Mineralogical speciation of sulfur in

acid sulfate soils from Luleå, Sweden

Niklas Gunnarsson

Natural Resources Engineering, master's

2018

Luleå University of Technology

Department of Civil, Environmental and Natural Resources Engineering

i

Preface This master thesis ends five years of intense studying and a wonderful time at Luleå University

of Technology. The thesis covers both geochemistry and mineralogy and was financially

supported by Interreg Nord in cooperation with the County Administrative Board of

Norrbotten. It was part of a larger project called “Coastal watercourses in Bottenviken: Method

development and ecological restoration- A cross-border Swedish-Finnish cooperation project”

that is in progress between 2015-2018.

I would like to start to express my gratitude to my external supervisor Fredrik Nordblad from

the County Administrative Board of Norrbotten for sharing his knowledge and experience

from working with these soils.

Next, I want to thank Bernhard Dold and Thomas Aiglsperger my advisors at the university

for their support and knowledge with sulfide-rich materials. A special thanks to Thomas for

sharing his mineralogical knowledge, for helping me with sample preparation and with

microscope and SEM-interpretation. Finally, I would like to thank my classmate Astrid

Theander for the company and helping hands with sampling during cold winter days.

Niklas Gunnarsson

Luleå, October 2018

ii

Abstract

Marine sulfide – bearing sediments that oxidize when in contact with oxygen and leach out

elements in high concentrations to small watercourses have been a problem for many years all

over the world especially around the Bothnian Bay. The purpose of this study was to further

investigate the sulfur mineralogy present in acid sulfate soils in the area of Luleå, Sweden. A

secondary aim was to see if elements leach out and accumulate in an acid sulfate soil closer to

the recipient. Samples were taken in two profiles (one oxidized and one waterlogged) from

four sites (sites A-D) and were analyzed for whole rock geochemistry. Two sites were further

investigated for mineralogy in polished samples with an optical microscope, Raman

spectroscopy and SEM-EDS. Each profile consisted of three layers: oxidation zone, transition

zone and reduced zone. The oxidation zone above the groundwater table was light grey with

brown iron hydroxide staining. Parts that lied under the water table were dark grey-black with

in general strong odor (“rotten eggs”) due to its sulfur content. It was usually straightforward

to distinguish and separate the layers from each other directly in the field, however in some

cases pH was needed for confirmation.

A general feature of investigated polished samples is the presence of abundant framboidal

pyrites that are common in reduced marine sediments. The transition zone was formed in sub-

oxic conditions and this feature is reflected by the mineralogy. Many morphologies of the

framboidal pyrite were observed in this layer and signs of both dissolution and formation

occur. In the sample from site C one could observe elemental sulfur in form of large (up to 50

µm) euhedral crystals. In the samples with pH<4, no sulfides occur as they have been replaced

by jarosite (site B). Site C lacks these sulfur-bearing hydroxides which is thought to be due to

a sulfur concentration of <0.2 %. Sulfur shows extensive leaching at most sites but at site B and

D1, it accumulates in the transition zone. Elements like cobalt (Co), nickel (Ni) and zinc (Zn)

are leached out or are accumulated further down in the profile. Elements that could have been

transported and have accumulated in the waterlogged profiles are Co, Ni, Zn and chromium

(Cr) and in some profiles also copper (Cu) and vanadium (V).

Keywords: acid sulfate soil, metal leakage, low pH, SEM, elemental sulfur, framboidal pyrite

iii

Preface ........................................................................................................................................ i

Abstract .................................................................................................................................... ii

Abbreviations and definitions ............................................................................................ iv

1 Introduction ..................................................................................................................... 1

1.1 Background ........................................................................................................................... 1

1.2 Aim ......................................................................................................................................... 1

2 Literature research........................................................................................................... 2

2.1 Forming acid sulfate soils .................................................................................................. 2

2.2 Oxidation of sulfides .......................................................................................................... 3

2.3 Improving ecological status ............................................................................................... 4

2.4 Trace elements in sulfidic sediments .............................................................................. 4 2.4.1 Association to organic matter ....................................................................................................... 5 2.4.2 Association to hydroxides ............................................................................................................. 5

2.5 Character of sediment ......................................................................................................... 7

2.6 Classification ........................................................................................................................ 7

2.7 Acid sulfate soils around the Baltic Sea .......................................................................... 7

2.8 Prediction .............................................................................................................................. 9 2.8.1 Comparison with ARD .................................................................................................................. 9

3 Materials and methods ................................................................................................. 10

3.1 Study site ............................................................................................................................. 10

3.2 Sample collection .............................................................................................................. 12

3.3 Sample preparation ........................................................................................................... 14

3.4 Polished samples ............................................................................................................... 15

3.5 Whole rock analyses ......................................................................................................... 15

3.6 Optical microscope, Raman spectroscopy and SEM-EDS ......................................... 16

4 Results and discussion ................................................................................................. 17

4.1 Gammelstad (Site B) ......................................................................................................... 19 4.1.1 Optical microscope and SEM ...................................................................................................... 19 4.1.2 Geochemistry of the profiles ....................................................................................................... 27

4.2 Björsbyn (Site C) ................................................................................................................ 30 4.2.1 Optical microscope and SEM ...................................................................................................... 30 4.2.2 Geochemistry of the profiles ....................................................................................................... 34

4.3 Sunderbyn (Site A) ............................................................................................................ 37 4.3.1 Geochemistry of profiles ............................................................................................................. 37

4.4 Björsbyn (Site D) ............................................................................................................... 39 4.4.1 Geochemistry of profiles ............................................................................................................. 39

4.5 Geochemical comparison ................................................................................................. 41

4.6 Classification ...................................................................................................................... 42

5 Conclusion ...................................................................................................................... 44

5.1 Future work ........................................................................................................................ 44

6 References ....................................................................................................................... 45

7 Appendix ........................................................................................................................ 51

7.1 Appendix 1: Profile descriptions .................................................................................... 51

7.2 Appendix 2: Mass balance ............................................................................................... 53

7.3 Appendix 3: Immobile element vertical distribution ................................................. 57

7.4 Appendix 4: Whole rock geochemical data ................................................................... 57

iv

Abbreviations and definitions

Abbreviation Explanation

AASS Actual acid sulfate soil (oxidation layer)

ABA Acid base accounting

AMD Acid mine drainage

AP Acid potential

ARD Acid rock drainage

ASS Acid sulfate soil

AVS Acid volatile sulfides

bgs Below ground surface

BP Before present

Eh Redox potential

GTK Geological Survey of Finland

ICP-SFMS Inductively coupled plasma mass spectrometry

LOI Loss on ignition

NAG Net acid generation

NNP Net neutralizing potential

NP Neutralization potential

OZ Oxidation zone

PASS Possible acid sulfate soil (reduced layer + transition zone)

RZ Reduced zone

SEM-EDS Scanning electron microscope-energy dispersive x-ray spectroscopy

SGU Swedish Geological Survey

TS Total solids

TZ Transition zone

XRD X-ray diffraction

100 ha 1 km2

A1,A2…D2 Profile names according to site

Chemical Elements Arsenic As Silicon Si

Aluminum Al Sodium Na

Cadmium Cd Sulfur S

Calcium Ca Titanium Ti

Cobalt Co Uranium U

Chromium Cr Vanadium V

Copper Cu Yttrium Y

Iron Fe Zinc Zn

v

Lead Pb Zirconium Zr

Magnesium Mg Heavy rare-earth elements (Tb (65)-Lu (71)) HREE

Manganese Mn Light rare-earth elements (La (57)-Gd (64)) LREE

Molybdenum Mo Rare-earth elements (La (57)-Lu (71)) REE

Nickel Ni

Minerals and chemical compounds

Anglesite PbSO4 Bicarbonate HCO3−

Barite BaSO4 Bisulfide HS−

Calcite CaCO3 Ferric iron Fe3+

Elemental sulfur S0 Ferrous iron Fe2+

Gibbsite Al(OH)3 Hydrochloric acid HCl

Greigite Fe3S4 Hydrofluoric acid HF

Goethite FeOOH Hydrogen gas H2

Hematite Fe2O3 Hydrogen ion (proton) H+

Ilmenite FeTiO3 Hydrogen peroxide H2O2

Iron(III) hydroxide Fe(OH)3 Hydrogen sulfide H2S

Jarosite KFe3(OH)6(SO4)2 Iron (II) hydrogen sulfide Fe(HS)2

Mackinawite Fe(1+x)S Lithium metaborate LiBO2

Magnetite Fe3O4 Nitric acid HNO3

Manganite MnOOH Organic matter CH2O

Marcasite FeS2 Oxygen gas O2

Pyrite FeS2 Polysulfide, anion Sn2−, S(n−1)

2−

Pyrolusite MnO2 Sulfate SO42−

Pyrrhotite Fe(1−x)S Sulfuric acid H2SO4

Schwertmannite Fe16O16(OH)12(SO4)2 Water H2O

Siderite FeCO3

Xenotime YPO4

Zircon ZrSiO4

1

1 Introduction

1.1 Background

Acid sulfate soils are “nasty” soils (Dent and Pons, 1995). H₂S is formed as a product from

sulfate reduction due to decomposing organic matter when oxygen is absent, that will later

favor the formation of iron sulfides (Berner, 1984). The term “acid sulfate soil” means that

sulfides come in contact with oxygen and start to oxidize, changing the character of the

sediment (Dent and Pons, 1995). This can be enhanced by anthropogenic activities that require

ditching such as farming or forestry that lower the groundwater table (Sohlenius et al., 2004)

or by shafting when building large structures (Pousette et al., 2007). The oxidation lowers the

pH to <4 leading to increased weathering and mobilization of aluminum (Al), to some extent

iron (Fe) and trace elements such as cobalt (Co), nickel (Ni) and zinc (Zn) (Åström, 1998;

Sohlenius and Öborn, 2004). If the acid leachate reaches smaller watercourses with minor

dilution effect, it has a significant effect on the water quality (Åström and Björklund, 1996;

Filppa, 2012). Especially Al is very toxic to the environment as aluminum hydroxide

precipitates with fish kills as a result (Erixon, 2009). Acid sulfate soils are spread over several

continents and cover more than 17 million hectares worldwide (Andriesse and van Mensvoort,

2006) where its environmental impact is a global problem. Lots of studies have been done in

other regions of the world, but it was not until recent years that the topic started to be

investigated further in Sweden.

1.2 Aim

Recent studies have concluded which and to what extent elements leach out from acid sulfate

soils but there is still a lack of knowledge about the soil mineralogy connected to these metals.

This master thesis was focused on investigating the sulfur mineralogy (both primary and

secondary) present in the acid sulfate soil stratigraphy from the Luleå area in Norrbotten. One

other aim was to compare sample geochemistry and connect trace element trends to sulfide

mineralogy. This was to increase the knowledge on both actual acid sulfate soil (AASS) and

potential acid sulfate soil (PASS) so that lowering of groundwater table and acid sulfate soil

management will have a minimum influence on nearby watercourses.

Questions to answer in this thesis were:

• What sulfide minerals can be observed in acid sulfate soil from the Luleå area?

• What secondary sulfur minerals are present in the acid sulfate soil?

• Are there any differences in sulfur mineralogy between the sites and what are they?

• Are elements leached out from a more oxidized profile, transported and then

accumulated in a waterlogged profile downstream?

To answer these questions, samples from four sites and in two profiles at each site were

collected and investigated on their whole rock geochemistry. Furthermore, polished sections

from two of the sites were prepared to study their mineralogy.

