Chemistry of the Hydrosphere

• Inherent properties that do not rely on chemical reaction• Temperature, density

• Color, taste, odor• Color due to presence of complex array of large organic molecules (fulvic and humic acids from

degradation of vegetative matter) complexed to certain metallic species

• Main metallic contributors to color are Fe3+ and Mn2+

• Turbidity, conductivity• Turbidity is a measure of the light scattering ability of a water sample due to the presence of

suspended particles – loosely correlated with total suspended solids

• Units: Nephelometric Turbidity Units (NTU)

• Drinking waters: < 1 NTU

• Rivers/streams: ~1 – 10 NTUs

• Storm events/erosion > 500 NTUs

• Lakes Wide range depending on “biological productivity”

Physical Properties of Natural Waters

• Conductivity is a measure of the ionic strength of a water sample

• Units: microSiemens (µS) cm-1

• Typically used as surrogate for total dissolved solids (TDS)

• High conductivity can be indicative of water contamination

• Total solids (dissolved, TDS, and suspended, TSS)• Typically estimated from conductivity

• Units: mg L-1 NaCl

• Conversion between conductivity and TDS is specific to water source, but good estimate is 0.65 mg L-1

NaCl per µS cm-1

• Distilled water: 0.5 – 30 µS cm-1

• Melted snow: 2-42 µS cm-1

• Potable water: 30 – 1500 µS cm-1

• Freshwater streams: 100 - 2000 µS cm-1

• TDS meters are actually conductivity meters in disguise

• Oxidation-reduction potential (ORP)• Measure of the oxidizing ability of a water sample

• Affects the speciation of many chemical constituents

• Typically measured with a Pt wire electrode

• Units: mV (vs SHE)

• Anaerobic (reducing) environment ORP < -200 mV

• Aerobic (oxidizing) environment ORP > 0 mV

• Properties dependent on chemical reaction• Alkalinity

• Two main definitions:

• Total alkalinity – total acid neutralizing capacity of a water sample

• Carbonate alkalinity – acid neutralizing capacity due only to carbonate buffer system

• Units:

• dKH, mEq CaCO3, ppm (mg L-1) CaCO3

• Hardness• Measure of important cations in samples of natural waters

• Hardness Index = [Mg2+] + [Ca2+]

• Units:

• ppm (mg L-1) CaCO3

• Natural waters contain dissolved and suspended materials that affect and are affected by water chemistry• Presence of different species is largely dependent on the water source, soil composition,

pH, human inputs, etc.

• Composition of typical rain water:

• Low dissolved organic carbon content (DOC < 1 mg L-1)

Composition of Rain Water

Ion SO42- Cl- NO3

- Na+ K+ Mg2+ Ca2+

ppm 2 8 0.5 4 0.3 0.3 2

mEq L-1 0.021 0.22 0.17 0.17 0.008 0.012 0.050

• Relatively high concentrations of dissolved gases:• O2: KH = 769.23 atm M-1

𝑑𝑂2 =𝜒𝑂2𝑝𝑡𝑜𝑡

𝐾𝐻=

0.21 1 𝑎𝑡𝑚

(769.23 𝑎𝑡𝑚 𝑀−1)= 2.7 𝑥 10−4 𝑀 = 0.27 𝑚𝑀

2.7𝑥10−4𝑀32𝑔

𝑚𝑜𝑙= 8.7 𝑚𝑔 𝐿−1

• CO2/H2CO3 ~ 1.32 x 10-5 M; 0.58 mg L-1

• SO3, NO2 ~ trace (but enough!)

• Remember, dissolution is a f(T, ionic strength)

Dissolved Gases in Rain Water (Reminder)

• Composition depends on• Rain constituents

• Soil/mineral substrate

• Plant/animal life

• Human activity

• Major cations include: Ca2+, Mg2+, Na+, K+

• Major anions include: HCO3-/CO3

2-, SO42-, Cl-, NO3

-, PO43-

• TDS ~ 20 – 150 mg L-1

• DOC ~ 1 – 10 mg L-1

• Human activity impacts include:• Agriculture (fertilizers), logging, mining, urbanization, …

Rivers and Streams

• Composition depends on:• Inflow constituents

• Residence time

• Plant/animal activity

• Human activity

• Deep, temperate lakes are characterized by stratification• Effects chemical speciation and chemical/biological processes

• Epilimnion – high dO2 (oxidizing)

• Thermocline

• Hypolimnion – low dO2 (reducing)

Lakes and Reservoirs

• Lakes classified as • Oligotrophic

• Very low concentrations of nutrients, such as nitrates, iron, phosphates, and carbon sources

• Mesotrophic• Intermediate level of productivity

• Eutrophic• High biological activity due to high concentrations of nutrients

• Composition depends largely on soil/rock type• Filtered through soil, sand and clay

• Usually low in micro-organisms

• Low dO2, reducing (ORP < -200 mV)

• TDS > 100 mg L-1

• DOC < 1 mg L-1

• Typically higher ion concentrations than in surface waters

Groundwaters

Ion SO42- Cl- NO3

- Na+ K+ Mg2+ Ca2+

ppm 1130 41 6 124 17 19 336

ppm 710 6 --- 6 5 44 258

Box Canyon Springs, CO1

Banff Springs, BC2

2 Grasby & Lepitzki, Can. J. Earth Sci., 39 (2002), 1349-1361.

1 https://www.boxcanyonouray.com/hot-springs/chem-analysis - accessed March 19, 2021

• Very high ion concentrations• In contrast to fresh surface waters, Na+ and Cl- are major ions (rather than Ca2+ and HCO3

-)• [Na+] ~ 12,000 ppm

• [Cl-] ~ 18,000 ppm

• TDS ~ 35,000 ppm (3.5%)

• DOC ~ 1 mg L-1 Carbon

• pH ~ 8.2

• Remember: At these high ion concentrations, ions no longer behave as “free” and “independent” chemical species

Marine and Ocean Waters

• Biological Oxygen Demand (BOD) – equals the amount of dO2 consumed as a result of oxidation of organic matter from natural sources (eg., sugars, fats, proteins, carbohydrates)• Also called Biochemical Oxygen Demand (BOD)

• Typical BOD for unpolluted surface water ~ 0.7 mg dO2 L-1

• Much lower than maximum dO2 of water at 25 oC

• Typical BOD for sewage typically several 100 mg dO2 L-1

• Soon depleted unless water is continuously aerated

• Chemical Oxygen Demand (COD) – equals the amount of dO2 consumed as a result of oxidation of organic (eg., petroleum, solvents, cleaning agents) and inorganic (ammonia, nitrite) matter from chemical sources

• COD• A closed water sample is incubated with a strong chemical oxidant under specific conditions

of temperature and for a particular period of time. A commonly used oxidant in COD assays is potassium dichromate (K2Cr2O7) which is used in combination with boiling sulfuric acid (H2SO4).

