Acid Pollution of Natural Waters

Outline of Topics• Background

– Acids and bases– pH ranges in the hydrosphere– Biological effects of acid pollution

• Sources of Acid Pollution– Acid deposition– Acid mine drainage

• Resisting Acidification– pH buffers– Dissolved inorganic carbon in natural waters– Measuring buffering capacity– Homogeneous and heterogeneous buffering

Acids and Bases• Lecture Questions

– What is an acid? What is a base?

– What is pH?

– Explain the difference between strong and weak acids/bases.

pH Ranges in the Hydrosphere

pH of Natural Waters: Lakes

Sources of Acid Pollution• Questions

– What pollutants cause acidification of natural waters, and how?

– What human activities generate these pollutants?

• SO2 and NOx emissions

– Generated by fossil fuel combustion (SO2 especially by coal burning)

– Atmospheric oxidation of SO2 leads to sulfuric acid and sulfate aerosol

– Atmospheric oxidation of NO leads to nitric acid and nitrate aerosol

– Acidification occurs during wet and dry deposition and due to spring snowmelt

• Acid mine drainage– Mining operations generate large quantities of mine waste

– Products of the aqueous oxidation of exposed minerals (mostly FeS2) in the waste leads to acidification of ground and surface water

Effects of Acidification on Aquatic Ecosystems

• Acute effects– Interferes with osmoregulation– Mobilization of toxic metals (Al, Mg, Zn)

• Chronic effects: reproduction problems– Lowers calcium levels in female fish, hinders ability to produce eggs– Fertilized fish eggs may develop abnormally– Frog and salamander eggs can be greatly affected by spring snowmelt

• Aquatic ecosystem effects– Can result in decline and algal species diversity and biomass– May effect ecosystem productivity– May decrease decomposer populations, affecting nutrient cycling and

increasing DOM/POM levels

Effects of Acid Pollution on Terrestrial Ecosystems

• Direct Effects– Damage to foliage

• Effects on Soil– The nature of soil– Soil pH– Loss of nutrients and other metals– Loss of ability to retain nutrients

Acid Mine Drainage• Picture to the left shows acid mine

drainage in tar creek, OK (at a superfund site)

• Mining operations expose minerals and mine waste to air and water

• Acid mine drainage is mostly due to the biologically-mediated oxidation of pyrite, FeS2

• Color is due to precipitation of oxidized iron (eg as Fe(OH)3)

Mechanism of Pyrite Oxidation• Overall process

)(SOH2)(Fe(OH))(O4

15)(OH

2

7)(FeS 423222 aqsgls

16H14Fe2SOOH814FeS 2242

322

O7H14Fe14HO2

714Fe 2

32

2

32 3Fe 3H O Fe(OH) (s) 3H

• Direct oxidation of S22- by O2 is too slow; instead it is oxidized by something else, usually Fe3+ that is present naturally in the environment:

• The Fe3+ is regenerated by reaction with dissolved oxygen– The Fe3+ is a catalyst for the generation of acid mine drainage

– This is the rate determining step of the entire process; it is an exothermic reaction

• Most acidity in the mine drainage is due to oxidation of disulfide anion to sulfuric acid, but some is due to hydrolysis by FeIII, for example:

Iron Cycling Controls Rate of Acid Generation

• Oxidation of disulfide is catalyzed by Fe3+: regeneration of catalyst is necessary

• Each cycle generates 2 protons of acidity

• Abiotic oxidation of Fe2+ by O2 is fairly slow ( = 1000d)

• Fe2+ oxidation is biologically mediated by bacteria who derive energy from the reaction; they are acidophiles who thrive in low pH environments. Reaction is 106 times faster when they are present.

3+14 Fe 2+14 Fe

2-2S

+14 H

+16 H

Thiobacillus Ferrooxidans

oxidation of disulfideproduces acidity

regeneration of Fe3+

Carbonate Chemistry in Natural Waters• Why is calcium bicarbonate, Ca(HCO3)2(aq), the dominant solute in most freshwaters

(so-called “calcareous” waters)?– Imagine a water body (lake, river, etc) in equilibrium with both calcite (the most common

carbonate mineral) and the atmosphere.

2+32 3 2 ( )CO ( ) CaCO ( ) H O( ) Ca ( ) 2HCO aqg s l aq

2H O

2 2CO ( ) CO ( )g aq 2H O 2+ 2-3 3CaCO ( ) Ca COs

2 2 2 3CO ( ) H O H CO ( )aq aq 2- -2 3 3 3H CO ( ) CO ( ) 2HCO ( )aq aq aq

Dissolution Processes

Hydration of CO2 Acid-Base Reaction

Water in equilibrium with atmospheric CO2 and CaCO3: pH = 8.3, [HCO3

-] = 1mM, [Ca2+] = 0.5mM (at 25oC).

Overall process: heterogeneous reaction between an acid and a base

Dissolution/evolution of gaseous carbon dioxide

Dissolution/precipitation of calcite (or another carbonate mineral)

Dissolved carbon dioxide reacts to give carbonic acid.

Carbonic acid from the air reacts with carbonate (a base) from mineral dissolution to give bicarbonate!

Ocean Acidification• What is ocean acidificaton?

– pH has decreased from 8.24 to 8.13 over the last 250 years.– Why?

• Carbonate chemistry! Atmospheric CO2 has increased since 1750.