2

2 Literature research

2.1 Forming acid sulfate soils

Acid sulfate soils (ASS) cover an area worldwide as large as 1.5*105 km2 (Pons et al., 1982) and

can be found in southeast Asia, West Africa, South America, eastern Australia and Europe

(Andriesse and van Mensvoort, 2006). In Finland, ASS is distributed over an area of 3500 km2

(Palko, 1994) which is the largest area in the world where up to 130 000 ha acid sulfate soil is

cultivated land (Yli-Halla et al., 1999). Around the Baltic Sea, these sediments were formed

during the Littorina Sea which was a former stage of the Baltic Sea with higher salinity than

present levels (Sohlenius, 1996). Today they can be found 0-100 m above current sea level

(Erviö, 1975; Palko, 1994) due to an isostatic uplift of 9-10 mm/year around the Bothnian Bay

(Johansson et al., 2004; Kaniuth and Vetter, 2004). This is a result from being forced down by

a 2-2.7 km thick ice sheet during the last ice age (Siegert et al., 2001).

Sulfide formation is the first step in the formation of acid sulfate soil where the reduction of

sulfate (SO42−) by bacterial activity can be illustrated by the reaction (Berner, 1984):

2CH2O + SO42− → H2S + 2HCO3

− (1)

where sulfate is reduced to hydrogen sulfide (H₂S) by using organic matter as an energy

source. The metastable iron sulfides that are most common in sediments are mackinawite

(Fe(1+x)S, x=0-0.11 simplified to FeS) and greigite (Fe3S4) (Morse and Rickard, 2004; Rickard

and Morse, 2005). It is considered that mackinawite and greigite are responsible for the black

color in acid sulfate sediments but there is only limited evidence that actually supports the

theory (Rickard and Morse, 2005). Formation of mackinawite has been described with two

reactions between ferrous iron and hydrogen sulfide in reduced conditions (Rickard, 1995):

Fe2+ + H2S → FeS + 2H+ (2)

Fe2+ + 2HS− → Fe(HS)2 → FeS + H2S. (3)

Greigite is considered to form from mackinawite by a solid- state transformation when oxygen

is present but also elemental sulfur could be used (Rickard and Luther, 2007; Rickard and

Morse, 2005). Production of iron sulfides may be limited by organic carbon for sulfate

reduction, sulfate or ferrous iron to react with (Berner, 1984). Mackinawite is quite common in

marine sediments such as ASS, but it is thought to be replaced fast by pyrite (FeS2) that is

formed by reaction between FeS and dissolved H2S or with polysulfides (Morse et al., 1987;

Rickard and Luther, 1997; Rickard, 1975). The transformation of mackinawite to pyrite can be

illustrated in the reactions by (Butler et al., 2004):

FeS + H2S → FeS2 + H2 (4)

FeS + Sn2− → FeS2 + Sn−1

2− (5)

3

where reaction (4) is dominating in anoxic conditions (Rickard and Luther, 1997) and reaction

(5) in partly oxic/anoxic conditions (Rickard, 1975). Iron sulfides are often only a minor part of

marine sediments although some systems may be enriched in sulfides. Generally, iron sulfides

in sulfide-rich sediments can be classed into two categories namely acid volatile sulfide (AVS)

and pyrite (Howarth and Jørgensen, 1984; Morse et al., 1987). AVS includes minerals such as

amorphous- FeS, mackinawite, greigite and pyrrhotite (Fe(1−x)S, x=0-0.2) (Cornwell and

Morse, 1987). They are presumed to give the black sulfide coating in reducing sediments. In

marine sediments pyrite is the most abundant iron sulfide (Berner et al., 1979) and it can form

either aggregates that occur in spherical shapes known as “framboids” or form separate

crystals (Sweeney and Kaplan, 1973). Enrichment of acid volatile sulfides has been connected

to low pyrite formation by several factors such as depletion of H2S due to a more rapid

formation of FeS in systems that are rich in iron (Gagnon et al., 1995). Other reasons that may

affect both rate and extent of pyrite formation are that H2S is converted to HS− as a dominating

sulfide specie in alkaline systems (Rickard and Morse, 2005).

2.2 Oxidation of sulfides

Oxidation of sulfide-rich sediments occurs when they get in contact with oxygen that will

produce sulfuric acid and lowers the pH<4 which creates the acid sulfate soil (Cook et al.,

2000). This process is accelerated by artificial drainage and excavation (Toivonen and

Österholm, 2011). Studies have concluded that oxidation of metastable iron sulfide gives

elemental sulfur as the first product of oxidation (Bloomfield, 1972a) that will oxidize further

to sulfate (Van Breemen, 1973). Half-times for oxidation of acid volatile sulfides have been

measured between 29-60 min (Burton et al., 2009, 2006) and the sulfides are considered to be

completely oxidized within a month (Purokoski, 1958). Two oxidants have been considered

important for iron sulfide oxidation: oxygen and ferric iron (Van Breemen, 1973). Sometimes

oxidation could reach down close to 2 m (Österholm and Åström, 2002) with secondary iron

minerals as a common precipitate in this layer (Collins et al., 2010). Oxidation of metastable

iron sulfide that occurs rather fast is illustrated by the reaction (Berner, 1962; Boman et al.,

2008):

10FeS1.1 + 24O2 + 26H2O → 10Fe(OH)3 + 11H2SO4. (6)

Pyrite that has a slower oxidation rate may instead oxidize according to (Hart, 1962):

4FeS2 + 15O2 + 14H2O → 4Fe(OH)3 + 8H2SO4. (7)

When pH drops below 4.5, ferric iron is more important as oxidant than oxygen and increases

the oxidation rate with a lower pH as a result (Nordstrom, 1982). The bacteria formerly called

Thiobacilli ferrooxidans (Van Breemen, 1973) nowadays Acidithiobacillus ferrooxidans has a

significant role since it catalyzes the oxidation of ferrous iron to ferric iron in natural systems

(Johnson and Hallberg, 2003). This oxidation of sulfides lowers the pH down to values 2.5-4

(Dent and Pons, 1995; Van Breemen, 1973).

4

Oxidized iron hydrolyzes and precipitates as iron hydroxides when pH rises above 3.5 (Van

Breemen, 1973). During pH<4 jarosite (KFe3(OH)6(SO4)2) (Brown, 1971) and schwertmannite

(Fe16O16(OH)12(SO4)2) (Sullivan and Bush, 2004) can be found. Both jarosite and

schwertmannite will later hydrolyze to goethite (FeO(OH)) and may do so over several months

or years (Burton et al., 2007). Oxidation of iron sulfides may cause deoxygenation and acidic

drainage water with a high metal release as a result (Burton et al., 2009; Morgan et al., 2012).

Metals can be mobilized in acid sulfate soils by either sulfide oxidation or by the weathering

of phyllosilicates that are accelerated by the more acidic pH (Åström and Björklund, 1996;

Stumm and Furrer, 1987). When pH in the recipient is getting below 5.2 several organisms will

not survive (Eilers et al., 1984). Studies show that elements released include aluminum (Al),

calcium (Ca), cadmium (Cd), cobalt (Co), copper (Cu), magnesium (Mg), manganese (Mn),

sodium (Na), nickel (Ni), silicon (Si), and uranium (U) are occurring as free ions as long as pH

is low and are therefore bioavailable (Nystrand and Österholm, 2013).

2.3 Improving ecological status

Some studies have tried to improve the quality of drainage water and to minimize the leaching

of trace metals by raising the pH in the soil with e.g. liming (Åström et al., 2007). A problem

in that study was that iron hydroxide particles coated the lime particles and therefore

hampered the improvement (Åström et al., 2007). Powell and Martens (2005) tested controlled

tidal flooding in combination with lime to remediate ASS with improved pH and trace metal

redistribution toward reducing depths as a promising result. Tidal flooding is likely by either

anthropogenic reflooding or sea-level rise due to climate change and will transport reactive

trace elements down to the redox boundary of the soil where they will be enriched (Keene et

al., 2014). Waterlogging of ASS has shown to decrease the Al concentration in porewater in

boreal climate and also that formation of iron sulfides was low despite increased reducing

condition (Virtanen et al., 2014). Waterlogging of acidified acid sulfate soils has been tried in

northern Australia with an improved water quality as a result (Johnston et al., 2009). This is

possible due to dilution of acidity, buffering ocean water and reducing conditions consuming

H+. Remediation with waterlogging where schwertmannite is present increased the pH up to

6.5 and reduced the iron to ferrous iron with a release of sulfate and alkalinity that later gave

precipitation of siderite (FeCO3) (Burton et al., 2007). In Finland, they have included a

regulation plan for groundwater adjustment to minimize sulfide oxidation from ASS (Uusi-

Kämppä et al., 2013).

2.4 Trace elements in sulfidic sediments

Metals are generally much more mobile when pH is low and therefore pH is a crucial factor

that controls metal mobility (Gundersen and Steinnes, 2003). Formation and sedimentation of

iron sulfides into marine sediments is one of the main removal paths for iron and sulfur from

marine waters (Berner, 1984). Both AVS and pyrite are sinks for trace elements with adsorption

and coprecipitation processes to regulate their mobility and bioavailability (Morse et al., 1987).

5

Huerta-Diaz and Morse (1992) and Scholz and Neumann (2007) have concluded that some

trace elements of geochemical importance can be included in sedimentary pyrite in varying

proportions. These trace elements could include large amounts of arsenic (As), cobalt (Co) and

manganese (Mn), moderate amounts of copper (Cu), nickel (Ni) and chromium (Cr) and small

amounts of lead (Pb), zinc (Zn) and cadmium (Cd) that may form their own sulfide minerals

instead. Huerta-Diaz and Morse (1992) also conclude that trace metal content in pyrite is

limited by pyrite production or trace metal depletion. Typical sulfide metals such as Zn, Ni

and Co have been shown to commonly be incorporated in iron sulfides (Volkov and Fomina,

1974). Cadmium has a trend of forming insoluble sulfides in anoxic, sulfide-bearing sediments

when H2S is present (Jacobs et al., 1987).

2.4.1 Association to organic matter

In the oxidation layer the elements zinc, copper and cadmium can be associated with organic

matter and then be released to the water column when the organic matter is being consumed

with diffusion upward or downward in the profile as a result (Scholz and Neumann, 2007).

Aluminum (Nordmyr et al., 2008) and chromium are usually also associated with organic

matter (Shaw et al., 1990) where chromium together with vanadium and nickel forms

dissolved complexes (Olson et al., 2017).

2.4.2 Association to hydroxides

Leaching of elements from both iron and manganese hydroxides are common when pH is

lowered resulting in their dissolution. The elements vanadium (Morford et al., 2005), cobalt

(Shaw et al., 1990) and chromium (Algeo and Maynard, 2004) are associated with manganese

hydroxides and will be liberated when the hydroxides are unstable. When iron hydroxides are

reduced, elements such as arsenic (As) are released to the environment (Widerlund and Ingri,

1995).

The group of lanthanoids (also known as the rare earth elements) that are the elements La (57)-

Lu (71) in the periodic table can be found in moderate to strongly elevated concentrations (ca.

36-4700 µg/l) in acidic water e.g.(Åström, 2001). Of importance for the mobility of lanthanoids

is the phyllosilicate matrix (clay minerals) holding the elements in varying accessory minerals

(Harlavan et al., 2009) but also the grain size where (sub-) micrometric particles offer a large

surface area for reactions (Åström et al., 2010). Organic matter has also been found to be an

important carrier of these elements (Åström et al., 2010). Figure 1 shows how elements are

liberated from different phases and get more accessible for uptake and leaching from less

stable phases when pH drops from near neutral to more acidic.

6

Figure 1: Element cycling among different phases depending on pH, where the major mobilization paths

are numbered (Claff et al., 2011). CDB=crystalline iron oxide phase.