• Because this chemical oxidant is not specific to oxygen-consuming chemicals that are organic or inorganic, both of these sources of oxygen demand are measured in a COD assay.

• BOD• The sample is kept in a sealed container fitted with a pressure sensor.

• A substance that absorbs carbon dioxide (typically lithium hydroxide) is added in the container above the sample level. Oxygen is consumed and carbon dioxide is released. The total amount of gas, and thus the pressure, decreases because carbon dioxide is absorbed by LiOH.

• From the drop of pressure, the sensor electronics computes and displays the consumed quantity of oxygen.

Measuring COD and BOD

• What happens in an anaerobic water environment?• Bacteria serve to decompose dissolved organic matter

• Fermentation – chemistry where both oxidizing and reducing agents are organic materials

2𝐶𝐻2𝑂 𝑎𝑞bacteria

𝐶𝑂2 𝑔 + 𝐶𝐻4(𝑔)

Hydrological Cycle

Plain font: Reservoir estimate, 1000 km3

Italics: Flow estimates between reservoirs, 1000 km3/yr

(97%)(1.1%)

1.9% Five main compartments:1) Oceans2) Ice/snow3) Atmosphere vapor/clouds4) Surface water5) groundwater

Trenberth K.E., et al., J. Hydrometeor., 8(4), 758-769 (2007)

Major Aquatic Chemical Processes

• Most important aqueous weak acid

𝐶𝑂2 𝑎𝑡𝑚𝑜𝑠𝑝ℎ𝑒𝑟𝑒 ⇌ 𝐶𝑂2 𝑎𝑞 𝐾𝐻 = 29.41 𝑀 𝑎𝑡𝑚−1

𝐶𝑂2 𝑎𝑞 + 𝐻2𝑂 ⇌ 𝐻2𝐶𝑂3 (𝑎𝑞) 𝐾𝑒𝑞 = 2 𝑥 10−3

• Very little free H2CO3 in water, so consider

𝐶𝑂2 𝑎𝑞 + 𝐻2𝑂 ⇌ 𝐻+ + 𝐻𝐶𝑂3− 𝐾1 =

𝐻+ 𝐻𝐶𝑂3−

[𝐶𝑂2]= 4.45 𝑥 10−7

as the key equilibrium established initially upon dissolution of CO2 (g)

• The 𝐶𝑂2 ⇋ 𝐻𝐶𝑂3− ⇋ 𝐶𝑂3

2− system in water may be described by the availabilities of the first (see above) and second protons of this weak diprotic acid

𝐻𝐶𝑂3− ⇋ 𝐻+ + 𝐶𝑂3

2− 𝐾2 =𝐻+ 𝐶𝑂3

2−

[𝐻𝐶𝑂3−]

= 4.69 𝑥 10−11

Water Acidity and dCO2

• Predominant species formed by dissolved CO2 depends on pH• The fraction of each species (α) present at a given pH can be calculated readily:

𝛼𝐶𝑂2=

𝐻+ 2

𝐻+ 2 + 𝐾1 𝐻+ + 𝐾1𝐾2=

𝐻2𝐶𝑂3

𝑆𝐶𝑂2

=𝐻2𝐶𝑂3

𝐻2𝐶𝑂3 + 𝐻𝐶𝑂3− + [𝐶𝑂3

2−]

𝛼𝐻𝐶𝑂3− =

𝐾1 𝐻+

𝐻+ 2+𝐾1 𝐻+ +𝐾1𝐾2=

𝐻𝐶𝑂3−

𝑆𝐶𝑂2

=𝐻𝐶𝑂3

−

𝐻2𝐶𝑂3 + 𝐻𝐶𝑂3− +[𝐶𝑂3

2−]

𝛼𝐶𝑂32− =

𝐾1𝐾2

𝐻+ 2+𝐾1 𝐻+ +𝐾1𝐾2=

𝐶𝑂32−

𝑆𝐶𝑂2

=𝐶𝑂3

2−

𝐻2𝐶𝑂3 + 𝐻𝐶𝑂3− +[𝐶𝑂3

2−]

• From these equations, we can derive a distribution of species diagram for the carbonate system as a function of pH.

• The diagram highlights the following:• For pH << pK1, αCO2 is essentially 1

• When pH = pK1, αCO2 = αHCO3

• When pH = ½ (pK1 + pK2), αHCO3 is at its maximum value of 0.98

• When pH = pK2, αHCO3 = αCO3

• For pH >> pK2, αCO3 is essentially 1

• How much does ionic strength impact species distribution?

• *For typical seawater with S=34, T=30 oC and AT=2234 µM

• How does this impact the distribution of species for the carbonic acid system?

Fresh water Seawater*

K1 5.012 x 10-7 1.567 x 10-6

K2 4.786 x 10-11 1.355 x 10-9

0

0.2

0.4

0.6

0.8

1

1.2

3 4 5 6 7 8 9 10 11 12

α

pH

0

0.2

0.4

0.6

0.8

1

1.2

3 4 5 6 7 8 9 10 11 12

α

pH

Freshwater

Seawater

H2CO3 HCO3- CO3

2-

• Knowing the saturation concentration of dCO2 and K1, we can calculate the pH of a pure water sample that has equilibrated with atmospheric CO2 (g).

