an<? the 

r -"»tio .Zca " -^Mn is plotted as ordinate, the data yield two straight l; nes, 

cne for the solutions containing 92 ppm Ca, and one for the solutions 

c 'mtaining 9.2 ppm Ca. The slopes of these lines provide a basir for 

c imputing the values of the exponent p in the equation. The value of 

If is the same for both sets of data. Because only one solution containing 1.0 ppm Ca was studied, there is no line for this concentration.

Because the experimental data are consistent with the idea that 

rianganese losses and gains are associated with ion exchange rather 

t"-an oxidation, it would seem that very little manganese oxidation
10.0

-X"ca

 ^Mn

1.0

0.10

0.1

0

Cc

0

a 1.Ofsprr

C

-rt

1.0

<1i^H

9^*P

C

C

^

. 'm ^t

0 «**

10.0 

Ca

Mn

-c

.^"

x^*"

G»

^P

<f ^

]

-S

00

s

^

LOO

FIGURE 1. Relationships among dissolved concentrations and adsorbed quantities of calcium and manganese in adsorption column containitg sand 2.

could have taken place. The experiments were run at a pH of 6.5, at 

which prior work showed the oxidation rate would have been very 

slow.

Figure 1 provides a basis for predicting certain features of the 

behavior of manganese in aquifers or in river or lal*e water in contact 

with sediment. Increases in calcium content of water in contact with 

sediment on which exchangeable manganese is pr^ent will tend to 

bring manganese into solution. The action of bacteria and the decay 

of organic matter in submerged soil often have the effect of bringing 

calcium into solution in immediate contact with the soil particles. This 

action could be a mechanism for increasing the manganese content 

of bottom water in reservoirs.

The results (fig. 1) further suggest that the ratio of dissolved 

manganese to calcium in ground water may be controlled by adsorption reactions in the aquifer material.

OXIDATION EQUILIBRIUM EXPERIMENTS

A volume of 3 liters of demineralized distilled water was brought 

to a manganese concentration of 10 ppm by adding manganese chloride. The pH was adjusted to 8.0 with sodium hyc'roxide. The solu-
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B19

tion was then transferred into six 500 mi-capacity polyethylene bottles. 

Five of these contained a measured quantity of sand 1, which had been 

thoroughly washed and dried beforehand. Amounts of sand present 

in the bottles ranged from 10 to 100 g. The remaining bottle contained 

no sand. The solutions were allowed to stand in the stoppered bottles. 

From time to time, measurements of pH and Eh were made on each 

solution and a part was removed and filtered through a plastic membrane filter having 0.45 /u, diameter pores. The amount of manganese 

remaining in solution in the filtrate was determined by the periodate 

procedure. The results of these observations are given in table 2.

TABLE 2. Manganese concentration, pH, and Eh of water in contact with

silica sand

Solution

1

t

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

4

6 _. __ _ ..... . .

Weight 

of sand

(grams)

10

30

50

75

100

Time

(hours)

0

17 

41 

137

440 

0

17 

41 

137 

440 

0

17

41 

137 

440 

0

17 

41 

137 

440 

0

17

41 

137 

440 

0

17 

41 

137 

440

Manganese 

concentration (ppm)

10

9.6 

9.6 

10 

9.2 

10

9.4 

10 

10 

9.6 

10

9.7 

10 

10 

9.2 

10

9.5 

10 

9.2 

9.0 

10

9.2 

9.8 

8.8 

9.6 

10

9.3 

10 

9.4 

9.4

PH

8.0

7.2 

7.0 

7.0

6.7 

8.0

7.2 

7.1 

7.0 

7.4 

8.0

7.5 

7.1 

7.0 

7.3 

8.0

7.1 

7.0 

6.8 

7.1 

8.0

7.0 

6.9 

6.7 

7.1 

8.0

6.8 

6.6 

6.4 

6.8

Eh (

Observed

0.36 

.39 

.40

.50

.37 

.40 

.43 

.46

.38 

.41 

.44 

.49

.41 

.42 

.45 

.48

.42 

.42 

.46 

.47

.43 

.44

.48 

.50

^

Estimated

0.43 

.47 

.47 

.53

.43 

.46 

.47 

.42

.38 

.46 

.47 

.43

.45 

.47 

.51 

.46

.47 

.49 

.53 

.46

.50 

.55 

.57 

.51

In an earlier experiment (Hem, 1963) with a similar set of solu- 

'ions, except that no sand was present, oxidation took place at a pH 

of 8.0 or more, as evidenced by some loss of manganese, and equilib4um was attained after the solution had stood for a time. The 

chemical system represented by the data in table 2 can be visurlized 

as involving divalent and oxidized manganese, dissolved oxygen and 

hydrogen ions, which will be related to each other at equilibrium by 

' he Nernst equation. Because the solutions are dilute and interfering 

ions are absent, conditions may be evaluated with a stability-field 

diagram such as that developed in the writer's previous research in
B20 CHEMISTRY OF MANGANESE IN NATURAL WATER

manganese chemistry (Hem, 1963). The value of Eh obtained from 

the graph for the observed dissolved manganese anc1 pH is entered in 

table 2 as "estimated Eh."

When equilibrium is reached in the solutions, tli^ calculated and 

observed Eh values should be in close agreement. Earlier work has 

shown that the experimental error in measuring J5h can be rather 

large, and this method of deciding whether or not equilibrium is attained leaves something to be desired. It is, nevertheless, a useful 

technique for studying oxidation reactions.