In a study from Finland they compared trace element concentrations between the three layers

of an acid sulfate soil (Åström, 1998). The conclusion was that manganese (Mn), zinc (Zn),

nickel (Ni) and cobalt (Co) were frequently released from the soil when it was oxidized but

also that small proportions of Zn, Ni and Co had reprecipitated further down in the transition

zone. They also concluded that only a small proportion of copper (Cu) was released and no

extensive leaching of iron (Fe), titanium (Ti), chromium (Cr) and vanadium (V) was observed.

This is thought to be due to that Cr, V and Ti are connected to weathering resistant fractions

or to organic matter and the small losses of Fe are thought to be due to precipitation of iron

hydroxides in the oxidation zone (Åström, 1998). In Sweden 15-45 % of Mn is lost from sulfide-

bearing sediments upon oxidation (Öborn, 1991). Factors affecting geochemistry of ASS are

grain size, mineral distribution and composition, concentration of organic matter and where

the groundwater table is located (Sohlenius et al., 2015).

7

2.5 Character of sediment

Acid sulfate soils are recognized by their odd color due to iron precipitates, bad odor due to

sulfur, sparse vegetation due to pH<4 and the speed of transformation from black to grey color

as a result of oxidation (Dent and Pons, 1995). The sulfur concentration in these soils could

vary between 0.6-2.1 % in the transition zone and the sulfidic layer while it is usually between

0.04-0.46 % in the oxidation zone (Åström, 1998). It is common that pyrite with 1-5 % is the

most abundant sulfide mineral in acid sulfate soils (Bloomfield, 1972a). Other minerals such

as marcasite (FeS2), mackinawite (FeS), greigite (Fe3S4), and amorphous iron sulfides could

also be present (Burton et al., 2008; Bush et al., 2004; Bush and Sullivan, 1997) with content

below 0.01 % (Bloomfield, 1972b) and up to 0.6 % (Berner, 1971). Åström (1998) found in their

study that the uppermost two layers had a pH between 3-5 while in the reduced layer a pH of

6-7. The low pH is primarily due to oxidation of sulfides that forms sulfuric acid but also by

hydrolysis of iron (Van Breemen, 1973).

Toivonen and Österholm (2011) concluded in their study that fine-grained phyllosilicates are

the main source for metals and that metal concentration variation is controlled by grain size.

2.6 Classification

Acid sulfate soil can be separated into “acid sulfate soil” and reduced sulfide-rich sediment

that may “potentially form acid sulfate soil” (PASS) (Brinkman and Pons, 1973). The

nomenclature used for acid sulfate soil layers is: actual acid sulfate soil (AASS) where pH is

low and iron hydroxide mottles occur and potential acid sulfate soil (PASS) where pH is near

neutral and has a sulfur content of at least 0.2 % (Edén et al., 2012). The Swedish Geological

Survey (SGU) classifies soil with in-field pH <4 as acid sulfate soil , samples with in-field pH

between 4.4< pH <6 and pH <4 after oxidation as semi-acid (TZ) and in-field pH > 6 and with

pH<4 after oxidation as reduced soil (Sohlenius et al., 2015). They also classify them into

different classes after field pH and sulfur content, where A+B are for pH<4, C for 4-4.5 and

pH> 4.5 is concluded not acid sulfate soil. For sulfur content, these classes were > 1 % class

one, 0.6-1 % class two and 0.2-0.6 % for class three, below that limit was not classified as sulfide

bearing (Edén et al., 2012; Sohlenius et al., 2015). The iron/sulfur ratio has been proposed as a

tool for classifying the acid potential where ratios < 3 is very acidifying and >60 as non-sulfidic

(Pousette et al., 2007).

2.7 Acid sulfate soils around the Baltic Sea

A rough areal estimate of the acid sulfate soils in Sweden was done to 140 000 ha (Öborn, 1994).

Mapping of acid sulfate soils in Sweden has had a low priority through the history and

therefore the area is not well defined (Öborn, 1989). Most acid sulfate soils occur along the

coast of the Baltic Sea and down around Lake Mälaren and Hjälmaren (Öborn, 1989). These

were deposited in the Litorina Sea (7000–4000 BP) (Sohlenius, 1996) but are now due to

isostatic land rebound lifted over the sea level (Figure 2).

8

Figure 2: Map showing the maximum sea level (light grey) along the Swedish coast where the majority of the

sulfide-rich sediments occur below this level, modified from (Sohlenius, 2011).

Acid sulfate soil form when the top meter is oxidized e.g. due to excavation or drainage for

agricultural use (Öborn, 1991). This was common in the early 19th century when farmlands

were prepared by draining wetlands accelerating the oxidation of underlying sulfides

(Österholm and Åström, 2004). They are characterized by low pH (3.9-4.5) and almost no

sulfur content due to the oxidation of sulfides (Sohlenius and Öborn, 2004). Potential acid

sulfate soil (PASS) both in Finland and Sweden is characterized by a black color with high AVS

content that indicates metastable iron sulfides (Backlund et al., 2005; Georgala, 1980). Around

Lake Mälaren the sulfide-bearing sediments are classified as muddy clay while along the coast

in northern Sweden as muddy silt (Sohlenius et al., 2004). The muddy clay is characterized by

a blue-grey color and is dominated by pyrite while the muddy silt along the northern coast

has black color due to dominating FeS minerals (Sohlenius et al., 2004). The soil is usually

characterized by a silicate matrix of granitic composition where soils in Sweden and Finland

show similarities (Åström, 1998; Öborn, 1989). A connected acid sulfate soil stratigraphy can

be seen along the Norrbotten coastline from Piteå up to Torneå where typical soil profiles

along this line have been described by (Fromm, 1965).

9

The most important mobilization factor is increased weathering when pH is lowered

(Sohlenius and Öborn, 2004). Mobile elements are mostly accumulated in the transition zone

where varying reducing and oxidizing conditions occur (Sohlenius and Öborn, 2003). When

the soil reaches above current sea level, iron monosulfides are oxidized in the top layer and

pyrite is preserved and therefore trapping elements such as Co, Ni and Zn (Boman et al., 2010).

Aluminum and copper were found to be strongly correlated with organic matter while Cd,

Co, Mn, Ni and Zn were likely associated with hydroxides (Nordmyr et al., 2008). In Finland

is iron not abundantly leached out from acid sulfate soils (Åström, 1998; Österholm and

Åström, 2002). This is due to iron hydroxide precipitation in the oxidation zone that may trap

elements such as arsenic that likes to adsorb to iron hydroxides when pH<7 (Wällstedt et al.,

2010). Areas with sulfide bearing sediments usually have elevated concentrations of sulfur,

nickel and copper in biota and have a lower resistivity than other fine-grained sediments

(Puranen et al., 1997). This makes it easier to find PASS by biogeochemistry and geophysics

(Sohlenius et al., 2004). Most sulfide bearing sediments found in Sweden lies close to the sea

and have been uplifted within the last 2000 years (Sohlenius, 2011). Porsöfjärden north of Luleå

has in periods been extremely affected by low pH and high metal concentrations (Erixon, 2009;

Filppa, 2012). Cadmium (Cd), nickel (Ni), manganese (Mn), copper (Cu) and cobalt (Co) have

been found to commonly leach out from ASS in general (Sohlenius and Öborn, 2004) while the

elements Co, Ni and Zn are typically found in drainage from acid sulfate soils in boreal climate

(Åström and Corin, 2000). In western Finland the concentrations of Zn, Ni and Co are 30-40

times higher in stream water draining sulfidic sediments compared to other sediments (Åström

and Björklund, 1996).

2.8 Prediction

In a study done at the University of Uppsala, they recommend that net neutralization potential

(NNP) should be investigated in sulfide-bearing sediments and if <5 also metal content should

be analyzed to prevent sudden acidification (Lax and Sohlenius, 2006). The company MRM

Konsult AB have developed a leaching test specifically for ASS where one reduced and then

several oxidized leaching tests are done with cycles of wetting and drying (MRM, 2007). In a

master thesis from the Luleå University of Technology, the MRM leaching test and Acid-base

accounting (ABA) test was compared for predicting ASS leaching where the MRM test was

concluded to give a more complete picture (Paulsson, 2012). Pousette (2010) describes a

leaching test developed by Luleå University of Technology for ASS that is similar to the MRM

leaching test where only the sample weight (10 g vs 25 g), water ratio and the number of

leaching steps were different for the test procedures.

2.8.1 Comparison with ARD

Acid-base accounting (ABA) test is the most common prediction method used for predicting

acid rock drainage (ARD) since it is fast and cheap. It is generally a result from calculating the

NNP as the difference between neutralization potential (NP) and acid potential (AP) where all

sulfur is assumed to be from pyrite and all neutralizing minerals as calcite (CaCO3) (Sobek et

al., 1978).

10

In most cases this is not accurate since the sulfur mineralogy is more complex (Dold, 2010).

This has led to a more high-resolution ABA-test where sulfate and iron hydroxide mineralogy

can be extracted from the sulfide mineralogy giving a more accurate AP (Dold, 2003). Skousen

et al. (1997) shows how to take siderite (FeCO3) into consideration to get a more accurate NP.

Another test that is frequently used is the NAG-test where hydrogen peroxide (H2O2) is added

to accelerate sulfide oxidation. The pH is measured in the leachate and then titrated stepwise

to neutral pH to get an acid forming capacity (Ian Wark Research Institute, 2002). Pyrite is the

most common sulfide mineral in mine tailings and is oxidized when oxygen is present (Dold,

2014). Mine tailings are very similar to ASS in many ways when looking at sulfide oxidation

and hydrolysis of minerals. One major difference is that ASS oxidizes much faster due to

dominating iron monosulfides compared to pyrite as dominating specie in mine tailings (see

section 2.2, (Dold, 2017)). See Dold, (2017) for more about prediction.

3 Materials and methods

3.1 Study site

This study has been carried out on actual acid sulfate soil (AASS) and potential acid sulfate

soil (PASS) from four locations around the city of Luleå in northern Sweden. The studied areas

are Site A close to Sunderbyn, Site B near Gammelstad, Site C and D both in Björsbyn (Figure

3).

Figure 3: Map showing sample points and studied areas.

11

Site A is located on a field ca. 40 m from a rainwater outlet, site B in a thicket ca. 30 m next to a

tributary to the Luleå river, site C profile 1 in a thicket and profile 2 close to the reeds behind

a garden business premises and site D profile 1 on a small meadow and profile 2 close to the

reeds ca 1 km from site C (Figure 4; Table 4).

Figure 4: Sample pictures from site B and C.

The geological units at site A and B are acidic intrusive rocks from the svecokarelian orogeny

ca. 2.85-1.87 Ga. At site C and D the units are a metasedimentary unit like greywacke or

quartzite, a basic intrusion like gabbro and an acidic intrusion all 2.85-1.87 Ga. Close-by there

is a younger acidic porphyritic intrusion, ca. 1.88-1.74 Ga (Figure 5). According to the soil map

from the Swedish Geological Survey, silt and clay dominate the sample areas.

12

Figure 5: Geological map of sample areas. The geological maps were generated via the SGU map generator

(http://apps.sgu.se/kartgenerator/maporder_sv.html).

3.2 Sample collection

Sample locations were chosen using a soil map from SGU that uses soil type and the shore-

level model as parameters. According to Sohlenius et al., (2015) this highlights areas with

silt/clay that have been uplifted the last 2000 years as the highest probability of finding AASS

and PASS.

13

Samples were taken in two profiles at each location, one on elevated levels showing extensive

oxidation and one closer to the shore that is still waterlogged with limited oxidation (Figure

6).

Figure 6: A simplified overview of the spatial relationship between profile 1, profile 2 and a nearby stream.