Recall, 𝐶𝑂2 𝑠𝑎𝑡 =𝑝𝐶𝑂2

𝐾𝐻=

390𝑥10−6𝑎𝑡𝑚

29.1𝑎𝑡𝑚

𝑀1

= 1.34 𝑥 10−5 𝑀

Now, 𝐾1 =𝐻+ 𝐻𝐶𝑂3

−

[𝐶𝑂2], but [H+] ~ [HCO3

-] and so [H+] = 2.38 x 10-6 and pH = 5.62

• What would happen if a strong mineral acid or base were added to this aqueous system?• Recalling that the predominant species in the carbonate system at “near-neutral pH” is the ampholyte

HCO3-, so it can react with both!

𝐻𝐶𝑂3− + 𝐻+ → 𝐻2𝐶𝑂3 → 𝐻2𝑂 + 𝐶𝑂2

𝐻𝐶𝑂3− + 𝑂𝐻− → 𝐻2𝑂 + 𝐶𝑂3

2−

• In either case, the HCO3- serves to “sink” or react away the added acid or base

• Note: This assumes that the HCO3- isn’t consumed completely

• Acidity – capacity of a sample of water to neutralize OH-

• Not usually a concern in environmental systems

• Sources of acidity include,• Acid rain (H2SO4, HNO3)

• Acidic ions in industrial waste (Al3+, Fe3+)

• Natural sources of acids (proteins, fatty acids, CO2)

• Agricultural sources (H2PO4-)

• Alkalinity – Ability of a body of water to neutralize acids• Much more common in environmental systems

• Two common definitions for alkalinity:• Carbonate alkalinity: When carbonate is the only contributor to alkalinity

• Commonly called phenolphthalein alkalinity

• Corresponds to titration with acid to pH 8.3

• Total alkalinity: • Also called acid neutralizing capacity

• Additional contributors to alkalinity present

• Commonly called methyl orange alkalinity

• Corresponds to titration with acid to pH 4.3

𝑑𝐾𝐻 = 𝐴𝑙𝑘𝑎𝑙𝑖𝑛𝑖𝑡𝑦 = 𝐻𝐶𝑂3− + 2 𝐶𝑂3

−2 + 𝑂𝐻− − [𝐻+]

𝐴𝑇 = 𝐻𝐶𝑂3− + 2 𝐶𝑂3

−2 + 𝐻𝑆𝑂4− + 𝐵 𝑂𝐻 4

− + 𝑂𝐻− − 𝐻+ + 𝑒𝑡𝑐 …

• Alkalinity impacts solubility of CO2

• Recall that CO2 acidifies the water, which is more effectively neutralized at higher alkalinity

• Alkalinity impacts solubility of metal salts• Hydrated metal ions are acidic, so greater alkalinity neutralizes/stabilizes them, resulting in

an equilibrium shift to greater dissolution• For example,

• As ability to neutralize H+ increases (due to higher alkalinity), equilibrium shifts to right and more Fe-salt could dissolve to form Fe(H2O)6

2+

𝐹𝑒 𝐻2𝑂 62+ ⇌ 𝐹𝑒 𝑂𝐻 3(𝑠) + 3𝐻+

• Equilibrium between dCO2 and CaCO3 is important in determining several water parameters, including• Alkalinity

• pH

• dCa2+

• The reaction between CaCO3 and dCO2 is

for which we can write the equilibrium expression

Dissolved CO2 and CaCO3 minerals

𝐶𝑎𝐶𝑂3 𝑠 + 𝐶𝑂2 𝑎𝑞 + 𝐻2𝑂 ⇌ 𝐶𝑎2+ 𝑎𝑞 + 2𝐻𝐶𝑂3− (𝑎𝑞)

𝐾′ =𝐶𝑎2+ 𝐻𝐶𝑂3

− 2

[𝐶𝑂2]=

𝐾𝑠𝑝𝐾1

𝐾2= 4.24 𝑥 10−5

• According to our balanced equation, [HCO3-] = 2[Ca2+]

• Plugging in a value of 1.3 x 10-5 M for dCO2 into K’• Assumes 390 ppm atmospheric CO2 and 25 oC water

Yields [Ca2+] = 5.1 x 10-4 M and [HCO3-] = 1.2 x 10-3 M

• Substituting these values into the Ksp expression yields [CO32-] = 8.7 x 10-6 M

• Finally, plugging all the calculated values into the equation

Yields [H+] = 5.54 x 10-9 M and pH = 8.3

Reality check…

𝐾1𝐾2 =𝐻+ 2 𝐶𝑂3

2−

[𝐶𝑂2]= 2.09 𝑥 10−17

• In summary, for a water system exposed to CaCO3 (i.e., limestone) our BOTE estimates yield:• [Ca2+] ~ 10-4 M

• [HCO3-] ~ 10-3 M

• pH ~ 8.3

• All within experimentally measured natural levels

• Consider two simple water samples:• Sample 1: DI water that has been degassed of CO2 and brought to pH 9.0

• Sample 2: DI water made to 0.010 mol L-1 NaHCO3 and pH 8.3

• pH of Sample 1 is higher, so might expect it to have a greater ability to consume acids (i.e., greater alkalinity)

• If we only consider carbonate alkalinity, then

Alkalinity = [OH-] + [HCO3-] + 2[CO3

2-] – [H3O+]

Alkalinity ≠ pH

• Sample 1:[OH-] = 10-5 M; [H+] = 10-9 M; [HCO3

-] = [CO32-] = 0

Alkalinity = 10-5 M = 10 µM

• Sample 2:[OH-] = 10-5.7 M; [H+] = 10-8.3 M; [HCO3

-] = 0.010 M

At a pH of 8.3, HCO3- is the only significant carbonate species, although there is a small

amount (9.4 x 10-5 M) of CO32- . So

Alkalinity = 10-5.7 M + 0.010 M + 2(9.4 x 10-5 M) – 10-8.3

= 0.010 M = 10 000 µM

• Conclusion: Sample 2, even though at a lower pH, has a greater capacity to neutralize acids