• Some of the CO2 dissolves in the oceans and reacts with carbonate.

• So what?– Represents an increase of almost 30% in H+ concentration.

• Remember that pH is a logarithmic scale.

– If it continues unchecked it can greatly affect calcifying organisms.

• These build shells of CaCO3

• Their shells can start to dissolve if the oceans become too acidic.– Some scientists predict this could happen on a wide scale in 10-20

years if CO2 emissions go unchecked.

Resisting Acidification: Buffering Acid Pollution• What is a buffer?

– A (pH) buffer is a solution that can absorb a “large” amount of acid or base with a relatively small change in pH

• How do they work?– Basically: added protons or hydroxide react with something

instead of just changing pH

• Buffering in natural systems– Soils and natural waters are naturally buffered against pH

change (some more so than others)– It is possible to overwhelm – temporarily or “permanently” – the

buffering ability of a water body– The carbonate system is the most important buffering

mechanism in natural waters• Carbonate system: all forms of dissolved inorganic carbon (DIC)

Buffering in Natural Waters• Question

– How are natural waters buffered against changes in pH?

– Some background:• The ability to resist acidification: pH buffering

• Soils and natural waters are naturally buffered against pH change (some more so than others)

• It is possible to overwhelm – temporarily or “permanently” – the buffering ability of a water body

• The carbonate system is the most important buffering mechanism in natural waters.

– The ability to buffer is quantified by the buffer index.• Determined from titration curves.

• Next slide illustrates by showing the titration curve of a carbonate salt

mmol added acid

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5

pH

2

4

6

8

10

12

Buf

fer

Inte

nsity

, mm

ol L

-1 p

H-1

0.0

0.5

1.0

1.5

2.0

2.5

Quantitative Assessment: Buffer Intensity

Titration curve for 1L of1mM carbonate solution

bicarbonatebuffer region

bufferintensity

Carbonate Buffer Intensity (CT = 1mM)

pH

2 3 4 5 6 7 8 9 10 11

Buf

fer

Inte

nsity

( )

, mm

ol L

-1 p

H-1

0.01

0.1

1

10

H+ bufferingcarbonate & OH-

buffering

bicarbonatebuffering

• Buffering of natural waters is often specified by measurements of its alkalinity– Alkalinity: the amount of acid that must be added (per unit volume) to

reach a pH of 4.5 (the pH when all bicarbonate has been protonated)– In a carbonate-only system, alkalinity is determined largely by the

concentration of the dissolved inorganic carbon (DIC)

• Typical values in natural waters are 1 mmol/L – in other words, it takes 1 mmol of protons to convert all the DIC in a 1L sample completely to CO2(aq)

• The above expression does not consider non-carbonate sources of alkalinity

• Alkalinity measurements typically only measure homogeneous buffering ability of the natural water

Resisting Acidification: Buffering Acid Pollution

23 3[ ] [HCO ] 2[CO ] [OH ] [H ]alk

Homogeneous Buffering of Natural Waters• Homogeneous buffering is the neutralization of added

acid/base by dissolved species– Mostly due to the bicarbonate buffer system:

• DIC (alkalinity) is about 1 mM in fresh water and about 2 mM in marine water.

• Borate species (75 M in seawater) provides approximately 10% of the buffering ability in marine waters

• Silicate is widespread in all natural waters (20 M in seawater, 100 M in rivers) and can be an important buffer in low-carbonate freshwater

3 2 3HCO ( ) H H CO ( )aq aq

Heterogeneous Buffering of Natural Waters

• Heterogeneous buffering is buffering due to the dissolution or evolution/precipitation of a gas or a solid

– Slower but more powerful than homogeneous buffering

• Common mechanisms:– Evolution/dissolution of gaseous CO2

3 2 2HCO ( ) H CO ( ) H Oaq g

– Precipitation/dissolution of calcite mineral

– Acid weathering of silicates (proton displacement) will occur over longer time periods. It is primarily important when carbonate minerals have been intensely weathered (ie, when they start disappearing).

23 3CaCO ( ) H Ca ( ) HCO ( )s aq aq

Heterogeneous vs Homogeneous Buffering

23 2 2CaCO ( ) 2H CO ( ) Ca + H Os g

• Homogeneous is much faster than heterogeneous buffering

• Buffering due to silicate weathering is slower than due to calcite and CO2 dissolution/formation

• Heterogeneous buffering ultimately provides a greater capacity to assimilate protons (ie, neutralize added acid)

• Neutralization of added acid by carbonate minerals continues until all dissolve and ultimately generate CO2:

Summary: Resisting Acidification• Carbonaceous waters

– The carbonate system (ie, DIC, carbonate minerals, and carbon dioxide) provides buffering in most natural water bodies

– Carbonate buffering occurs whenever water is in contact with air and with carbonate minerals

– Homogeneous buffering due to DIC is the “immediate” response to acidification

– Heterogeneous buffering provides long-term buffering

• Long-term acidification– Would eventually overwhelm the carbonate system

• No more homogeneous/heterogeneous buffering involving the carbonate system

• All carbonate minerals would dissolve (weather)• Then all carbonate and bicarbonate would be transformed to carbonic acid

and CO2

– Silicate buffering can then occur• But it is slow, and silicate minerals weather too• Eventually transformed to quartz, which does not weather (or buffer) well.

Buffering of Soils (?)