Brown: humus, light grey-orange: OZ, dark grey-black: TZ-RZ.

An Edelman auger was used for collecting samples at different depths. Soil samples were first

placed in a complete profile on plastic bags to visually distinguish and measure at which depth

each layer was located. In general the color was enough to separate layers (Figure 7) but to

ensure these visual assumptions the pH was measured with a WTW pH 330i pH-meter and a

Sentix 41 pH-electrode.

Figure 7: A: An example of a brownish acid sulfate soil (AASS). B: An example of a dark grey

potential acid sulfate soil (PASS).

14

To do the measurement a piece of material from a layer was mixed with a small part of milli-

Q water and stirred to a paste in which the pH was measured. This was done at least once per

layer but at some of the locations with more uncertainty, it was necessary with more frequent

measurements. After measuring the pH, one representative sample per layer (350-1000 g) was

collected in plastic bags. After collection, the samples were stored in a cold storage at the

university laboratory for 1-3 weeks until sample preparation.

3.3 Sample preparation

The most important minerals in ASS are the iron sulfides, so minimizing mineralogical

changes to samples during sample preparation is therefore important (Claff et al., 2010).

Drying of anoxic soils is thought to likely affect redox-sensitive elements that are present

(Bordas and Bourg, 1998; Rapln et al., 1986). Despite the risk, it is still common to use drying

as a pre-treatment method (Åström, 1998). When all samples had been collected they were

placed in aluminum forms and weighted before put into an oven where they were dried for 2-

3 days at 50 °C. When all samples were dry they were weighted again to reconstruct the initial

water content of each sample. During the drying, the color changed from a dark grey-black to

more light grey-brown on reduced samples (Figure 8).

Figure 8: A: Sample (site C profile 1, RZ) before drying. B: Sample (site C profile 1, RZ) after drying.

One piece from each layer, both profiles from site B and C (total 12) were chosen for preparing

polished samples. Homogenization was done with a glass bottle against a flat surface that

firmly crushed the sample into evenly sized grains. When the sample was homogenous

enough it was split with a riffle splitter at least three times until the remaining part was ground

with an agate mortar to reduce the grain size further. Ground samples of representative

sample aliquots were put into small plastic bags marked with sample-name that were

subsequently sent to ALS Scandinavia AB for whole-rock analysis. The remainder of the

samples were put back into plastic bags and stored in the cold storage.

15

3.4 Polished samples

The 12 pieces from site B and C were put into plastic molds for molding the pieces into epoxy.

One-part resin and one-part hardener were stirred together in a cup until entirely mixed. The

molds were filled up close to the top and then placed in a vacuum pump to take away all

bubbles from the epoxy. When all bubbles had been removed, they were hardened for 24 hours

before polishing. Polishing was done in a Struers LaboPol-6 polishing machine at 1200 rpm

with a lubricant as a cooling agent. The first three samples (Site B1 (RZ), Site B2 (RZ) and site

C2 (RZ)) were going through the entire polishing program with 4 minutes each on P500, P800

and P1200 abrasive paper as a start and then using 15, 9 and 3 microns diamond paste on

separate disks for polishing (Figure 9).

Figure 9: Polishing machine and polished samples after rough polishing.

A problem was that the samples were too soft, and a relief occurred. Therefore, on the

remaining samples, grinding with abrasion paper for 30 seconds at a time was done without

the additional use of diamond polishing. The samples C2 (OZ), B1 (OZ) and B2 (OZ) got deep

cavities from the polishing and were covered in a new layer of epoxy and then re-polished

with no improved result. Sample C1 (RZ) was also remolded and re-polished since it cracked

when trying to free it from the mold.

3.5 Whole rock analyses

ALS Scandinavia prepared the samples for analysis with two methods. The first one was total

digestion with a mixture of hydrochloric acid (HCl), nitric acid (HNO3) and hydrofluoric acid

(HF) that is mixed with a small part of a sample in a test tube that is heated until digestion is

reached. The other method is mixing lithium metaborate (LiBO2) with the sample and heating

it to 1000 °C melting it to a pearl that is digested in nitric acid so-called fusion analysis. The

samples from total digestion were then analyzed with 2nd gen ICP-SFMS, Element 2/XR and

from fusion analysis with first generation ICP-SFMS. Analysis with ICP-SFMS were been done

according to SS-EN ISO 17294-1, 2 (mod) and EPA-method 200.8 (mod). Samples were dried

in 50 °C but were total solid (TS) corrected to 105 °C according to SS 02 81 13-1.

16

After that loss of ignition (LOI) was determined for nearly all samples where samples were

heated in an oven to 1000 °C and their weight before and after were compared.

To highlight element enrichment or depletion it is possible to calculate a coefficient τ with the

equation (Anderson et al., 2002):

τi,j =Cj,w∗Ci,p

Ci,w∗Cj,p− 1. (8)

Where (j) is the mobile element and (i) is the immobile element in weathered (w) and parent

material (p). Positive values denote enrichment of the mobile element and a negative value a

depletion (Chadwick et al., 1990). When the coefficient reaches -1 it implies total depletion and

around zero neither enrichment or depletion (Ling et al., 2014).

3.6 Optical microscope, Raman spectroscopy and SEM-EDS

All samples were thoroughly investigated with reflecting light using a Nikon Eclipse LV

100POL optical microscope equipped with a Nikon DS-Fi1 camera. Interesting reflecting

grains were photographed and marked for further investigation with SEM and Raman at the

university. Raman spectroscopy was done on samples from site B profile 2 (OZ+TZ) and site

B profile 1 (OZ+RZ) using a Bruker Senterra equipped with an Olympus BXFM optical

microscope. Raman spectra were obtained using a 532 nm laser beam (d=1 micron) at 0.2 mW

with 45 s integration time and 3 repetitions. The monochromatic laser light (wavelength=530

nm) is sent on the sample, reacting with a mineral grain (inelastic scattering) that changes the

energy level of the laser photons that can be measured as Raman shifts (Colthup et al., 1975).

This shift is characteristic for each Raman-active mineral and gives information about the

mineral structure, hence minerals that have similar chemistry in scanning electron microscopy

e.g. iron oxides (hematite Fe2O3; magnetite Fe3O4) or iron sulfides (pyrite FeS2; marcasite FeS2)

can be distinguished. Scanning electron microscopy with energy-dispersive spectroscopy

(SEM-EDS) was performed on samples from site B profile 1 (OZ-RZ), from site B profile 2 (OZ)

and from site C profile 1 (TZ) to determine the chemical composition and get high-resolution

images of interesting grains. For this a ZEISS Merlin FEG-SEM was used at 20 keV accelerating

voltage and 1.1 nA beam current. The scanning electron microscope shoots a high-energy

beam of electrons at a selected part of a sample. The electrons of the beam interact with the

sample and produce secondary electrons, backscattered electrons and X-rays as a result of the

reaction (Reimer, 1998). These electrons are collected by detectors that compute an image of

the sample. The emitted X-rays are used in the energy dispersive spectroscope that gives a

distinct energy signature for every element that is collected in a spectrum making it possible

to interpret the mineralogy (Reimer, 1998).

17

4 Results and discussion Most profiles in this study and in earlier research consist of three layers with an oxidation zone

near the ground surface, a transition zone further down and a reduced zone at the bottom.

Figure 10 visualizes these layers in photographs and in idealized profile sketches with typical

colors and gradual transitions between layers. Typical colors in this soil are light grey with

orange mottles in the oxidation layer that turn darker grey further down when it gradually

turns into a more reduced environment. The bottom layer is often dark grey-black caused by

iron sulfides and/or by organic matter. The stratigraphy of profile 1 at each site starts with a

15-30 cm humus layer followed by a 43-110 cm thick oxidation layer.

Underlying is a transition zone (40-82 cm) that is depending on the groundwater table that lies

65-110 cm below ground surface (bgs). For profile 2 the humus layer is 0-30 cm in thickness

with a 15-45 cm underlying oxidation layer. The transition zone varies here between 0-74 cm

with a groundwater table between 0-40 cm below ground surface. One general trend that can

be seen at each site is that at site A the oxidation layer and transition zone are equal in

thickness, at site B the oxidation zone is thicker than the transition zone, at site C the oxidation

zone is thinner than the transition zone and at site D the oxidation zone is thicker than the

transition zone. Since the transition zone is formed through groundwater fluctuation one can

interpret that where the oxidation zone is thicker than the transition zone the groundwater

table is permanently lowered on a stable level whilst where the opposite result is seen the

groundwater level is varying more naturally with weather and flood.

For better visibility of the dominating geochemical processes, site B and C will be first

discussed with evidence of the mineralogy that will be linked to the chemical trends followed

by interpretation of processes at site A and D with a final comparison between all sites.

18

Fig

ure 10: F

ield p

rofiles w

ith id

ealized sk

etches fo

r clarification

. Bro

wn

: hu

mu

s,

oran

ge-lig

ht g

rey: O

Z, d

ark

grey

: TZ

and

black

: RZ

.

19

4.1 Gammelstad (Site B)

4.1.1 Optical microscope and SEM

In general, the samples from profile 1 have light grey color with a 1-5 mm brown oxidation

rim but the sample from the OZ is orange/red-grey in color. Samples from Profile 2 have a

light brown/beige-dark grey color with a 1-4 mm brown-grey rim. Samples from this site have

a trend where large silicate grains occur high up in the profile and decrease in numbers with

depth.

Oxidation zone (profile 1)

The entire sample was oxidized which gave no oxidation rim. In this position no sulfides were

found in the microscope. The sample contained lots of iron hydroxides and one of them was

jarosite (Figure 11).

Figure 11: A+B: Secondary electron image. C: Raman spectrum of grain with reference spectrum for

comparison. D: EDS spectrum with a picture from the optical microscope.

Jarosite is usually formed at pH < 3 but is stable at pH< 4 (Brown, 1971). One major problem

of forming jarosite is usually the source for potassium but it is quite common in heavily

weathered sulfide-rich mine tailings (Dold, 2014). As can be seen in Figure 11A and 11B the

large white spot is composed of many small jarosite grains. This area was first detected in

optical microscope due to its characteristic yellow color. This was later confirmed by the

Raman spectrum that corresponds well to the reference spectrum for jarosite, except for a small

shift to the right (Figure 11C). Later also the chemical composition was confirmed in SEM

where the EDS spectrum showed in Figure 11D matches the composition of jarosite

(KFe3(OH)6(SO4)2) except for a minor sodium content.

20

Transition zone (profile 1)

Further down in the profile one can find microcrystals of pyrite that together forms the

spherical shape “framboidal pyrite”. Many observations of framboidal pyrite in a variation of

sizes were done with an optical microscope. Some of them had a growth rim around the main

framboidal core (Figure 12A).

Figure 12: SEM image of framboidal pyrites with a sulfide rim. D: Chemical composition of the rim.

The circular grains are the framboidal pyrite that is composed of several small cubic-pyramidal

pyrite grains. The EDS spectrum shown in Figure 12D confirms that also the rim around is

composed of iron sulfides indicating the grain has been growing. It is uncertain which iron

sulfide this rim is formed from but since the iron peak is small this supports that it could be

pyrite rather than an iron monosulfide. Some observations of framboidal pyrite with other

morphologies have also been done (Figure 13).

21

Figure 13: A+B: Framboidal pyrite with internal structure. C: Backscatter image of unevolved framboidal

pyrite. D: Secondary electron image of unevolved framboidal pyrite.