• Alkalinity is often expressed in terms of ppm CaCO3

• But, remember that each CO32- can neutralize 2 H+, so

µM H+ acceptors ppm CaCO3 ppm Ca2+ Sensitivity to acids

<200 <10 <4 High

200 – 400 10 – 20 4 – 8 Moderate

>400 >20 >8 Low

1 𝑚𝑔 𝐶𝑎𝐶𝑂3

𝐿

1 𝑚𝑜𝑙 𝐶𝑎𝐶𝑂3

100 𝑔= 10 𝜇𝑀 𝐶𝑂3

2−

• Alkalinity: Ability of a body of water to neutralize acids

• Addition of strong acid shifts carbonate equilibrium toward carbonic acid (i.e., increased acidity)• Reduces ability of water to neutralize further additions of acid

• Acid rain can be neutralized by soil• Extent of neutralization depends strongly on type of soil/bedrock

• Limestone (CaCO3) efficiently neutralizes acid rain

• Granite does not efficiently neutralize acid rain

Surface Water Acidification

𝐴𝑙𝑘𝑎𝑙𝑖𝑛𝑖𝑡𝑦 = 𝐻𝐶𝑂3− + 2 𝐶𝑂3

−2 + 𝑂𝐻− − [𝐻+]

• Consider Lake Champlain• Average total alkalinity: 50 ppm CaCO3

• Average depth: 20 m

• Mean annual rainfall: 935 mm

• Average rain pH: 4.9

• Is Lake Champlain sensitive to acid rain effects?

• Two questions we need to ask are:• (1) How much acid can the lake neutralize?

• (2) How much acid falls into the lake?

• If (2) > (1), the lake will acidify…

Importance of Alkalinity

• Per km3, the moles of “ANC” are:

(𝑉𝑐𝑜𝑙𝑢𝑚𝑛) 𝐴𝑙𝑘𝑎𝑙𝑖𝑛𝑖𝑡𝑦 = 1𝑘𝑚2 0.020𝑘𝑚50 𝑚𝑔 𝐶𝑎𝐶𝑂3

𝐿

0.020 𝑥 109𝑚3 1000𝐿

𝑚3

50 𝑚𝑔 𝐶𝑎𝐶𝑂3

𝐿= 1 𝑥 109 𝑔 𝐶𝑎𝐶𝑂3,

Which corresponds to 1 x 107 moles CO32-.

Remembering that each mole of CO32- neutralizes 2 moles of acid, this results in the ability of

the lake to consume 2 x 109 moles H+ per km3.

• The amount of rain falling on 1km2 is 106 𝑚2 0.975 𝑚 = 0.975 𝑥 106 𝑚3 = 0.975 𝑥 109 𝐿

Which corresponds to

0.975 𝑥 109𝐿 10−4.9 = 1.2 𝑥 104 𝑚𝑜𝑙 𝐻+

So, assuming no other inputs/outputs to/from lake, it would become acidified in 2 𝑥 107

1.2 𝑥 104 ~1670 𝑦𝑒𝑎𝑟𝑠. (Phew!!!)

• Areas most affected by acid rain are those having granite or quartz bedrock

If pH or rain had remained at 1980’s level of ~4.0, the time estimate to acidification of Lake Champlain would have been ~10x faster!

• Oxidation – reduction • dO2 is most important oxidizing agent in surface waters

• Concentration depends also on “thermal pollution”

• Most common substance oxidized by dO2 is organic matter

• Representing organic matter as polymerized carbohydrate (i.e., plant fiber), CH2O – this is BOD

• Dissolved NH3, NO2- and NH4

+ (result of biological activity) are also oxidized by dO2 – this is COD• Eventually leads to NO3

-

Chemistry of Natural Waters

𝐶𝐻2𝑂 𝑎𝑞 + 𝑂2 𝑎𝑞 → 𝐶𝑂2 𝑔 + 𝐻2𝑂

• Can also get chemistry that does not occur readily in aerobic environment

𝐹𝑒3+ 𝑖𝑛𝑠𝑜𝑙𝑢𝑏𝑙𝑒 + 𝑒− → 𝐹𝑒2+ 𝑠𝑜𝑙𝑢𝑏𝑙𝑒• Often can have aerobic (epilimnion) and anaerobic (hypolimnion) conditions in same body of

water

• Anaerobic conditions in lakes typically do not last indefinitely

• Inversion of water layers during winter/summer cycles

Species typically in their most oxidized forms

Species typically in their most common reduced forms

• Used to quantify the chemical reducing or oxidizing nature of natural waters• pE – negative base 10 log of the effective “concentration (i.e., activity) of electrons in water”

𝑝𝐸 = − log 𝑎𝑒

• Large negative pE indicates large value for electron activity in solution → reducing conditions• This is the situation in anoxic water, such as a swamp

• Large positive pE indicates small value of electron activity in solution → oxidizing conditions• This is the situation in well aerated surface water

• In practice, pE ranges from about -12 to 25• Boundaries are imposed by water stability to redox

Electron Activity Measured as pE

• Let’s take the example of iron redox in water:

𝐹𝑒3+ 𝑎𝑞 + 𝑒− ⇌ 𝐹𝑒2+ 𝑎𝑞

[Note: we will write most ½ reactions as reductions, regardless of spontaneous direction of the reaction]

𝐾𝑒𝑞 =𝑎𝐹𝑒

2+

𝑎𝐹𝑒3+ 𝑎𝑒−

from which we can solve1

𝑎𝑒−=

𝐾𝑒𝑞 𝑥 𝑎𝐹𝑒3+

𝑎𝐹𝑒2+

and

𝑝𝐸 = − log 𝑎𝑒− = log 𝐾𝑒𝑞 + log𝑎𝐹𝑒

3+

𝑎𝐹𝑒2+

Also, we know that Δ𝐺𝑜 = −2.303𝑅𝑇 log 𝐾𝑒𝑞 = −𝑛𝐹𝐸𝑜

Calculation of pE

Combining the last two equations, we arrive at

𝑝𝐸 =𝑛𝐹𝐸𝑜

2.303𝑅𝑇+ log

𝑎𝐹𝑒3+

𝑎𝐹𝑒2+

Substituting in values for “F” and “RT” and for 1 mole of electrons transferred gives