Although initial Eh was not measured, from the Eh-pH diagram 

it is evident that the initial solution was unstable and would tend to 

lose manganese by oxidation. Only a small amount of oxidation, 

however, is enough to lower the pH of the unbuffrT*ed solution substantially. When the pH goes down, the oxidation reaction almost 

stops but a considerable time is required to arrive at a state of 

equilibrium at the final pH.

The results in table 2 indicate that the above course of events was 

followed. The total loss in manganese, which wai1 small in all the 

solutions, occurred in the initial 17 hours, while t>e solutions were 

standing overnight. The decrease in pH also occurr*! almost wholly 

in this period. The Eh measured in most samples at 17 hours, however, was below the equilibrium value. The slow return to equilibrium 

which followed entailed mostly increases in the ret'ox potential and 

some minor fluctuation of pH.

Minor differences in behavior, possibly owing to presence of sand, 

could be discerned. Amounts of manganese lost we~e about the same 

in all solutions, but the solutions with more sand had higher Eh's. 

There was no definite indication in any of these data that the sand 

had an important effect on the rate or amount of oxidation occurring, 

and no further experiments were made with the quartz sand.

The next series of experiments was made using s^nd 2. The sand 

was prepared for use by stirring it in a 1.0 molar sodium chloride 

solution for several hours, the mixture was allowed to stand overnight, 

and the excess salt was washed from the sand with demineralized 

distilled water until the washings were chloride-free. This procedure 

was intended to fill exchange positions on the sand surface with sodium. 

Presumably, these positions would later be occupied by manganese.

Data given in table 3 were obtained in the first experiment with sand 

2. The initial conditions imposed on the solutions reported on in 

table 3 were different in two respects from those effecting the data 

in table 2: (1) The initial pH in the solutions was 7.7 instead of 8.0; 

at this lower pH the rate of manganese oxidation would be lower. 

(2) Sand 2-had some adsorption capacity for cations which might 

be expected to remove manganese from solution.
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B21

"ABLE 3. Manganese concentration, pH, and Eh of unbuffered solutions in 

contact with feldspathic sand

Solution

..

E

10............ ........... .

r.. - --- ---.

12 .-.....-...............

Weight 

of sand 

(grams)

None

10 

30 

50 

75 

100

Time 

(hours)

0 

18 

42 

114 

210 

354 

0 

18 

42 

114 

210 

354 

0 

18 

42 

114 

210 

354 

0 

18 

42 

114 

210 

354 

0 

18 

42 

114 

210 

354 

0 

18 

42 

114 

210 

354

Manganese 

concentration (ppm)

10 

9.8 

10 

10 

9.7 

9.9 

10 

9.2 

9.0 

8.4 

7.4 

7.9 

10 

8.3 

7.6 

6.4 

6.3 

5.9 

10 

7.5 

5.7 

4.9 

3.9 

3.8 

10 

7.0 

5.2 

4.3 

3.5 

3.8 

10 

6.2 

3.8 

2.8 

2.4 

1.4

PH

7.70 

7.77 

7.60 

7.25 

7.09 

7.07 

7.70 

7.59 

7.11 

7.23 

7.24 

7.28 

7.70 

7.41 

7.13 

6.77 

6.74 

6.95 

7.70 

7.34 

7.05 

6.87 

6.95 

6.92 

7.70 

7.43 

7.20 

7.00 

7.13 

7.22 

7.70 

7.27 

7.01 

6.92 

6.96 

6.98

Eh (volts)

Observed

0.47 

.41 

.47 

.44 

.45 

.49 

.47 

.43 

.49 

.45 

.46 

.47 

.47 

.44 

.50 

.47 

.48 

.48 

.47 

.45 

.50 

.47 

.47 

.49 

.47 

.44 

.49 

.46 

.46 

.47 

.47 

.46 

.50 

.46 

.47 

.48

Est'niated

0.34 

.36 

.43 

.46 

.46

.37 

.46 

.44 

.44

.45

.40 

.45 

.52 

.53 

.49

.42

.48 

.51 

.51 

.51

.40 

.45 

.50 

.48 

.46

.44

.50 

.52 

.52 

.52

The results obtained with sand 2 show that the solutions in contact 

r~ith this sand decreased considerably in dissolved manganese over the 

? 54 hours of the experiment. The rate of loss of manganese was rapid 

for the first few days and became slower thereafter. Figure 2 s^ows 

1 he rate of loss of manganese in four of the solutions. Except for the 

solution having 100 g of sand, the dissolved manganese in the solutions did not change significantly after about 200 hours.

The loss of dissolved manganese from these solutions can b^ attributed either to adsorption by the cation-exchange surfaces of the 

?and or to oxidation of manganese, or to both. Any appreciable oxidation effect would be an indication of catalytic action by the sand surface, for no manganese loss occurred at this pH in the absence of sand. 

Because the pH changes in the unbuffered solution are small, it seems 

probable that there was only a small amount of oxidation, and that 

most of the effects observed in these solutions were the result of ion 

exchange.
B22 CHEMISTRY OF MANGANESE IN NATURAL WATER

EFFECTS OF SAND ON OXIDATION FATE

Iii order to obtain more information as to the relative importance 

of adsorption and oxidation in the loss of manganese from solutions 

in contact with sand 2, several additional test rung were made. The 

first set of solutions was made up as before, and sodium bicarbonate 

was added to give a concentration of 120 ppm bicarbonate. The 

bicarbonate was added in order to stabilize the pH in these solutions 

and to prevent it from decreasing as a result of oxidation and hydrolysis of manganese. The results of this series of tests are in table 4.