Figure 13A and 13B show a framboidal pyrite without the growth rim, instead it has a well-

organized pattern like it has been growing one rim at the time until it reaches its spherical

shape. This can be compared to the unorganized core in Figure 12. The grain in Figure 13C

and 13D is chemically the same as pyrite and has a similar shape to a framboid but could be

either at the formation or in the dissolution phase of formation since the microcrystals are

welded together which makes it near impossible to distinguish individual grains. No

secondary iron minerals were documented with images but some barite (BaSO4) was observed

(Figure 14).

Figure 14: Barite grain with confirmation from the EDS spectrum.

22

Barite is thought to precipitate with bacterial help (Hanor, 2000), this due to its incorporation

with the redox-sensitive sulfur form sulfate. The observation of barite indicates that some

oxygen has reached down to this level which is in accordance with the sub-oxic transition

zone.

Reduced zone (profile 1)

In the reduced zone, black layers filled with framboidal pyrites can be seen with microscope

(Figure 15) but also with the naked eye. Especially in the central part of the sample a lot of

both small (4 µm) and large (16 µm) framboidal pyrites were observed.

Figure 15: Framboidal pyrite layer in an optical microscope with reflected light (10x ocular).

All the rounded black dots are framboidal pyrite grains that are arranged along a certain line

that appears black. In some cases, the framboidal pyrite had formed aggregates of different

sizes and shapes called polyframboids (Figure 16).

Figure 16: Framboidal pyrite forming aggregates of different shapes.

23

Figure 16A shows a rod-shaped aggregate with framboidal pyrites where each framboid has

been compressed together. They are showing a varying size distribution from 5-30 µm where

some grains show signs of a hexagonal cross-section (Figure 16B). It is also interesting to note

that, they have a 120° angle between each framboid. Figure 16C and 16D show another large

aggregate with framboidal pyrite where they are staked on each other in a larger heap. One

can see in Figure 16D that the pyrite microcrystals are growing on the outside of the framboid,

making it grow. Some grains have holes where microcrystals have been or going to be later.

Sometimes framboidal aggregates keep their spherical shape and grow larger instead (Figure

17).

Figure 17: Polyframboid formed by smaller framboids.

Figure 17A shows that this sample contains a lot of framboidal pyrites (all white dots). The

larger framboidal aggregate is formed from 10 µm framboidal pyrites (Figure 17C). As can be

seen in the well-shaped framboids, the microcrystals have been welded together making it

hard to distinguish each microcrystal. Around this, an unstructured pattern of pyrite

microcrystals lies on the surface (Figure 17D). This is most likely the case before the framboidal

texture is established. Like in the transition zone, some barite was found also down in the

reduced layer (Figure 18).

24

Figure 18: Diffuse barite crystal.

The barite grain has an anhedral shape and is diffuse which could indicate that it is becoming

unstable in more reducing environments. According to the stability diagram (Tokunaga et al.,

2016), barite is stable at all pH but is depending on the redox potential where it is stable

between Eh -0.65-0.25 V.

Oxidation zone (profile 2)

Like for profile 1, there are large weather-resistant silicate grains present. In this sample, only

a few unoxidized framboidal pyrites were observed. Barite that was found in profile 1 was

also observed in this layer (Figure 19).

Figure 19: Diffuse barite grain.

Like in the reduced layer in profile 1, barite shows a diffuse morphology which could indicate

metastability and that this grain is near dissolution. Other interesting grains that were found

are some that may be either churchite and/or xenotime (Figure 20).

25

Figure 20: Churchite or xenotime grains (𝑌𝑃𝑂4).

Although both churchite and xenotime both have the same chemistry it is more likely that the

grains in Figure 20 are xenotime. This since xenotime is a more common mineral, which could

contain small amounts of HREE and is associated with more common minerals such as zircon

(ZrSiO4), ilmenite (FeTiO3) and magnetite that also were observed. The yttrium peak (k alpha)

and the phosphorus peak are nearly at the same energy level which makes them hard to

distinguish from each other.

Transition zone (profile 2)

Large silicate grains that are visible with the naked eye are also present in this sample. More

framboidal pyrites than in the oxidation layer but less than in profile 1 were observed. In this

sample oxidation did not affect the pyrite grains like it did in the oxidation zone. Examples of

sparse packed framboidal pyrites taken with 50x ocular in the optical microscope can be seen

in Figure 21.

26

Figure 21: Poorly packed framboidal pyrites (50x ocular, optical microscope).

Very few framboidal pyrites were observed in the TZ sample and when observed they were

in cross-section with microcrystals that were not densely packed as in profile 1.

Reduced layer (profile 2)

In the reduced layer, the amount of framboidal pyrites increases only slightly. Black layers

filled with framboidal pyrites like the layers in profile 1 can be seen in one corner (Figure 22).

Figure 22: Framboidal pyrites (approx. 4 µm) in layers (5x ocular, optical microscope).

In this sample many of the framboidal pyrites in these lines were laying in an oxidation rim

and showed signs of oxidation. One could also see areas with iron hydroxides over the sample

that went into the central part of the sample from the rim.

27

4.1.2 Geochemistry of the profiles

The geochemical composition of the profiles was compared to study if mobile elements from

the oxidized profile accumulate downstream in the waterlogged profile. Soil samples at site B

were taken near a pedestrian bridge next to a small river between the neighborhoods

Gammelstad and Sunderbyn. The height above mean sea level were ca. 2.5 m and 1.5 m with

ca. 30 m apart. Element concentrations from profile 1 can be seen in Figure 23 below.

In this profile, pH is close to 3.8 in the oxidation zone due to intense sulfide oxidation. This

has led to the low sulfur concentration seen in Figure 23 which explains why no observations

of sulfides were done. For the elements sulfur, manganese, molybdenum and arsenic one can

see a trend with leaching through the entire profile where sulfur has leached 1.5 % and

manganese as much as 950 ppm. This could mean that at least Mn but possibly also Mo and

As have been incorporated in sulfides in the soil and released with sulfur when oxidation

starts. One can see that iron is retained in the oxidation zone which is also supported by the

observed iron hydroxides in the polished sample. Under the optical microscope jarosite was

found among the hydroxides in the oxidation zone which is supported by a pH<4 observed in

Figure 23, that keeps jarosite stable. One can also see that vanadium follows this trend since it

is known to adsorb to e.g. ferrihydrite (Larsson, 2014) which is also known to occur at low pH

(Naeem et al., 2007). Copper shows a similar trend which could indicate adsorption or

coprecipitation to iron hydroxides (Benjamin and Leckie, 1981) and/or association to organic

matter (Scholz and Neumann, 2007). Further down in the transition zone, pH increases >6

which gives a significant increase in sulfur concentration leading to the framboidal pyrites

observed here. Varying between oxidizing and reducing conditions with ground water level

is most likely the reason for the different morphologies (growth and dissolution) observed in

the transition zone. Some pyrites could be observed in the transition zone but in the reduced

layer, the amount increased rapidly. This correlates well with the increasing sulfur

concentration which increases with depth. One can see that iron follows sulfur in the reduced

parts which could indicate its association with iron sulfides.

Another trend observed is a slight increase of aluminum in the transition zone that could be

related to either precipitation of hydroxides or formation of clay minerals. It is though

important to note that Cr, the light rare earth elements (LREE) and the heavy rare earth

elements (HREE) follow this trend.

Figure 23: Element concentrations site B profile 1.

28

Chromium is a trivalent cation and will follow the REE that are commonly associated with

clay minerals (Åström et al., 2010; Harlavan et al., 2009). Also Zn and Co follow this trend and

some research suggest adsorption to illite (clay) particles (Chester, 1965). Zinc and Cobalt

follow each other and have a trend with leaching from the OZ (16 and 4 ppm) with small

accumulation in the transition zone (3 and 1 ppm). Nickel shows the same leaching rate

through the entire profile and do not follow any other trend. Lead shows increasing

concentration upward which could be due to either small amount of anglesite (PbSO4) that has

precipitated or association to organic matter.

Barite was observed in both the transition zone and reduced zone. A pH between 6.5-7 in these

layers gives a redox potential (Eh) of -0.25 V at lowest for barite to be stable (Tokunaga et al.,

2016). Framboidal pyrite is also stable in these two layers and according to the stability

diagram for pyrite (Figure 5.6) (Ingri, 2012), it is stable only in a small window. If connecting

the redox potential (Eh) of -0.25 V for barite to the stability for pyrite it gives a stable pH of 6-

8 which overlaps the existing pH in these layers. Rickard (1970) illuminates the ongoing debate

about the formation of framboidal texture for pyrite that has been in progress for several years.

Wilkin and Barnes (1997) debate about the possibility of formation through either reaction

forming greigite or replacement of greigite that commonly has a framboidal texture. Some of

the framboidal pyrites down in the reduced layer had oxidized during storage and then started

to get a coating of hematite that showed during Raman spectroscopy. Transition zone sample

had a 4-6 mm oxidation rim while reduced sample a 1-3 mm rim. This profile was stored for

12 days before preparation making the oxidation rate 0.33-0.5 mm/day and 0.083-0.25 mm/day

for the transition zone and reduced zone. The difference in oxidation rate could indicate that

the transition zone is dominated by more easily oxidized iron sulfides and the reduced zone

dominated by framboidal pyrites but this is just speculation. Element concentrations from

profile 2 can be seen in Figure 24 below.

In the waterlogged profile, pH is near 4.5 in the oxidation layer and increases further down.

Sulfur continues to leach out (0.8 %) even though the total concentration is lower than in profile

1. Very few framboidal pyrites were observed in both the oxidation zone and transition zone

which correlates well with the low sulfur content. Elements that follow sulfur exactly are Co

(3 ppm) and Cr (26 ppm) suggesting they could be included into sulfides. The sulfur

concentration increases with nearly 1% from the transition zone to the reduced layer. The

increase in number of framboidal pyrites is far too low to correspond to that sulfur increase.

Figure 24: Element concentrations site B profile 2.

29

Therefore, the source of sulfur must be some other mineral than framboidal pyrite like iron

monosulfide or nanoparticulate pyrite. In profile 1 was barite found in the reduced parts but

in profile 2 it was found in the oxidation zone. With a pH of 4.5 in the oxidation zone this gives

a redox potential (Eh) of -0.05 V at lowest in this sample for barite to still be stable (Tokunaga

et al., 2016). For pyrite to also be stable in this condition the redox potential could vary between

-0.05-0 V (Figure 5.6) (Ingri, 2012). In the transition and reduced zone when pH goes up to 7-

7.5 the redox potential could vary between -0.35-(-0.20) V for pyrite to be stable.

Like in profile 1 one can see that iron is retained in the OZ due to precipitation of iron

hydroxides although it was not that prominent in the polished sample like it was in profile 1.

Elements that follow the iron trend are Zn, Ni, Cu, Mo and V that could be due to adsorption

or coprecipitation to the iron particles. Manganese and the light rare earth elements (LREE)

show a leaching trend through the entire profile. Arsenic and the heavy rare earth elements

(HREE) show a trend with accumulation in the transition zone (5 and 1 ppm). Aluminum

shows a trend with the highest concentration in the oxidation layer but instead of decreasing

downward it is stable between the transition zone and reduced layer. Lead is stable through

the entire profile.

Oxidation zone has a 1-2 mm rim, transition zone a 2-4 mm rim and reduced zone a 0.5-1 mm

rim. After storage for 18 days this gives an oxidation rate of 0.056-0.11 mm/day, 0.11-0.22

mm/day and 0.028-0.056 mm/day. The oxidation zone sample from profile 2 was already

oxidized when taking the sample even though it was more reduced than the OZ in profile 1.

Like in profile 1, the transition zone shows a higher oxidation rate which could indicate

minerals that oxidize faster than pyrite such as iron monosulfides but this has to be further

investigated since also grain size affects the oxidation rate. It looks like profile 1 is dominated

by framboidal pyrite but not excluding iron monosulfides and profile 2 is dominated by more

easily oxidized minerals like iron monosulfides but has some framboidal pyrite as well.