𝑝𝐸 =𝐸𝑜

0.0591+ log

𝑎𝐹𝑒3+

𝑎𝐹𝑒2+

Under standard conditions (i.e., “o”), 𝑎𝐹𝑒3+ = 𝑎𝐹𝑒

2+ = 1 and

𝑝𝐸 = 𝑝𝐸𝑜 =𝐸𝑜

0.0591

From a table of standard reduction potentials, we find that

𝐹𝑒3+ + 𝑒− ⇌ 𝐹𝑒2+ 𝐸0 = 0.771 𝑉

And

𝑝𝐸0 = 13.0

• From standard reduction potentials:

𝐹𝑒3+ 𝑎𝑞 + 𝑒− ⇌ 𝐹𝑒2+ 𝑎𝑞 𝐸𝑜 = +0.771 𝑉

𝑝𝐸 =0.771𝑉

0.0591𝑉+ log

𝑎𝐹𝑒3+

𝑎𝐹𝑒2+ = 13.0 + log(𝑄)

• If standard reduction potentials are not available:• We can use the relation

log 𝐾𝑒𝑞 =𝑛𝐸𝑜

0.0591= 𝑛𝑝𝐸𝑜

or 𝑝𝐸𝑜 =log 𝐾𝑒𝑞

𝑛

Methods for Calculating pE

• For example when several reactions combine to produce an overall half-reaction

𝐹𝑒 𝑂𝐻 3 𝑠 + 3𝐻3𝑂+ 𝑎𝑞 + 𝑒− ⇌ 𝐹𝑒2+ 𝑎𝑞 + 6𝐻2𝑂

This reaction is the sum of

𝐹𝑒 𝑂𝐻 3 𝑠 ⇌ 𝐹𝑒3+ 𝑎𝑞 + 3𝑂𝐻− 𝑎𝑞 dissolution

𝐹𝑒3+(𝑎𝑞) + 𝑒− ⇌ 𝐹𝑒2+(𝑎𝑞) reduction

3𝐻3𝑂+ 𝑎𝑞 + 3𝑂𝐻− ⇌ 6𝐻2𝑂 acid-base

• We can find (from a table) that Ksp(Fe(OH)3) = 9.1 x 10-39 and so log Ksp = -38

• Also, we know that for the standard reduction of Fe3+ to Fe2+ occurs at a potential of +0.771V, so pEo = log Kred = 13.0

• Finally, for the acid base, we have 𝐾𝑎𝑏 =1

𝐾𝑤3 = 1042 and log Kab = 42.

• Now, multiplying all the Ks (or adding all the log Ks) yields

pErxno =

log 𝐾𝑟𝑥𝑛

1= -38 + 13.0 + 42.0 = +17

• Most generally applicable is use of standard thermodynamic properties (ΔGo)

• Use Hess’s Law to calculate ΔGrxno for the overall reaction

𝑝𝐸𝑜 =−Δ𝐺𝑜

2.303𝑛𝑅𝑇

For example,

𝑆𝑂42− 𝑎𝑞 + 10𝐻3𝑂+ 𝑎𝑞 + 8𝑒− ⇌ 𝐻2𝑆 𝑎𝑞 + 14𝐻2𝑂

From tables of ΔGo, we calculate ΔGrxno = -231.98 kJ

𝑝𝐸𝑜 =− −231.98 𝑘𝐽

1000𝐽𝑘𝐽

(2.303)(8.314 𝐽 𝑚𝑜𝑙−1𝐾−1)(298.2𝐾)(8 𝑚𝑜𝑙 𝑒−)= 5.08

• Three ways to calculate pE based on available data1. From standard reduction potentials

𝑝𝐸 =𝐸𝑜

0.0591 𝑛−

1

𝑛log 𝑄 (if equation written as a

reduction)

2. From overall equilibrium constant for reaction

𝑝𝐸𝑜 =log 𝐾𝑒𝑞

𝑛

3. From Gibbs Free Energy change of reaction (ΔGo)

𝑝𝐸𝑜 =−Δ𝐺𝑜

2.303𝑛𝑅𝑇

Recall:

• A local industry is dumping wastewater that contains 26 ppm Cr(III) into a local stream• Knowing that it is the Cr(VI) species that is most toxic to humans, is this a serious problem?

• Make two assumptions about the stream: pH = 6.5 and stream is well aerated (i.e., saturated with dO2)

• For the chromium system, the ½-reaction is

𝐶𝑟2𝑂7−2 𝑎𝑞 + 14𝐻3𝑂+ 𝑎𝑞 + 6𝑒− ⇌ 2𝐶𝑟3+ 𝑎𝑞 + 21𝐻2𝑂

Example: Cr(III)/Cr(VI)

𝑝𝐸"𝑠𝑦𝑠𝑡𝑒𝑚" = 𝑝𝐸𝑜 −1

6log

𝐶𝑟3+ 2

[𝐶𝑟2𝑂72−] 𝐻3𝑂+ 14

= 23.0 −1

6𝑙𝑜𝑔

1

10−6.5 14 −1

6𝑙𝑜𝑔

𝐶𝑟3+ 2

[𝐶𝑟2𝑂72−]

• But, the pE”system” is established by O2 oxidation in this well aerated environment, soWe will develop the pE expression for this scenario soon:

𝑝𝐸"𝑠𝑦𝑠𝑡𝑒𝑚" = 𝑝𝐸𝑂2

𝑜 −1

𝑛log

1

𝑝𝑂2 𝐻3𝑂+ 4 = 20.8 +1

4𝑙𝑜𝑔 0.21 10−6.5 4 = 14.1.

And now

• Solving for 𝐶𝑟3+ 2

𝐶𝑟2𝑂72− yields 1.7 x 10-38!!!

• ALL the Cr(III) has been oxidized to Cr(VI) and we have a real problem!