TABLE 4. Manganese concentrations, pH, and Eh of solutions containing 

bicarbonate ions in contact with feldspathi') sand

Solution

13. 

14-._-. . 

15-..... - -.-..- 

16 

17-- 

18

Weight 

of sand 

(grams)

10

30

50

75

100

Time 

(hours)

0

7 

24 

48 

97 

216 

0

7 

24 

48 

97 

216 

0

7 

24 

48 

97 

216 

0

7 

24

48 

97 

216 

0

7 

24 

48 

97 

216 

0

7 

24 

48 

97 

216

Manganese 

concentration (ppn)

10

10 

10 

9.7 

10 

9.5 

10

8.6 

8.7 

8.1 

8.1 

6.9 

10

7.9 

7.4 

6.5 

5.1 

3.2 

10

7.4 

6.0 

4.9 

3.6 

2.1 

10

6.1 

4.4 

2.5 

2.3 

0.60 

10

5.7 

3.8 

2.2 

1.2 

.25

pH

8.03

8.07 

8.12 

8.29 

8.43 

8.60 

8.03

7.95 

8.07 

8.18 

8.38 

8.39 

8.03

7.88 

7.92 

7.89 

8.18 

8.17 

8.03

7.82 

7.78 

7.75 

8.07 

7.90 

8.03

7.75 

7.84 

8.04 

8.06 

8.09 

8.03

7.71 

7.70 

7.93 

8.10 

8.02

Eh (volts, 

observed)

0.40

.40 

.38 

.39 

.40 

.37 

.40

.41 

.39 

.40 

.40 

.39 

.40

.42 

.40 

.43 

.41 

.41 

.40

42 

.40 

.43 

.42 

.43 

.40

.43 

.40 

.41 

.41 

.41 

.40

.43 

.41 

.42 

.41 

.41

The pH of the solutions in this experiment fluctuated about 0.3 pH 

unit (except for the solution containing no sand). Tils compares with 

a fluctuation of nearly a full pH unit in other tests. The rate of 

manganese loss was also more rapid than in the earlier experiment 

at a lower pH. The rate also was steady during most of the experiment. In figure 3, the relation time to manganese concentration is 

shown graphically. The rate of manganese loss is also apparently 

a function of the amount of sand present.
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B23

NOmil/M 83d SiaVd Nl 'NOI1VH1N30NOO 3S3NVONVW
10.0

FIGURE 3. Rate of loss of manganese from solutions in contact 

120 ppm, initial pH 8.03.
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B25

The redox potential at equilibrium in the solutions containirg bicarbonate was not calculated. With one exception, the manage nese 

activity in each solution and the observed pH intersect in an area where 

manganese activity is controlled by the solubility of MnCO3 . This 

isolubility is not a function of Eh. However, later experiments 

phowed that manganese carbonate was not precipitated from these 

solutions to any significant extent.

The difference between the results shown in figure 3 and those from 

'he previous experiment given in figure 2 suggests that oxidation of 

Mie manganese was a more important factor in solutions 13-18 than 

: n 7-12. The difference in pH between the two experiments is erough 

"o indicate considerably more oxidation in solutions 13-18 than in 7-12. 

'lowever, it seems unlikely that the exchange capacity of the ssnd is 

r^ery strongly influenced by pH in this range.

The results are most readily explainable as a combination of effects 

of cation exchange and oxidation. In both sets of solutions the decrease in dissolved manganese early in the experiment was rapid and 

i s attributable to adsorption, perhaps with a small amount of concurrent oxidation. The rate at which manganese content of solutions 

represented in figure 2 decreased approached zero as the exchange reacion neared equilibrium. The solution containing 100 g of sand, 

however, showed a slow and steady loss of dissolved manganese, as

 ^presented by the straight part of the lowest curve in figure 2. This 

1 oss is attributable to oxidation, and appears to be a function of the 

."and surface available. Very little manganese was lost from the soluions containing less than 100 g of sand after the ion-exchange rea ction

ad reached equilibrium.

The rate of manganese oxidation is known to be slowed hi the presence of bicarbonate (Hem, 1963), and even though the pH of solu4 ion 13, which contained no sand, rose to 8.6, very little oxidation had 

occurred after 216 hours. 

The rates of manganese loss shown in figure 3 are more rapid than

hose in figure 2, but the amount of manganese which remained in 

rolution when the curves merged into straight lines was about the 

^me in all the solutions that contained the same amount of sand. 

This general agreement supports the hypothesis that the initial loss of

 nanganese is the result of ion exchange and the steady, slower rate 

attained later in the experiment was the result of oxidation, catalyzed 

' y the sand surface and its relatively immobile layer of adsorbed manganese ions.
10.0

5.0

1.0

0.5

0.1

100 200 300 400 

TIME. IN HOURS

FIGUBE 4. Rate of loss of manganese from solutions in contact with sand 

concentration 2,000 ppm; Initial p
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B27

Nine additional sets of solutions were prepared, by similar procedures, to evaluate the effect on- manganese loss of such factors ar pH 

md different amounts of sulfate and bicarbonate. The results obtained in a solution containing 2,000 ppm of sulfate, added as sodium 

r^lfate, are given in figure 4. Large amounts of sulfate were found 

previously (Hem, 1963) to retard the oxidation rate, and the sodium 

rould tend to compete with manganese for ion-exchange position on 

: \e sand. The results in figure 4 are, therefore, to be expected manganese loss was very slow in that particular experiment.