If comparing the concentrations of Mn and S in the two profiles one can see that it is lower in

profile 2 (840 ppm and 0.95 %, respectively) compared to the first profile (1480 ppm and 1.9 %,

respectively). Their trends are quite similar in both profiles and overlap each other. Profile 2’s

reduced samples are taken closer to the surface than those of profile 1 which could lead to

similar concentrations if taken deeper down even though the profile lies closer to the mean sea

level. When comparing element concentrations, one can see that many of the elements show

higher concentrations in the waterlogged profile. There is a difference of 1-20 ppm between

the oxidation zones for Zn, Co, Ni, Cu, V, and LREE. For Cr, As and Pb higher concentrations

can be seen in the more reduced parts of the profile. Of these elements Zn and Co have leached

out more from the OZ than are being accumulated in the transition zone, Ni has leached out

through entire profile 1 and Cu and V have leached out from the transition zone. Zinc, nickel

and cobalt show the best possibilities for transportation to profile 2 but this has to be

monitored before any conclusions can be drawn. Arsenic has been liberated but leaching too

little to be able of giving such an increase downstream, there could either be a naturally higher

concentration or more transported from other acid sulfate soils upstream the river. All other

elements are having naturally higher concentrations compared to profile 1.

30

4.2 Björsbyn (Site C)

4.2.1 Optical microscope and SEM

Samples from profile 1 are mostly light grey-grey colored where the sample from OZ is brown

in color due to iron hydroxides. All samples have a 0.5-3 mm brown-light grey oxidation rim

surrounding the main matrix. Profile 2 has black-grey colored samples with a 4 mm diffuse

oxidation rim in the oxidation layer and a thin rim partially surrounding the transition zone

sample. Profile 1 has large silicate grains that can be seen with the naked eye which becomes

fewer with depth.

Oxidation zone (profile 1)

The sample from the oxidation zone is covered with iron hydroxides that give a dark brown

color but no jarosite was found on the surface. A beige oxidation rim surrounds the polished

sample and a breccia with large silicate grains lies outside the rim. One can observe many

euhedral reflecting grains that could be silicate grains.

Transition zone (profile 2)

This zone contains smaller framboidal pyrites compared to site B that are evenly distributed

over the sample. Out in the oxidation rim the pyrite has been altered but it is not possible to

say how far the oxidation extends. Some parts of the sample contained thin layers with a

different color than the original matrix (Figure 25A).

Figure 25: Layering due to mobilization and sulfur concentrating in framboidal pyrite.

Figure 25A shows one of the layers that appears brighter than the surrounding matrix. When

comparing the chemistry in point B inside the layer with point C in the matrix one can see that

the layer has lower sulfur concentration than the matrix (Figure 25B and Figure 25C). One can

also see a line with framboidal pyrites that lies within the layer, indicating that sulfur has been

concentrated in framboidal pyrites (Figure 25A).

31

In one area of the sample one can see a typical cubic pyrite grain together with a framboidal

pyrite sphere (Figure 25D). Another observation in this sample was a zoned iron sulfide

(Figure 26).

Figure 26: Zonation of metastable iron sulfide.

Zonation in Figure 26 has most likely been formed as a result from oxidation where the bright

spots (B) are the intact iron sulfide. Further out in zonation (spectrum C and D), sulfur and

iron concentrations decrease while other elements increase their concentrations (e.g. Figure

26C). This grain has a diffuse morphology that could indicate dissolution. This makes it

impossible with only the chemistry to interpret if it is a pyrite grain or an iron monosulfide

and should therefore be called metastable iron sulfide until more investigation has been done.

Iron peaks are very small which supports pyrite as the correct mineral rather than a

monosulfide. The most interesting observation in this sample was euhedral grains of elemental

sulfur (Figure 27).

32

Figure 27: Elemental sulfur next to a framboidal pyrite grain. ES= elemental sulfur, Py=pyrite.

Figure 27A shows three (50 µm) grains of native sulfur where the spectrum in Figure27D

shows the single sulfur peak from the grain in focus in Figure27C. In Figure27C one can also

see that the grain is quite soft since the electron beam from the instrument has burned three

holes in it. This is only a small area with native sulfur of many with more sulfur grains.

Reduced zone (profile 1)

The reduced layer has fewer framboidal pyrites than observed in the transition zone. Many of

the observed pyrites show signs of oxidation, both in the matrix and in the oxidation rim

(Figure 28). Some iron hydroxides have precipitated both in the rim and in the matrix of the

sample.

Figure 28: Framboidal pyrites with oxidation surface (20 x ocular, optical microscope).

The few framboidal pyrites found in the sample had an orange color and a rim around them

(Figure 28). This rim could either be an oxidation rim or a growth rim like those seen at site B.

33

Oxidation zone (profile 2)

The sample is black with a 4-5 mm thick dark grey oxidation rim. Larger silicate grains are still

intact, and no iron hydroxides have been observed. Some framboidal pyrites are preserved

due to a higher groundwater table that limits oxygen penetration and slows down the

oxidation in this sample. Instead the pyrite grains had grown larger (Figure 29).

Figure 29: Framboidal pyrites, not densely packed (50 x ocular, optical microscope).

Like site B profile 2 the framboidal pyrites are seen in a cross-section and are not as densely

packed like they usually are. The tiny sparkling grains in the background (in B) could either

be iron monosulfides or some silicate grains.

Transition zone (profile 2)

Grey colored sample with a 0.5 mm oxidation rim only at the downside of the sample. Many

small (4-8 µm) framboidal pyrites were observed evenly distributed over the sample (Figure

30). There was at least 15-25 framboidal grains per 5x ocular view.

Figure 30: Distribution of (up to 6 µm) framboidal pyrites (black dots) (10 x ocular, optical microscope).

Reduced zone (profile 2)

Some large areas were covered with iron hydroxides but are most likely only surface staining.

In this sample, only 1-3 framboidal grains per 5x ocular were observed compared to 15-25 in

the TZ (Figure 31).

34

Figure 31: Framboidal pyrites (20x ocular, optical microscope).

Some of the framboidal pyrites show signs of orange alteration on the surface and have a thin

rim around the grains that could be a result of oxidation or growth. Grains here are more

densely packed than in the oxidation zone.

4.2.2 Geochemistry of the profiles

Soil samples from site C were taken in a thicket close to the water, behind “Stadsträdgården”

in Björsbyn. Profiles were taken ca. 5.5 m and 2.5 m above mean sea level with ca. 30 m apart.

Element concentrations for profile 1 can be seen in Figure 32 below.

In this profile, pH reaches down to 3.8 near the ground surface. In the sample from the

oxidation zone one can see lots of iron hydroxides but no jarosite. The low pH supports the

presence of jarosite like at site B and the first thought is that potassium is the limiting factor

but when comparing potassium concentration between site B and C (Table 13 and Table 14)

they differ very little. That leaves the limiting factor to sulfur concentration. Site B has 0.45 %

in the oxidation zone and site C < 0.2 % which then most likely is the reason jarosite has not

formed at site C profile 1. In the OZ nearly all elements have been liberated, due to the low

pH. One major trend that can be seen is that sulfur is leached out (0.95 %) and reprecipitated

further down (1.2 %). This means that at least 0.25 % of the sulfur is already in the TZ from the

beginning before a possible reprecipitation. Manganese (213 ppm and 1170 ppm, respectively),

the light rare earth elements (2 ppm and 54 ppm, respectively), molybdenum (0 ppm and 13

ppm, respectively) and arsenic (2 ppm and 1.5 ppm, respectively) follow this trend and

support that it has to be a naturally higher occurrence of elements further down than are

leached out from the OZ.

Figure 32: Element concentrations site C profile 1.

35

In this profile, most processes occur in the TZ. Mobilization and concentration of sulfur into

pyrite, oxidation forming zonation in iron sulfides and large native sulfur grains have been

observed in this sample. The pH in this layer is near 7 which leads to a redox potential (Eh)

between -0.35-(-0.25) V for pyrite to be stable according to the iron stability diagram (Figure

5.6) in (Ingri, 2012). The most peculiar is the presence of native sulfur in the TZ. When looking

at the sulfur stability diagram (Figure 5.7) in (Ingri, 2012) at pH 7 with the required redox

potential for pyrite to coexist, one can see that the stable sulfur species in this system should

be H2S and/or HS−. Elemental sulfur is only stable in a narrow redox window at pH 0-5. Only

other documented finding of native sulfur with SEM was in the sulfate-methane transition

zone of marine sediments in China (Lin et al., 2015), which is a different environment than in

this study. Some possibilities are that native sulfur is an oxidation product from metastable

iron sulfides (Bloomfield, 1972a), from oxidation of hydrogen sulfide or from reduction of

sulfate (Berner, 1984). Boman et al. (2008) had a similar discovery near Vasa, Finland with

elemental sulfur together with pyrite in the TZ although it was only measured through a

separation sequence. Their conclusion was that no metastable iron sulfides are present since

they would have reacted with the elemental sulfur and formed pyrite like proposed by e.g.

(Morse et al., 1987). Therefore, in this sample it had to be more elemental sulfur present than

it was metastable iron sulfide and all the latter have been converted into pyrite. The sulfur

content reaches up to 2.4 % in the TZ with both framboidal pyrite and native sulfur as observed

species. The reduced layer has a sulfur concentration right below 1.4 % and it is reflected on

the fewer framboidal pyrites observed in this layer compared to the TZ. For the reduced layer

that has a pH of 8.5, the redox potential could vary between -0.4-(-0.3) V for pyrite to still be

stable (Figure 5.6) (Ingri, 2012). A major trend seen is for iron where leaching (OZ 1.6 %) occurs

through the entire profile but less in the more reduced part. This can be compared to site B

where iron was precipitating as hydroxides in the oxidation zone instead of leaching. The

elements following this trend are Zn (OZ 56 ppm), Co (OZ 10 ppm), Ni (OZ 28 ppm), Cu (OZ

22 ppm), Cr (OZ 30 ppm) and V (OZ 30 ppm). Since iron and sulfur are not following each

other this could indicate that the majority of the iron is incorporated in other minerals than

sulfides hence the trend. The low pH makes aluminum and lead to leach out through the entire

profile. At site B these two elements had increasing concentrations upward instead of leaching.

The only element that increases upward in this profile is the heavy rare earth elements (HREE)

which could be due to incorporation in weather-resistant minerals or association to clay

minerals. Oxidation rims are 0.5-1 mm in the oxidation zone, 0.7-3 mm in the transition zone

and 0.7 mm in the reduced layer. With storage for 20 days, this gave an oxidation rate of 0.025-

0.05 mm/day, 0.035-0.15 mm/day and 0.035 mm/day for each layer. Element concentrations for

profile 2 can be seen in Figure 33 below.

36

Like profile 2 at site B, pH in this waterlogged profile is higher than in profile 1. At this site

pH is near 7 near the ground surface and it increases above 8 with depth. This makes it possible

to observe intact framboidal pyrites in the oxidation zone although the sulfur concentration is

as low as in profile 1. This could also explain why only larger grains are still intact. With a pH

close to 7 this gives a redox potential between -0.35-(-0.25) V for pyrite to be stable (Figure 5.6)

(Ingri, 2012). In this profile one can see the same trend as in profile 1 for sulfur (0.4 % and 1.5

%, respectively) together with Mn (110 ppm and 1063 ppm, respectively), Mo (0 ppm and 7

ppm, respectively) and As (0.5 ppm and 5 ppm, respectively), with leaching from the OZ and

accumulation in the TZ. This supports the theory of a layer in the transition zone naturally

enriched in sulfides which connects these two profiles. The amount of framboidal pyrites

correlates well with the sulfur content that reaches up to nearly 2.4 % in this layer. In the

transition zone pH gets close to 7.8 which gives a redox potential between ca -0.37-(-0.27) V

for pyrite to be stable (Figure 5.6) (Ingri, 2012). With depth the pH increases above 8 and with

it the sulfur concentration decreases down close to the same level as in the OZ. This sample

has also framboidal pyrite but fewer than the transition zone which is reflected on the lower

sulfur concentration. At this pH the redox potential in the reduced zone lies between -0.4-(-

0.3) V for pyrite to be stable (Figure 5.6) (Ingri, 2012).