14.1= 23.0 −1

6𝑙𝑜𝑔

1

10−6.5 14 −1

6𝑙𝑜𝑔

𝐶𝑟3+ 2

𝐶𝑟2𝑂72− = 7.8 −

1

6𝑙𝑜𝑔

𝐶𝑟3+ 2

𝐶𝑟2𝑂72−

• Reducing conditions – low pE

2𝐻2𝑂 + 2𝑒− ⇌ 𝐻2 𝑔 + 2𝑂𝐻−(𝑎𝑞)• We can write a modified Nernst equation for this ½ reaction:

• From table of standard reduction potentials, Eo for this ½-reaction is -0.828 V, so 𝑝𝐸𝑜 =

−0.828𝑉

0.0591= −14.0

Setting the boundary condition pH2 = po= 1 atm, gives

Which defines the lower boundary line for water stability

Water Stability Boundaries

𝐸 = 𝐸𝑜 −1

𝑛𝑙𝑜𝑔

𝑝𝐻2

𝑝𝑜 𝑂𝐻− 2

𝐸 = 𝐸𝑜 − 𝑙𝑜𝑔 𝑂𝐻−

𝑝𝐸 = 𝑝𝐸𝑜 + 𝑝𝑂𝐻 = −14 + 𝑝𝑂𝐻 = −𝑝𝐻

• Oxidizing conditions – high pE

6𝐻2𝑂 ⇌ 4𝐻3𝑂+ 𝑎𝑞 + 𝑂2 𝑔 + 4𝑒−

• We can write a modified Nernst equation for this ½ reaction:

• From table of standard reduction potentials, Eo for this ½-reaction is 1.229 V, so 𝑝𝐸𝑜 =1.229𝑉

0.0591= 20.80

Setting the boundary condition pO2=po=1atm, gives

Which defines the upper boundary line for water stability

𝐸 = 𝐸𝑜 +1

𝑛𝑙𝑜𝑔

𝑝𝑂2

𝑝𝑜 𝐻3𝑂+ 4

𝐸 = 20.80 + 𝑙𝑜𝑔 𝐻+

𝑝𝐸 = 20.80 − 𝑝𝐻

Because we’ve written the ½-reaction as an oxidation

Remember, defined for the reduction ½-reaction

• Template of water bodies to use with Pourbaix diagrams

Pourbaix Diagram for Water

• Provide a visual representation of the zones of dominance for the various oxidation states of an element in water• Situation is somewhat more

complex than indicated above

because we also have to take

into account pH effects

• Solid lines indicate combination

of pH and pE where the concentrations

of the species on either side are equal

Pourbaix (pE-pH) Diagrams

𝐹𝑒2+ 𝑎𝑞 + 2𝑂𝐻− ⇌ 𝐹𝑒 𝑂𝐻 2 (𝑠)

𝐹𝑒3+ 𝑎𝑞 + 3𝑂𝐻− ⇌ 𝐹𝑒 𝑂𝐻 3 (𝑠)

• Nitrogen is an essential nutrient in natural waters

• Organic forms of nitrogen are of concern to human health

• In solution, the most important intermediates between the extremes are nitrite ion and molecular nitrogen

pE-pH Diagram for Nitrogen

Oxid.Num.

-3 0 +1 +2 +3 +4 +5

(aq) NH4+

NH3

N2 NO2- NO3

-

gas NH3 N2 N2O NO NO2

Increasing levels of Nitrogen oxidation

Application of Pourbaix Diagrams

Acid mine water

Rain, lakes

RiversOceans

Submerged soils, swamps, etc

Marine sediments

• For example, let’s consider the reduction of NO3- to NH4

+, which proceeds according to the reaction

1

8𝑁𝑂3

− +5

4𝐻+ + 𝑒− ⇌

1

8𝑁𝐻4

+ +3

8𝐻2𝑂

For this reaction, we can use standard tables of reduction potentials to calculate the standard reduction potential for the overall reaction:

(Eo) = +0.881 V, yielding a pEo = 14.90, indicating a mildly oxidizing environment

• When not under standard conditions (i.e., concentrations/activities different from unity),

𝑝𝐸 = 𝑝𝐸0 − 𝑙𝑜𝑔𝑁𝐻4

+

𝑁𝑂3−

18 𝐻+

54

= 14.90 −5

4𝑝𝐻 −

1

8𝑙𝑜𝑔

[𝑁𝐻4+]

[𝑁𝑂3−]

• Oceans:• pE ~10 and pH ~ 8, so

10 = 14.90 −5

4(8) −

1

8𝑙𝑜𝑔

[𝑁𝐻4+]

[𝑁𝑂3−]

[𝑁𝐻4+]

[𝑁𝑂3−]

~10−41

• Marine sediment:• pE ~ -4 and pH ~ 8

[𝑁𝐻4+]

[𝑁𝑂3−]

~108

• When dO2 is high, the reduction of O2 to water is the dominant equilibrium determining overall electron availability (high pE)

1

4𝑂2 + 𝐻+ + 𝑒− ⇌

1

2𝐻2𝑂

• In these cases, pE of the water is related to the acidity and to the pO2:

𝑝𝐸 = 20.80 +1

4𝑙𝑜𝑔 𝐻+ 𝑝𝑂2 = 20.80 − 𝑝𝐻 +

1

4log(𝑝𝑂2)

• For a neutral sample of water (i.e., pH 7) that is saturated by O2 from air (i.e., pO2=0.21 atm), pE = 13.6

• pE < 13.6 is reducing

• In a practical application, the pE can be calculated from:

• The electrode potential measured for whatever process determines the electron availability in the water sample

• The concentrations and/or pressures of the dominant redox pairs

𝑝𝐸𝑜 =𝐸𝑜

0.0591(look familiar)

• Note: pE is defined “per electron”, so balanced reaction coefficients may appear “weird”

• The environmental behavior of an element or compound depends on the particular form of the species that is present. In the aqueous environment, the species distribution depends on a number of factors, including solution pH, pE, and the nature and availability of complexing ligands.

• In order to determine the species distribution of a particular substance, it is usual to make an assumption that all forms are in equilibrium with the surroundings. Furthermore, it is frequently assumed that concentration = activity. Accurate calculations require activity estimations based on detailed knowledge of the solution composition and ionic strength.

• Single variable distribution diagrams show how the concentrations of different species change as a function of the one defined variable.

• Two variable diagrams indicate regions on a two-dimensional plot where individual species predominate. No detailed information about concentrations within the domains is given.