From one of the other sets of solutions it was learned that a biM.rbonate concentration of 1,000 ppm brought about manganese renoval from all solutions at the same rate, whether sand was present 

:" not. These results indicate precipitation of manganese carbor ate; 

:he solubility product was considerably exceeded in these solutions. 

Cther sets of solutions that gave faster manganese loss rates than the 

results in figure 3 had sulfate concentrations of 100 ppm and b; carl: onate concentrations of 20 ppm.

The cation-exchange capacity of sand 2, measured by means of 1.0 

riolar manganese chloride, was 0.66 meq per 100 g. This value is not 

directly applicable to the much more dilute solutions used in these 

experiments. The amount of manganese in solution in each of the 

bottles was 5.0 mg at the start of the experimental runs. Manipulations during the study required the removal of most of the solution 

f "om the bottle to make Eh, pH, and manganese determinations, and 

c nly an approximate figure can be calculated for the amount of dissolved manganese removed from the system in this way. Hence, the 

total amount of manganese retained in the sand could not be accurp.tely 

calculated. The results are of value as indicators of reaction rrechanisms and oxidation rates as affected by the sand surfaces, but are 

r ot susceptible of strict quantitative interpretation.

To obtain a strict control of pH in the further study of oxidation 

nte, the problem was approached in a different manner. A solution 

c ontaining 5.0 mg Mn+2 and 50 mg SO4~2 in a volume of 500 ml was 

prepared from stock solutions of manganous chloride and sodium 

sulfate and boiled distilled water. It was adjusted to pH 8.50 by 

rapid addition of standard potassium hydroxide, and was maintained 

at this pH by addition of the base as required. The solution was 

stirred with a mechanical stirring device, and from time to time porti ons were removed with a pipet, filtered through a plastic membrane 

flter, and analyzed for manganese. The amounts of base added and 

elapsed time were recorded. The experiment was repeated at pH 

j .82 and at pH 9.28. The sulfate was present to decrease the oxids.tion 

n,te to a level at which it would be easy to study.
LITER PER MOLE A OF THOUSANDTHS IN CONCENTRATION, MANGANESE 

A.HXSIK3H3 JlO HH.LV.M.
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B29

The results of the experiments are shown graphically in figure 5. 

Then the pH was held at 8.50, the loss of manganese was slov but 

appeared to take place at a constant rate throughout the test. The 

lialf-time for the oxidation at this pH, calculated on the basis of an 

extrapolation of the results given in figure 5, is 306 minutes, equivalent 

1 o a first-order rate constant of 2.3 X 10~3 min -1. At pH 8.8P and 

9.28, the oxidation rates were not constant. In each of these two runs, 

after a short period of slow loss of manganese, the rate increased and 

si much higher rate of loss was maintained until almost all the manTanese had disappeared from solution. The last few observations 

'nade in each of these tests suggest a possible further change in rate, 

but these results were probably influenced to some extent by difficulty 

: n measuring small manganese concentrations accurately. Consequently, their significance is questionable.

The increase in rate is attributed to the autocatalytic effect of preipitated manganese oxide and has been observed before in this work 

''Hem, 1963). If the rates observed during the first 10 minutes at pH 

\28 and during the first 20 minutes at pH 8.88 are extrapolated, they 

;^ive reaction half-times of about 15 minutes and 120 minutes, re^ecI'ively. The first-order reaction-rate constants are 4.6X10~2 and

-W .8X10~3 , and may represent the uncatalyzed oxidation rates at these 

">H levels.

Although later rates observed at these two pH's were not completely 

"onsistent, both sets of data show a period of fairly constant reaction

-ate. The rate constants for these sets are 2.5 X10-1 at pH 9.2P and

-.I X10-2 at pH 8.88. The amount of increase in the rate brought about 

%y the autocatalytic effect of oxide surfaces is substantial.

Two additional experiments were made to explore the effect of sand 

£ on the rates of loss of manganese. These experiments were made 

Tinder conditions identical to those of the preceding ones, except that 

"5 g of sand 2, which had been previously saturated with Mn+2 by 

exposure to 1.0 MnCl2 and washed free of chloride with distilled 

^ater, was added to each solution before adjusting the pH to the level 

r t which it was held constant. The sand grains were kept in motion 

1 :y the stirrer. The rates of manganese loss in these solutions at pH 

P.80 and 9.00 are shown in figure 6.

The autocatalytic effect which was so pronounced in the solutions 

1' hat contained no sand is missing from the results shown in figure 6. 

'''he reaction half-time at pH 8.80 is 133 minutes and at pH 9.0 is 63 

i linutes, and a steady rate was maintained throughout the experiment 

in both solutions. When the difference in pH is considered, the rates 

ixTere not significantly different from the initial ones observed in solutions that contained no sand. The reaction rate constants are 

:* .1X10-2 at pH 9.0 and 5.2 X1Q-3 at pH 8.80.
MINUTES 

2of sand g 

/ I 

x/ 7
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B31

Figure 7 is a graph showing first-order rate constants for the manganese oxidation reaction in aerated water, in relation to pH for solutions which contained no sulfate or bicarbonate (from laboratory 

studies made for Hem, 1963), and for solutions containing 200 ppm 

sulfate. The solid points represent the data for samples which contained sand. The effect of pH and sulfate have been noted previously.

The apparent inhibiting effect of the sand on the autocatalysis is 

of some interest. Several possible explanations may be suggested. 