One can see that iron leach out (OZ 1.8 %) like in profile 1 but small amounts are accumulated

in the transition zone (0.2 %) indicating participation in sulfide formation in the transition

zone. This is different than for profile 1 that showed no accumulation further down. The

leaching of iron without retaining concentration due to hydroxide precipitation could indicate

that this site contains less Acidithiobacillus ferrooxidans bacteria that oxidizes Fe2+ to Fe3+

(Nordstrom, 1982). This means Fe2+ is kept in solution and is transported downward until it

reaches the sulfur front and precipitates as sulfides or is transported out of the system where

it can precipitate as iron hydroxides. In this profile Ni, Cu, Pb and V follow the same trend as

iron did in profile 1 with only intense leaching in the OZ. Zinc together with the light rare

earth elements show a continuous leaching pattern from top to bottom while chromium shows

intense leaching in the OZ and TZ. Aluminum shows the same trend as it has done before with

increasing concentration upward due to either higher silicate content or association with

organic matter. The heavy rare earth elements have a trend with liberation from the TZ (2

ppm) and retaining their concentrations in the OZ and cobalt is stable or shows a tiny increase

upward throughout the entire profile. This could indicate incorporation into more weather

Figure 33: Element concentrations site C profile 2.

37

resistant silicate minerals that are containing both Al and HREE and possibly small amounts

of Co. Oxidation rims are 4-5 mm around the oxidation sample and 0.5 mm around the

transition zone. With 20 days storage, this gives a 0.2-0.25 mm/day oxidation rate in the

oxidation zone and a 0.025 mm/day oxidation rate in the transition zone.

If comparing element concentrations between the two profiles one can see that Zn (7 ppm and

28 ppm, respectively), Ni (10 ppm and 5 ppm, respectively) and Co (9 ppm and 1 ppm,

respectively) are higher in the OZ and TZ in profile 2. Copper is 2 ppm and vanadium is 5

ppm higher in the TZ and Cr is (20 ppm and 45 ppm, respectively) higher in the OZ and RZ.

All these elements are being leached out from profile 1. Of these elements it is possible that all

of them could be transported from profile 1 to profile 2 even though most of them also show

leaching from profile 2. Cobalt has an increasing concentration but in profile 2 conclusion is

that it probably is included in some weather resistant mineral making it less presumable that

the increase is due to transportation from the upstream profile. The other elements either show

no leaching trend upstream or show no increase in element concentrations.

4.3 Sunderbyn (Site A)

4.3.1 Geochemistry of profiles

Soil samples for site A were taken on the field across from the “Kläppen school” in Sunderbyn.

Profiles from site A were taken 6 m and 1 m above mean sea level with ca. 100 m apart. The

element concentrations in profile 1 can be seen in Figure 34 below.

It can be observed that pH is close to 4.5 in the OZ leading to intense liberation of mostly S

(1.45 %) but also Mn (235 ppm), Zn (12 ppm), Co (4 ppm), Ni (8 ppm) and Cu (2 ppm) are

liberated from this layer. Like at site B iron is only liberated from the transition zone (0.2 %).

This is due to precipitation of iron hydroxides in the oxidation zone limiting the loss compared

to the RZ. A similar trend was also observed in studies from Finland (Åström, 1998; Nordmyr

et al., 2008). Iron follows the trend of sulfur in the TZ and RZ which could indicate

incorporation in sulfides further down. Arsenic follows this trend and is likely to be associated

with the iron particles (Widerlund and Ingri, 1995). Lead is also following this trend, but are

commonly precipitating as anglesite (PbSO4) particles at least in an acid mine drainage (AMD)

environment (Dold, 2010). Soil studies instead suggest the possibility of adsorption to

ferrihydrite or organic matter (Gustafsson et al., 2011). Downward in the profile, pH increases

Figure 34: Element concentrations with depth from site A profile 1.

38

to near neutral stopping the liberation of several elements. At this pH one can see that 36 ppm

of the liberated manganese accumulates in the TZ. This could be related to precipitation of

manganese hydroxides/oxides such as MnOOH and/or MnO2 (Stumm and Morgan, 1996).

Elements that follow this trend and possibly are associated with these hydroxides are Zn (15

ppm), Co (2 ppm), Ni (4 ppm) and Cu (1 ppm), which is in agreement with studies by (Algeo

and Maynard, 2004) and (Shaw et al., 1990). Some of these elements (Zn, Co and Ni) that

accumulate in the transition zone are following earlier research results from this environment

(Åström, 1998). Since more Zn has been accumulated in the transition zone than released by

oxidation this could be explained by a naturally higher Zn content in this layer.

Chromium and the light rare earth elements (LREE) have a similar trend with accumulation

in the transition zone without liberation from the OZ. Aluminum shows a decreasing

concentration with depth which could be due to either more silicate minerals, precipitation of

gibbsite (Al(OH)3) or association with organic matter higher up in the profile. The only

elements that show a similar trend as Al are the heavy rare earth elements (HREE) since these

elements are stable as trivalent cations. When looking at the mass balance in Figure 38 one can

see that aluminum show a slight accumulation in the transition zone probably due to clay

mineral formation which then most likely Cr and the REE are related to. Molybdenum and

vanadium show also decreasing concentrations with depth which could indicate that these

elements are associated with weather-resistant minerals. Element concentrations from profile

2 for comparison are shown in Figure 35 below.

In the waterlogged profile, pH is close to 6.2 due to the higher groundwater table that limits

oxygen penetration and slows down the oxidation. All elements show leaching from the

oxidation layer despite the high pH. If comparing immobile element concentrations (Ti and

Zr) with depth (Figure 46) one can see that Ti shows a similar trend like Fe and S while Zr

follows Mn. This could indicate a dilution effect due to organic matter in the top layer, which

gives a false leaching result. Iron (1.7 %), sulfur (1.5 %) and arsenic (4 ppm) are strongly

liberated from the oxidation layer while they are stable further down in the profile. Manganese

(494 ppm) and molybdenum (3 ppm) show a similar trend but are leaching more in the

transition zone than the other elements. In the transition zone, pH is slightly decreased.

Figure 35: Element concentrations with depth from site A profile 2.

39

Aluminum is accumulated (0.3 %) in this layer which could indicate precipitation of aluminum

particles like gibbsite (Al(OH)3) or formation of clay particles that gives a higher concentration.

One can also see that Zn (24 ppm), Co (4 ppm), Ni (7 ppm), Cu (1 ppm), Cr (7 ppm), V (3 ppm),

the LREE (18 ppm) and the HREE (1 ppm) are following this trend. Zn, Co, Ni and sometimes

Cu are commonly trapped in the transition zone (Åström, 1998; Nordmyr et al., 2008). The REE

are commonly associated with clay mineralogy and therefore would an enrichment support

such a hypothesis (Åström et al., 2010; Harlavan et al., 2009).

If looking at each element and compare the profiles one can see that only Co, Ni and Cr reach

higher concentrations in profile 2. Only Co and Ni of these show a leaching pattern with an

excess of mobile concentrations from the adsorption in the transition zone. This could mean

that these elements are transported downstream and accumulated in the waterlogged profile.

Due to the dilution effect, this is only kept as a possibility.

4.4 Björsbyn (Site D)

4.4.1 Geochemistry of profiles

Site D lies on a small meadow within 1 km from site C in Björsbyn. Profiles were taken 3 m

and 2 m above mean sea level with ca. 30 m apart. Element distribution through profile 1 can

be seen in Figure 36 below.

At this site, pH is higher (⁓ 4.8) near the ground surface than further down where it reaches 3

before increasing toward neutral pH. Like at site C sulfur leaches out (0.7 %) and is

accumulated in the transition zone (1.3 %). This backs up the theory of a sulfide-rich layer in

the transition zone. Site C and D are connected through this layer. Elements that follow sulfur

are Mo (0 ppm and 7 ppm, respectively), As (3 ppm and 9 ppm, respectively) and the LREE

(12 ppm and 40 ppm, respectively). This is very similar to site C where also Mn followed this

trend. Chromium shows minor accumulation in the transition zone (2 ppm) that could indicate

incorporation into sulfides. Iron has a trend where it leaches quite intense in the oxidation

layer (1.1 %) and further down stabilizes. which is similar to the trend seen at site C. Cobalt

and nickel follow the trend of iron (10 ppm and 24 ppm, respectively) which could be due to

incorporation in pyrite (Huerta-Diaz and Morse, 1992; Scholz and Neumann, 2007). Zinc,

copper and manganese show a continuous leaching trend (OZ 74 ppm, 26 ppm and 388 ppm,

respectively) through the entire profile. This could indicate that Zn and Cu are associated with

manganese particles and when pH is lowered they are dissolving making them go into

Figure 36: Element concentrations site D profile 1.

40

solution. Another trend is aluminum that is intensely leached out. This time lead and

vanadium follow this trend indicating that they could be associated with clay particles.

Element concentrations for profile 2 can be seen in Figure 37 below.

This waterlogged profile has like profile 2 at the sites A and C a pH above 6 in the OZ. The

transition zone is missing in this profile and only a very thin (10 cm) OZ is present which

makes it hard to interpret the trends. One can see one dominating trend in this profile, with

increasing concentrations upward which could be due to association with organic matter

(Nordmyr et al., 2008; Scholz and Neumann, 2007; Shaw et al., 1990) and some of them possibly

due to incorporation into weather-resistant minerals (Mo, Pb, V, HREE). The other trend to see

is one with minimal leaching and/or stable trend. In general, this profile has very low sulfur

content through the entire profile. This may indicate that the profile had oxidized and now are

reduced again although sulfides are not formed again as it did in Littorina time (Virtanen et

al., 2014). One possibility that would be great to investigate further is if the sulfur

concentration increases further down beyond the profile depth. This to determine if this profile

has been an old oxidized layer and the real reduced layer is located even further down. If

checking the trend for Zr and Ti (Figure 47), they follow the stable/minimum leaching trend

which means that this profile may have been affected by organic dilution like profile 2 at site

A.

Since profile 2 is without a transition zone, trends are not possible to completely compare

between the two profiles. Comparing the waterlogged profile concentrations to those of the

oxidation profile, one can see that Zn, Co and Ni concentrations are 2-2.5 times higher in the

oxidation layer and vanadium is 16 % higher. This could mean that these elements have been

transported from profile 1 and been accumulated in profile 2 like the case with accumulation

in the TZ presented by (Åström, 1998). The total concentrations of Cu, Al, Cr and Pb are higher

through entire profile 2 where the RZ in profile 2 has similar concentrations as the OZ and/or

the TZ in profile 1. Remaining elements either show no leaching trend upstream or are in

general lower downstream such as sulfur that is much lower than in profile 1.

Figure 37: Element concentrations site D profile 2.

41

4.5 Geochemical comparison

Site C has the highest Mn (1900 ppm and 1700 ppm, respectively), S (2.3 % and 2.15 %,

respectively), Ni (43 ppm and 41 ppm, respectively) and V (103 ppm and 103 ppm,

respectively) content of all sites and the lowest of As (5.3 ppm and 5.7 ppm, respectively).