• The pH and pE of the aqueous environment are two key properties that define the nature of chemical species. The pE is a measure of redox status and is high when the water is well aerated and low where oxygen is excluded or has been consumed.

In summary…

• Used to quantify the chemical reducing or oxidizing nature of natural waters• pE – negative base 10 log of the effective “concentration (i.e., activity) of electrons in water”

• Low pE – reducing environment

• High pE – oxidizing environment

• pE is governed by the dominant redox equilibrium reaction in the water sample• Much like the strongest acid dictates the pH of a solution

The pE Scale

• Gross classification as “dissolved” or “particulate”

• Found in every body of water on Earth

• Typical concentrations of few ppmC

• Only small fraction of all aquatic carbon is OM

• However, its influence in environmental chemistry is far beyond its mass contribution

• To put this in perspective, sedimentary carbonate material (i.e., minerals, soil, etc) ~ 20 000 000 Pg!

Chemistry of the Hydrosphere: Organic Matter (OM)

Mass of C (Pg)

Living organisms 3

Dissolved Organic Carbon (DOC) 1000

Dissolved Inorganic Carbon (DIC) 37000

• Natural• Terrestrial- and hydrosphere-derived – primarily from plant and/or microbial residues

• Terrestrial-derived can be transported to aquatic systems (eg, runoff)

• Anthropogenic• Human inputs, such as sewage, bulk effluents, agricultural chemicals, medicinals, etc.

• Delineation of the two is not always obvious:• Chloroform is a known anthropogenic pollutant

• Chloroform (as well as >1500 other organochlorines) is also formed through natural processes by living organisms

Sources of Organic Matter

• Small molecules• For example, monosaccherides or low MW organic acids

• Chemical structure and properties are amenable to individual specific study• Individual polluting species, like chemical pesticides, fall into this category

• Macromolecules• For example, humic and fulvic acids

• Treated as classes in terms of their general structural properties and reactivity

• Characterization based on operational definition (i.e., particular analytical protocol) and noton fundamental structural properties

• Many natural forms of OM fall in the macromolecule category

Categories of Organic Matter

• Toxicity of specific organic compounds (later)

• Reaction with other aquatic species• For example, mono- and di-methyl tin are formed inside the fish gut upon intake of inorganic

tin pollutant

• Toxicity can be enhanced (as is the case for Sn) or reduced in the presence of OM

• Consumption of oxygen (BOD)

Issues Related to aqueous OM

• Makes up about 50% of DOM in surface water

• Chemical composition not discrete at molecular level

• Subdivided into three operational classes• Fulvic acid (FA) – fraction of HM that is soluble in aqueous solutions spanning all pH values

• Humic acid (HA) – insoluble in aqueous solutions at pH <2

• Humin (Hu) – insoluble at all pH values

• Formation of HM is not well understood• Degradative pathway

• labile macromolecules, such as carbohydrates and proteins

• Synthetic pathway

Humic Material

• Formation of HM is not well understood• Degradative pathway

• Labile macromolecules, such as carbohydrates and proteins, are lost (i.e., liberated) during microbial attack

• Refractory compounds/biopolymers, such as lignin, melanins and cutin, are selectively transformed to high MW material that is precursor to Hu

• Further oxidation of these materials produces molecules small enough and hydrophilic enough to be soluble

• Synthetic pathway

• Both pathways are likely important, depending on environmental conditions

Formation of Humic Material

Plant material

HuminHumic

AcidFulvic Acid

Small molecule

Plant material

Small molecule

Fulvic AcidHumic

AcidHumin

C O H N ash

45 – 60 25 – 45 4 – 7 2 – 5 0.5 – 5

HM Composition and Structure Elemental analysis of HM (mass %)

Carbon content usually increases in the series

Fulvic acid < Humic acid < Humin

Oxygen content usually follows the reverse trend

Carbon content usually greater in soil humus than in humate from lakes/oceans

Reverse is true for O and N

HM can vary in age depending on environment

20 years for streams (highly aerated)

500 – 1000 years in soil (less aerated)

>1000 years for buried deposits (eg, peat/coal) in anaerobic environment

• Standard reduction potentials for reactions important in HM oxidation in the environment

• Note: The Eo values reported are for the reactants in their standard state. So,𝑂2 𝑔 + 4𝐻+(𝑎𝑞) ⇌ 2𝐻2𝑂(𝑙)

• So, when considering E and pE of reaction, we will use concentrations (M) and partial pressure (p)

Reducing Half-Reaction Eo (V)

𝑂2 + 4𝐻+ + 4𝑒− ⇌ 2𝐻2𝑂 1.229

2𝑁𝑂3− + 12𝐻+ + 10𝑒− ⇌ 𝑁2 + 6𝐻2𝑂 1.25

𝑀𝑛𝑂2 + 4𝐻+ + 2𝑒− ⇌ 𝑀𝑛2+ + 2𝐻2𝑂 1.23

𝐹𝑒 𝑂𝐻 3 + 3𝐻+ + 𝑒− ⇌ 𝐹𝑒2+ + 3𝐻2𝑂 1.005

𝑆𝑂42− + 9𝐻+ + 8𝑒− ⇌ 𝐻𝑆− + 4𝐻2𝑂 0.248

𝐶𝑂2 + 8𝐻+ + 8𝑒− ⇌ 𝐶𝐻4 + 2𝐻2𝑂 0.17

• Ecological redox sequence

Stigliani, 1988. Changes in valued “Capacities” of soils and sediments as indicators of nonlinear and time-delayed environmental effects, Fig. 16

Aer

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esp

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ion

Den

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fica

tio

n

Mn

red

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Fe

red

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SO

42

-re

du

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n

Met

han

og

enes

is

• HM as a proton acceptor• Acidic character associated largely with –COOH and PhOH groups

• pKa of –ROOH is about 2.5 – 5, so essentially deprotonated in most water bodies, giving HM a negative charge• Major source of anionic charge in the dissolved phase

• Provides charge balance for high concentrations of cations typically found in water

• HM reaction with small organic molecules• Reaction forces include

• Van der Waals forces (always present)