The deposition of oxide on the sand surfaces evidently results in a 

material whose properties differ from the freely suspended particles 

of oxide that occur in solutions that contained no sand. Ultiirately, 

perhaps a deposit would be built upon the sand grains that wo^ld be 

as effective in catalyzing the oxidation as the free suspended oxide, 

but the amount formed in this experiment was evidently not suf^ient 

to do so. In addition to effects stemming from the physical nature 

of the oxide deposit, the oxidation at the sand surface may favor 

formation of an oxide chemically different from that formed ir solution. As noted previously (Hem, 1963), the oxide formed be-tween 

pH 8.5 and 9.5 is a mixed product, having a decreased proportion of 

tetravalent manganese at higher pH. It seems possible that the sand 

surfaces may alter what is a complex reaction mechanism and may 

promote the formation of the tetravalent oxide rather than tH less 

highly oxidized species.

Possibly a larger amount of sand in the reacting solution would have 

changed the results. Because of mechanical difficulties in maintaining 

continuous motion of the sand, amounts much greater than those used 

in the experiment could not be utilized. The adsorbed manganese on 

the sand added another factor. Some of this manganese may have 

been replaced by sodium from the solution, but the ratio of manganese 

to sodium in the solution should have been high enough to prevert, any 

important amount of replacement. Obviously, sand having a capacity 

of adsorbing more manganese would have obscured the results.

The results of this set of rate studies suggest that the sand arected 

the course of the oxidation reaction by inhibiting the autocatjrlyzed 

process which occurred in the absence of the sand. A further indication of this effect can be obtained from comparisons of relatior ships 

of OH"1 demand to manganese oxidation.

The amount of base consumed and corresponding amounts of manganese precipitated were calculated from the titration data. To obtain the amount of manganese lost by precipitation, the total anount 

present at each time of sampling for manganese content was calculated. 

This calculation required correcting for manganese removed in the 

sampling process. A satisfactory stoichiometric relationship was not 

obtained, however, perhaps partly because carbon dioxide was adsorbed
FIRST-ORDER RATE CONSTANT (k
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B33

by the alkaline solutions from the air. Nevertheless, the results did 

suggest that considerably less of the standard base was required per 

unit of manganese oxidized at pH 9.28 than at pH 8.5. The amount 

of base consumed in the oxidation experiments where sand was present 

was greater than would have been expected at the same pH in the 

absence of the sand-grain surfaces. This fact is a further suggestion 

that the sand surfaces may promote the formation of MnO2.

Although the magnitude of the catalytic effect of 15 g of sand 2 in 

500 ml of solution appears to be small, the effect observed in the earlier 

experiments where 100 g of sand was present is evidently appreciable. 

The effect in the presence of the larger amount of sand and of bicarbonate ions can only be evaluated approximately. The rates obtained 

in earlier work (Hem, 1963) at pH 9.0, 9.1, and 9.3 are about the same, 

or slightly slower, than those in figure 6 representing solutionr containing 200 ppm of sulf ate. If these rates are extrapolated to pH 8.0, 

a value for k of about 1X10"4 min~1 is obtained. In the presence of 

100 g of sand at this pH, the rate constant estimated from figure 2 

is 1.9 X10"4. However, the actual oxidation rate at pH 8.0 in the presence of bicarbonate is not readily determinable, and the extrapolated 

'T alue is probably not reliable.

ADSORPTION COLUMN EXPERIMENTS

The sands used in these experiments have a mean particle diameter 

of about 0.5 mm. The surface area per gram is therefore abovt 47 

sq cm. The area of 100 g of sand is 4,700 sq cm. In a bottle- containing 500 ml of water, this gives an area of sand of about 9.4 sq cm 

^)er milliliter. When the sand was just saturated with water the weight 

composition was 23.5 percent water and 76.5 percent sand. Hen°-e, in 

i.n aquifer that was composed of this type of sand the surface area 

oer milliliter of water would be about 140 sq cm. This area is about 

15 times that available per milliliter of water in the bottles containing 

100 g of sand.

In many aquifers a considerable amount of fine-grained material is 

^resent. Because the surface area of a sphere is proportional to the 

square of the diameter, an increase in proportion of small particles 

: ncreases the area per unit weight very markedly. At the same time,

 ^he small particles decrease the space available for water in a unit volume of aquifer material. Consequently, the surface area of solids per 

nilliliter of water in many aquifers is far greater than 140 sq cm, a

 ^alue which is probably near the minimum for granular aquifers.

In order to duplicate more closely the solid-liquid relationships that 

occur in partly saturated or intermittently saturated sediments, some 

curther experiments were performed with the adsorption column used
B34 CHEMISTRY OF MANGANESE IN NATURAL WATER

earlier to determine the cation-exchange behavior of sand 2 toward 

manganese and calcium. A considerable volume of manganese solution was passed through the column and the cation-exchange positions 

on the grains were filled with manganese ions. The column was then 

drained and was allowed to stand undisturbed in the laboratory for 

several months. During this period, the upper part of the sand had 

become completely dry. Work with the column was resumed without 

disturbing the sand. In these resumed experiments, solutions of manganese chloride were used which contained about 100 ppm HCOa"1 

(added as sodium bicarbonate) to stabilize the pH. No calcium was 

present in the solutions. In general, the solutions were more alkaline 

than the ones used previously and had a pH ranging from 8.0 to 9.0.

It was immediately evident that manganese was teing absorbed by 

the column to a greater extent than before, and eTT en though large 

volumes of solution were passed rapidly through the column, the manganese content of the effluent did not reach the level of the original 

solution. Rather, the manganese concentration of tl °s effluent seemed 

to reach a constant value considerably below the initial concentration. 