Profile 1 has the highest Mo content (15 ppm) and the lowest Cr (65 ppm) and REE (LREE 119

and HREE 9 ppm) content.

Profile 2 has the highest Cr (139 ppm) and lowest Mo (1.2 ppm) content. Since a transportation

trend of Cr was seen from profile 1 (Figure 32) and an increase in concentration occurred

downstream (Figure 33), it strengthens the probability for Cr to has been transported, leaving

profile 1 with the lowest concentration of this element. Profile 1 at Site D shows the lowest

concentrations of Al (4.9 %), S (0.07 %), Zn (48 ppm), Co (6 ppm), Cu (8 ppm), Ni (12 ppm), Pb

(13 ppm) and V (66 ppm) all in the OZ. In the TZ it has the highest concentration of As (18

ppm) and in the RZ the highest concentrations of Zn (122 ppm) and Pb (21 ppm) where the

latter ones possibly have been transported downward in the profile and enrich the reduced

layer while depleting the oxidation layer of those elements. Profile 2 shows enrichment of Cu

(40 ppm) in the OZ making it possible to has been transported downstream, depleting it in the

oxidation profile. This profile also has the highest concentrations of rare earth elements

(253+21 ppm) compared to the other profiles, located in the OZ. The oxidation zone in profile

1 at site A has been enriched in aluminum (6.69 %). The transition zone at profile 2 is enriched

in cobalt (18 ppm) while the oxidation zone is depleted in iron (2.62 %) and manganese (439

ppm) compared to the other sites.

When comparing these chemical trends, note that the aluminum concentration mostly

increases upward despite many papers highlighting aluminum leaching out, leaving high

concentration in waterways e.g. (Filppa, 2012; Nordmyr et al., 2008). When comparing this

trend with the mass balance diagrams in 7.2 to denote if this is true, one can see that this is in

most cases not true. This could be due to complications during the analysis with precipitation

of aluminum minerals as a result or loss of volatile elements concentrating the solution in

aluminum. To investigate this further an x-ray diffraction (XRD) analysis has to be done to see

what matrix mineralogy the samples have. For example, profile A1 Figure 34) shows that Al

increase upward while in Figure 38 it is believed that aluminum accumulates in the transition

zone which may indicate the formation of clay particles that Cr and LREE adsorb to. In most

profiles, the mass balance shows a trend of leaching for Al other than in D1, B2 and A1 where

it is accumulated in the transition zone or is leached from the transition zone (section 7.2). A

possible reason for aluminum to plot with increasing concentration upward is the association

to organic matter (Nordmyr et al., 2008).

In profile 1 at site D (Figure 36) nearly all elements including Fe and Al are being liberated

from the OZ but according to the mass balance (Figure 44) they should instead be accumulated

in the transition zone like sulfur. At site B lead increases upward in both profiles (Figure 23

and Figure 24) which could indicate association with organic matter but the mass balance

(Figure 40 and Figure 41) suggest an association with iron hydroxides in the oxidation layer.

Profile B1 (Figure 23) shows accumulation of Zn, Co and Al in the TZ, mass balance (Figure

42

40) instead shows liberation of these elements although Zn and Co still follow the Al trend. In

profile 2 at site B (Figure 24), Co and Cr follow sulfur while the HREE are accumulated in the

TZ. The mass balance (Figure 41) suggests that Co and Cr follow iron and that the HREE

leaches through the entire profile instead. In profile A2 lead is slightly accumulated in the TZ

but in mass balance (Figure 39) it increases upward in the profile. In profile A1 (Figure 34), Mn

and Cu show a slight accumulation in the transition zone and V increases upward but the mass

balance (Figure 38) suggest that they all are leached out through the entire profile.

In profile 1 at site C (Figure 32), the HREE have a trend with increasing concentrations upward

which is according to the mass balance (Figure 42) instead an accumulation in the TZ. In profile

2, the HREE are instead thought to be leached only from the TZ and Co to be relatively stable

through the entire profile. According to (Figure 43) the HREE are leaching through the entire

profile and Co follows iron with leaching only in the TZ. Mass balance is a good way of

correcting errors that are probably related to interference with organic matter since the ratio is

based here on zirconium concentration that is thought to be immobile. This despite the fact

that zircons have been found that are clearly affected by weathering in the low pH.

4.6 Classification

Since potential acid sulfate soil (PASS) contains iron sulfides that easily oxidize when exposed

to oxygen, pH drops fast. This effect can be seen in Table 1 when comparing field pH with pH

after oxidation. The pH in the oxidation zone in profile 1 at site D was taken on a different

depth than the actual sample was taken and therefore not possible to compare.

Table 1: Table over pH measured in the field and after oxidation in the lab.

If comparing sulfur content to where the pH has changed the most in each profile one can see

that at sites A, C and profile B1 the pH has been lowered the most where sulfur content is

highest. This probably follows directly to sulfide content in the profiles especially since all

samples were collected at different dates and have therefore been oxidized for a different time.

Profiles B2, C2 and D2 have a trend where the pH shows a large drop even though the sulfur

content is low. This could mean that the remaining sulfur is included in iron monosulfides that

oxidizes within hours when exposed to oxygen. Profile D1 shows a decreasing pH trend with

depth which is not correlating with the sulfur trend. This could mean that sulfur minerals with

slower oxidation rate than iron monosulfides occur in the transition zone rather than in the

reduced layer. This since pH has dropped faster in the reduced layer.

43

When classifying each sample according to the SGU-standard (section 2.6) the OZ from

profiles B1-D1 are classified as AASS (field pH< 4). The OZ from profile A1 and the TZ from

profiles A2 and D1 are classified as semi-acid while the TZ and the RZ of profiles A1-B1 are

classified as reduced. The RZ of profiles A2 and C2, the TZ of profile C1 and the OZ of profiles

C2-D2 also classified as reduced. Some layers did not reach pH<4 after oxidation and are

therefore not classified as acidifying but are most likely sulfidic due to large pH drops. When

classifying each profile according to the GTK-standard presented by (Edén et al., 2012) where

starting depth of PASS, minimum field pH and also the maximum sulfur content are taken

into account the result follows in Table 2.

Table 2: Classification according to GTK system.

Class 1 means the PASS horizon starts within 1 m and class 2 within 1-1.5 m from the ground

surface. When classifying according to minimum field-pH class D represent a pH > 4.5, class

B a pH between 3.5-3.9 and class A pH < 3.5. Total sulfur content is taken for the uppermost

40 cm of the PASS horizon where (I) are representing a total sulfur content > 1 % and (III) a

content between 0.2-0.6 %. Pousette et al. (2007) propose classifying the acidic potential of

samples with Fe/S ratio, leading to this study’s acid potential presented in Table 3.

Table 3: Acidic potential classification.

If the Fe/S ratio is < 3 the soil is very acidifying and > 60 non-sulfidic. This means that the

transition zone at site C and profile D1 have the potential of being very acidifying and site A

close of being. Many of the reduced layers are also showing a high potential of being

acidifying. All oxidation layers are showing a low potential of being acidifying which is since

most of the sulfur has already leached out from them.

44

5 Conclusion Framboidal pyrite was observed at both site B and C in all samples except for the oxidation

zone samples from profile 1. This is due to extensive oxidation that leads to a low pH. In these

samples secondary iron hydroxides precipitate. In profile 1 at site B the secondary mineral

jarosite was observed amongst the hydroxides but was not observed at site C although the

sample had the potential of forming it. The low sulfur concentration is thought to prohibit this

formation at site C. Barite was only observed at site B (profile 1 TZ-RZ and profile 2 OZ) but

if site C had been further investigated with SEM it might have been found also there. Elemental

sulfur was only found in the TZ at site C and although site B was thoroughly investigated for

this specie no other observation was achieved. This may indicate either the formation or

oxidation of iron sulfides. Different morphologies and stages of evolution of framboidal pyrite

were observed in the samples, where most variation occurred at site B.

A combination of iron sulfides is thought to be present at site B even though the RZ looked

pyrite dominated and although no iron monosulfides were observed they cannot be excluded.

This also gives rise to the debate about whether iron monosulfides gives the black color or just

pyrite since the black lines were filled with framboidal pyrites. The waterlogged profiles retain

the pH high enough for framboidal pyrite to still be intact and in profile 2 at site D it would be

interesting to see if the sulfur concentration increases further with depth. According to the

oxidation rate, site B oxidizes faster than site C on a small scale which could indicate different

sulfide mineralogy. At site B and in profile A1 iron is retained due to hydroxide precipitation

in the OZ while it leaches out at site C and in profile D1. This could indicate a lower activity

of oxidation bacteria at these two sites that prevents precipitation of iron hydroxides. Sulfur

leaches out from sites A and B while accumulating in the TZ at site C and in profile D1. This

is directly reflected on the amount of framboidal pyrites observed at site B and C. The usually

related sulfide elements Zn, Co and Ni show either accumulation in the transition zone like in

other research (Åström, 1998), relation to iron or total liberation through most profiles.

Elements that could be associated with iron hydroxides at different sites are As, Pb, Cu, V, Zn,

Co, Ni and Mo. Many of the elements show the same trend in both profiles at a location. At

site B and C where pH is low most elements not associated with hydroxides, heavy minerals

or organic matter are liberated with large intensity. Aluminum is in many profiles most likely

associated with organic matter. Elements that could be transported and accumulated

downstream are Co, Ni, Zn and Cr and at some of the profiles also Cu and V.

5.1 Future work

In the future, monitoring geochemical parameters and continual groundwater sampling to

measure mobile element concentrations are recommended to increase the knowledge of mass

fluxes in the system. Addition of tracers to the groundwater is proposed to study if elements

travel downstream to the more waterlogged part and conclude if it accumulates further out.

Other work could be to investigate elemental sulfur closer since it has been observed in SEM

at site C and therefore may do some more experimental studies to investigate its stability and

evolution further. Further studies are needed to investigate the parent material in order to

understand if elements are liberated from either sulfides or the silicate matrix.

45

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7 Appendix

7.1 Appendix 1: Profile descriptions

Table 4: Coordinates from profile sites.

Table 5: Profile description site A profile 1.

Table 6: Profile description site A profile 2.

Table 7: Profile description site B profile 1.

Table 8: Profile description site B profile 2.

52

Table 9: Profile description site C profile 1.

Table 10: Profile description site C profile 2.

Table 11: Profile description site D profile 1.

Table 12: Profile description site D profile 2.

53

7.2 Appendix 2: Mass balance

Figure 38: Mass balance site A profile 1.

Figure 39: Mass balance site A profile 2.

54

Figure 40: Mass balance site B profile 1.

Figure 41: Mass balance site B profile 2.

55

Figure 42: Mass balance site C profile 1.

Figure 43: Mass balance site C profile 2.

56

Figure 44: Mass balance site D profile 1.

Figure 45: Mass balance site D profile 2.

57

7.3 Appendix 3: Immobile element vertical distribution

Figure 46: Zirconium and titanium distribution site A profile 2.

Figure 47: Zirconium and titanium distribution site D profile 2.

7.4 Appendix 4: Whole rock geochemical data

58

Tab

le 13: Sam

ple ch

emistry

from

total d

igestio

n an

d m

easured

with

ICP

-SF

MS

. Green

are recom

men

ded

valu

es by

AL

S to

use fo

r interp

retation

.

59

Tab

le 14: Sam

ple ch

emistry

from

fusio

n an

alysis an

d m

easured

with

ICP

-SF

MS

. Green

are recom

men

ded

valu

es by

AL

S to

use fo

r interp

retation

.