Furthermore, the total amount of manganese picked up by the column, 

about 75 mg, greatly exceeded the exchange capacity of the sand it 

contained. Evidently, manganese was now being rapidly oxidized 

in the column. As the experiments continued, the upper 10 cm or 

so of sand in the column became dark brown. Tl is color was the 

result of deposition of a coating of oxidized manganese on the sand 

grains, aad MnO2 was shown to be present by X-ray diffraction analyses of the material.

The approximate rates of oxidation of manganese as solutions 

passed through the sand ranged from a half-time of about 10 minutes 

at pH 8.6 to about 120 minutes at pH 8.0. These values are inexact 

because the actual contact time of solution with the column is uncertain. Probably most of the oxidation occurred in the upper part 

of the column. However, these rates are appreciable faster than the 

ones indicated for pure water in figure 7. They are much faster than 

the ones observed for solutions containing bicarbonate in the absence 

of sand (Hem, 1963). The catalytic effect of the oxide coating on 

these sand grains is, thus, definitely indicated.

When solutions having a pH greater than 9.0 were passed through 

the column, some of the manganese oxide was dislodged from the sand 

grains and the effluent contained a brown colloidal suspension of manganese oxide.

When solutions containing 1.0 ppm Mn, having a pH below 7.0, were 

passed through the column, the oxidation reaction appeared to be reversed and somewhat higher concentrations of manganese were pres-
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B35

rit in the effluent than had been in the original solution. Measure-

<mts of pH, Eh, and manganese content for two initial solutions and

: c the effluent discharged after constant manganese concentration

as been reached are given below. The calculated values for Eh ^ere

i'^termmed from graphs of pH, Eh, and manganese activity publirhed

vrlier (Hem, 1963). The close check obtained between observed

" nd calculated Eh suggests the solution in the column had reached a

3f ate of oxidation equilibrium after about 600 ml had passed through.

___ Eh (volts')

Solution 

Input _________________

Effluent_--_____.___. 

Input.________________

Effluent..--._-____

pH

7.21

7. 78

6.35

7. 18

Measured i

0.47

.45

.50

.49

Calculated

0.49

.45

.49

Mn (ppri)

1. C

. 1

1. C

1. 2

The amount of oxygen present in aerated water at a temperature of 

2''"'° C is more than sufficient to oxidize 10 ppm of manganese. HowT'er, if larger concentrations of manganese are present, the availability 

TC oxygen may become the principal factor affecting oxidation rate.

The control of pH in the column experiments was not very exact 

and the time the solution spent in the part of the column where oxida- 

^ on was most active is uncertain. However, the average time required

*">r half the manganese to be oxidized at pH 8.5 in the column appeared 

V be about 12 minutes. In an earlier experiment at pH 8.5 in aerated 

instilled water, the half-time for manganese oxidation was about 30 

n inutes. In the presence of 100 ppm bicarbonate activity the rate was 

Ti uch slower than this. The results of the column experiments are not 

sufficiently refined to show the exact magnitude of the catalysis effect 

T speeding the oxidation of manganese, but they do indicate that the 

VTect is much stronger than where the area of sand per unit volume 

)" solution is small.

An indication as to the reversibility of the oxidation process can be 

obtained from the column results. Establishment of equilibrium by 

r troducing solution at pH 7.21, which involved oxidation and precipirtion, required about the same time as the reverse equilibrium obYined by introducing a solution whose pH was 6.35. In the attain- 

"ent of this equilibrium, manganese was reduced and dissolved.

APPLICATION OF LABORATORY AND THEORETICAL 

DATA TO NATURAL CONDITIONS

The laboratory work described in this paper gives some general 

r dications as to the mechanisms that may be involved in the deposition

->* manganese oxides in natural environments and in the reduction 

fid solution of such deposits in surface and ground water. A trief

 r capitulation of the possibilities is in order.
B36 CHEMISTRY OF MANGANESE IN NATURAL WATER

Because the work was concerned with inorganic chemical processes, 

the effects of aquatic plant or animal life were not considered extensively, though it is not to imply that the behavior of manganese as 

observed in nature is not extensively influenced by these life processes. 

The fundamental energy relationships that may b^ used to evaluate 

the chemistry of manganese are applicable to all environments.

The mineral fragments that constitute stream sediments, soils, and 

the sedimentary rocks have some cation-exchange capacity. In normal 

fresh-water environments the exchange positions are generally occupied by the common divalent cations of natural water. Manganese is 

likely to be present in this adsorbed material in somewhat greater 

proportion than in the solution at equilibrium. Adsorption-desorption equilibria may have considerable influence on the amount of 

manganese in solution in water in contact with sediments.

Deposition of manganese oxide coatings on mineral grains can be 

initiated when a bed of sand or a subsoil stratum through which 

manganese-bearing water has been moving becomes dry and remains 

so for a time. Renewed movement of manganese-bearing aerated 

water through this material can bring about localized deposits of 

manganese oxide in the areas where Eh is high enough. The oxidation 

process will deplete the dissolved oxygen of the wr.ter, and probably 

deposition of oxide will not occur very deeply in the zone of saturation, 

and in some instances would be possible only in the soil or subsoil 

zones where water moves generally by unsaturated flow.

The behavior of divalent manganese that may be brought into 

contact with manganese oxide will be controlled by three fundamental 

variables, pH, Eh, and dissolved manganese activity. Increases in 

any one of these variables above the level at which the initial deposit 

of oxide was formed will bring about further deposition of oxide, and 

decreases in any one of the three variables will cause the deposited 

oxide to go back into solution. These general principles should help 

explain the occurrence of manganese oxide deposits in lakes, streams, 

and elsewhere, and the occurrence of water in reservoirs, both surface 

and underground, which contains excessive quantities of manganese. 

Although literature relating to these occurrences is abundant, it is 

generally difficult to determine from published c'escriptions which 

factors are principally responsible for high mangar^se concentrations 

and the formation of oxide deposits, because field observations and 

measurements are almost always inadequate. Measurement of Eh in 

the field is generally difficult and in some environments is not practicable. The variability of pH and manganese content, as well as of 

other dissolved constituents in stream water, generally requires more 

sampling and study than are given.
DEPOSITION AND SOLUTION OF MANGANESE OXIDES R37

Manganese oxide can be deposited rapidly from solution in the 

T3sence of suitable surfaces until a condition of equilibrium as pre- 

; ",ted from Eh-pH diagrams is reached (Hem, 1963). Thus, at pH

*.? manganese would be deposited from aerated water which had more 

Ji an about 0.10 ppm of manganese in solution. However, at pH 7.2 

M ".nganese oxide deposition would not be expected unless the solution 

ontained more than 1.0 ppm. These observations correlate reasonr.bly 

veil with the observed behavior of manganese in ground water. In 

^'ound-water aquifers or in water near the bottoms of reservoirs and 

akes, which is prevented from circulation by thermal stratification, 

v iucing conditions under which considerable amounts of Mn+2 will be 

;f ?,ble in solution can be expected. There is little difficulty in finding 

rechanisms to explain waters having a high manganese content 

x" ,en, more than one reasonably plausible explanation can be given. 

"' ie deposition of manganese oxide on the beds of flowing streams is 

', more difficult process to explain, although some of the instances 

sported may be due to fairly simple processes.

Ingols and Wilroy (1962) have reported some observations on the 

ranganese content of the Chattahoochee River near Atlanta, Ga. 

During the summer months, manganese reached high concentrat; ons 

r the deeper parts of a reservoir in which water was stored for generaion of power. As a result, water that was discharged into the r; ver 

1 iring the late summer and early fall had as much as 15 ppm man- 

;r,nese. The manganese was precipitated as oxide in the riverbed 

V r a few miles below the dam the first year the reservoir was u^ed, 

)iit was not carried far downstream in solution. The precipitation, of 

r ost of the manganese can be readily explained as a result of an in- 

.rease in Eh and pH of the discharged water as it was aerated1 by 

lowing down the river channel. The average pH of the river witer 

.1; the Atlanta raw-water intake (Lohr and Love, 1952) was reported 

x be 6.9 during 1950; the maximum observed value was 7.2 and the 

rinimum was 6.5. Evidently the Eh of the river water is sufficiently 

rgh for removal of manganese.

It is not unreasonable to postulate that organisms in river witer 

r ay bring about micro-environments adjacent to or within their struc-

-ure which attain and maintain higher pH and Eh levels than tl xr>se 

x : the main water body, and that these growths thereby oxidize the 

manganese and convert it to MnO2 . This hypothesis has been fre- 

^iiently suggested and, as noted earlier in this report, it was proposed 

FT Ljunggren (1953) as the explanation for manganese oxide deposits 

xi~ streambeds in southern Sweden.
B38 CHEMISTRY OF MANGANESE IN NATURAL WATER

COPRECIPITATION OF IRON AND MANGANESE

Before concluding that biochemical reactions arc the best explanation for all occurrences of manganese precipitates ir streams, it should 

be noted that the manganese oxide precipitates in streambeds often 

form without obvious indications of biological activity. Certain 

aspects of manganese chemistry in the presence of iron may be of 

interest in this connection. Manganese oxide strea mbed deposits for 

which analyses are available contain considerable amounts of iron, 

and iron can be assumed to be present in the water.

In previous work (Hem, 1963), it was shown that when ferric 

hydroxide was precipitated from a mixed solution of manganous 

and ferrous ions, the ferrous iron content quickly dropped below detection, and the manganese in solution also decreased although at a slower 

rate, so long as pH was above 6.5. According to theoretical solubility 

calculations for pure manganese solutions, no precipitation of manganese should have taken place; hence, the precipitation of iron must 

have had a considerable influence.

Mixed deposits of manganese and iron oxide in streambeds could 

perhaps be brought about by coprecipitation, originating on a coating 

of ferric hydroxide adhering to the surfaces of rocks in the stream. 

Such an adherent coating; would form when the iron was precipitated 

at a pH low enough to impart a positive electrostatic charge to the 

particles of ferric hydroxide. A pH below 6.0 would be required 

for this precipitation of positively charged material, but a considerably higher pH would be required to cause coprecipitation of manganese. In the dilute, poorly buffered water of a stream draining an 

igneous or metamorphic forested terrane, the necessary pH fluctuations probably could occur. The manganese content of the deposits 

could be increased at the expense of iron by a replacement process 

which could take place when the stream characteristics were favorable.

The chemical system in a segment of river where manganese oxide 

is being deposited is more strongly influenced by oxygen than a groundwater system or the system at the bottom of a lake. In the river segment, the Eh of the water is a function of dissolved oxygen which 

normally is present in excess and which will not H appreciably affected by oxidation or precipitation reactions involving manganese 

in solution. The pH, although not constant, is fixed by carbon dioxide 

equilibria. The amount of ferrous iron in solution will fluctuate in 

response to pH and Eh. The activity of Fe+2, however, will be a very 

small fraction of the total amount of iron apparently present. Most 

of the latter will consist of colloidally dispersed or molecular Fe 

(OH) 3 or complexed forms of ferrous iron. The activity of ferrous 

ions can be estimated from the pH and Eh (Hem, 1960c) for equilib-
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B39

rium conditions. Biochemical utilization of iron may tend, at times, 

to depress the ferrous ion activity below the equilibrium value, and 

any ferrous iron produced by local reactions in the streambed vill be 

carried off by the moving water. One may expect the Fe+2 activity 

in the stream to be at, or possibly at times below, the equilibrium level 

predictable from the Eh and pH of the water, in the presence of solid 

Fe(OH) 3 .

Ferric hydroxide particles have a positive surface charge in water 

whose pH is low enough, and under these conditions would be attracted to negatively charged surfaces such as those of mineral^1 in a 

streambed. Once formed, a thin layer of ferric hydroxide might 

participate in a reaction with manganous ions in the stream water, 

which can be represented :

2Fe(OH) 3 c+Mn+2 +2 H+1 =MnO2 c+4 H2O + 2 Fe+2

From the standard free-energy values for the participating: substances, the equilibrium constant for this reaction is calculated as 

10 5 - 67. The reaction is normally displaced to the left however, 

when activity of Fe+2 is very low and activity of Mn+2 is relatively high, 

some replacement of ferric hydroxide by manganese dioxide will 

occur.

The negatively charged surface of the MnO2 would tend to attract 

more ferric hydroxide, and the process could be repeated to build up 

a deposit of the two substances. The proportion of manganese to iron 

would depend on the length of time that conditions were favorable 

for manganese deposition and on rates of reactions involved.

The mass-law equation for the above reaction

m [Mn1C -1-91 +2 TTT4-19 A v ] [H+]2

can be used to calculate equilibrium activities for manganese r,t va- 

-ious activities of Fe+2 and H+1 . For example, if [Fe+2 ] is 10'12 

molal, at pH 6.5 the solubility of manganese will be 10~5 - 33 molar or 

about 0.26 ppm, and at pH 6.0 the solubility of manganese will be 

only 10-6 - 33 or 0.026 ppm. Thus, if the activity of ferrous iron, >^hich 

appears as the second power in the equation, can be maintained at a 

low enough value, the precipitation of manganese can be expected even 

though the activity of manganese in solution is well below the point 

where oxide precipitation would be expected when iron is absent 

(Hem, 1963).

Calculated activity of Fe+2 at various pH levels (Hem, 1960c, p. 

50) indicates that at pH 6.0 the Eh would have to be +0.69 v, and 

at pH 6.5, +0.60 v to maintain [Fe+2 ] at 10'12 molal. These Eh values
B40 .CHEMISTRY OF MANGANESE IN NATURAL Y'ATER

are probably somewhat above levels commonly reached in stream waters but may not be impossible in strongly aerated water in this pH 

range.

Although Ljunggren (1953) did not obtain Eh measurements, the 

pH's of .streams he studied in Sweden were between 6.2 and 6.8 during 

one period of summer low flow and generally near or somewhat below 

6.0 at higher flow a year later. The amount of manganese in solution 

was generally less than 0.10 ppm at the low-flow period, but manganese 

concentrations were not determined in the later observations. The 

above calculations suggest that coprecipitation may play a part in the 

deposition of oxides in those streams.

According to Krause (1928), the isoelectric point for ferric hydroxide is at pH 5.2. The observations of the writer (Hem and Skougstad, 

1960) suggest that the isoelectric point may be between pH 5.5 and 

6.0 in solutions resembling natural water of low dissolved solids 

concentration. When the pH of the stream water is below the isoelectric point, adherent precipitates of ferric hydroxide may be 

formed in the bed. Replacement of part of the ferric hydroxide by 

manganese dioxide will be favored when the water is at a higher pH. 

This theoretical model obviously needs to be tested by detailed study 

of a stream in which manganese and iron deposition occurs before it 

can be verified. Among the kinds of observations needed are accurate measurements of pH and Eh in the precipitation zone, periodic 

sampling upstream and downstream from the precipitating section 

to determine seasonal changes in dissolved manganese activity, and 

detailed, examinations of the deposits and the surfaces on which they 

form to identify any organisms associated with them.

SUMMARY AND CONCLUSIONS

Because manganese oxide deposits can occur in widely differing natural environments, it seems logical to postulate more than one mechanism for their formation. Circulation of manganese from soil to trees 

and back to the soil by leaf-shedding and similar biochemical mechanisms involving other plant types will tend to lep.k some soluble 

manganese, both to surface runoff and to ground water. Biochemical 

processes probably are frequently associated with both the solution 

and the deposition of manganese, where energy relationships are 

favorable for the organisms that participate.

The oxidation of manganous ions in aerated water is very slow at 

near-neutral pH, but the oxidation rate can be increased by inorganic 

catalysis. Many of the mineral surfaces to which natural water is 

exposed have some cation-exchange sites, which may strongly attract 

manganese ions. The surfaces of feldspathic sand grains, for example,
DEPOSITION AND SOLUTION OF MANGANESE OXIDES B41

adsorb manganese, and when they are saturated with adsorbed manganese ions, the surfaces catalyze the oxidation process.

Coprecipitation of manganese oxide and ferric hydroxide can take 

place in aerated water whose pH is below Y.O, if ferrous iron activity 

is kept low. Such a process may be important in the deposition of 

mixed oxides in streams and lakes.

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B42 CHEMISTRY OF MANGANESE IN NATURAL WATER

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