• Electrostatic forces

• Hydrogen bonding

• Hydrophobic interactions

• Salt linkage or ligand exchange

Forms and Reactivity of HM

• Electrostatic forces dominate when have positively charged (or highly polar) small molecules

• Hydrogen bonding possible when small molecule has electronegative atoms like N or O

• Complexation of metal ions• Common functional groups found in HM include:

• All are effective chelators of metal ions

• For example,

• Strength of interaction between HM and metal ion depends on:• Nature of the metal ion – higher charge → stronger bond

• pH of ambient solution – higher pH → stronger bond• More anionic sites to interact with metal ion

• Ionic strength of solution – higher ionic strength →weaker bond• More “other ions” to complex with HM

• More anions to react with metal cation

• Availability of functional groups – depends on nature and concentration of HM

Complexation Strength

• Typical concentrations (i.e., complexation capacity) for some types of water are:• Rivers: 1 – 2 µM

• Lakes: 2 – 5 µM

• Ponds: 5 – 15 µM

• Swamps: >15 µM

• Stability constants: 𝐾𝑓′ =

𝑀𝑒𝑡𝑎𝑙−𝐹𝐴

𝑀𝑒𝑡𝑎𝑙𝑓𝑟𝑒𝑒 [𝐹𝐴𝑓𝑟𝑒𝑒]

• FAfree = concentration of available complexing functional groups on HM

• Because need deprotonated function groups to complex with metal ions, this value depends on pH

Mg2+ Ca2+ Mn2+ Co2+ Ni2+ Cu2+ Zn2+ Pb2+

Kf’ (a) 1.4x102 1.2x103 5.0x103 1.4x104 1.6x102 1.0x102 4.0x103 1.1x104

(a) Conditional stability constants determined at pH 5.0

• Most common organic complexing agents are detergents• Ethylenediaminetetraacetic acid (EDTA) – used for industrial cleaning and in textile and

paper industries

• Nitrilotriacetic acid (NTA) – phosphate substitute in detergents

• Both have very strong stability constants with metal ions• Can either precipitate the ions as complexation salts or enhance solubility (and corresponding

mobility) as complexed metallic ions

Metal Complexation with Anthropogenic Organic Matter

pH, Mn+

𝑝𝐾𝑎1 = 2.0

𝑝𝐾𝑎2 = 2.7

𝑝𝐾𝑎3= 6.2

𝑝𝐾𝑎4 = 10.3

• The effective formation constant (Kf’)

depends on pH because chelatorsmust be deprotonated to be effective.• For example,

For EDTA,

EDTA Chelation

(11.4)(9.5)

Numbers in parentheses indicate formation constants with NTA

𝐾𝑓′ = 𝐾𝑓 𝛼4

• NTA is an interesting case study• Present in detergents at concentrations of ~15% by mass

• Acts as an efficient binder for Ca2+ and other ions, softening the water and improving the action of the surfactants in the detergent

• Triprotic carboxylic acid: pKa1=1.66, pKa2=2.95, pKa3=10.28

• Forms 1:1 complex with most metal ions under appropriate pH conditions• Note that in natural water systems, NTA exists mainly as the di-anion

NTA Chelation

𝑝𝐾𝑎1 = 1.66

𝑝𝐾𝑎2= 2.95

𝑝𝐾𝑎4 = 10.28

Metal Ion

Log Kf

(pH 11)Log Kf’ (pH 8)

Mg2+ 5.43 3.14

Ca2+ 6.45 4.17

Mn2+ 7.44 5.16

Fe2+ 8.83 6.55

Zn2+ 10.66 8.38

Pb2+ 11.4 9.18

Cu2+ 12.96 10.68

Fe3+ 16.26 13.98

Very stable complexes formed with most metal ions

Free NTA in wastewater is degraded by biological processes during water treatment

Complexed NTA is resistant to these same degradation processes

Complexation increases the metal lifetime in solution (i.e., mobilization lifetime)

Typical levels just below waste streams reach concentrations of several hundred ppb NTA

Conditional formation constant: 𝐾𝑓′ = 𝛼3𝐾𝑓 =

𝐾𝑎1𝐾𝑎2𝐾𝑎3

𝐻+ 3+ 𝐻+ 2𝐾𝑎1+ 𝐻+ 𝐾𝑎1𝐾𝑎2+𝐾𝑎1𝐾𝑎2𝐾𝑎3

• Water – aquo complexes• Simplest form in which metal ion can exist in water

• Often consider metal ions coordinated with water as “free ions”

• Strength of aquo complex is inversely related to charge density on the ion (C pm-1)• Stronger aquo complex means lower chemical activity of the ion

• Metal-aquo complexes have acid/base character

𝑀 𝐻2𝑂 𝑎𝑏+ + 𝐻2𝑂 ⇌ 𝑀 𝐻2𝑂 𝑎−1 𝑂𝐻 𝑏−1 + + 𝐻3𝑂+

often written simply as

𝑀𝑏+ + 𝐻2𝑂 ⇌ 𝑀𝑂𝐻 𝑏−1 + + 𝐻3𝑂+

Other Complexing Agents

• General acidity trends for metal-aquo complexes include:• Higher charge – greater acidity

• Transition metals more acidic than alkali/alkaline earth metals

• Smaller, more highly charged ions yield higher acidity aquocomplexes

• Other inorganic complexing agents include:• Chloride

• Ammonia

• Phosphate

• Cyanide

• Sulfide, sulfite, sulfate

Metal Ion

pKa

Na+ 14.48

Mg2+ 11.42

Ca2+ 12.7

Al3+ 5.14

Fe2+ 10.1

Fe3+ 2.19

• Traditional classification of metals (and semi-metals)• Type A metal ions (“hard sphere metals”)

• Stability of complexed ion is positively correlated with 𝑍2

𝑟for both the metal ion and the

ligand• Complex stability then follows trend Mg2+ > Ca2+ > Sr2+ > Ba2+

• A preference for O- and F-containing ligands

• For example, Al(H2O)5(OH)2+ is an important soluble aluminum species

• No sulfides formed

• Poor affinity for ammonia and cyanide

• Form sparingly soluble precipitates with OH-, CO32- and PO4

3-

• For example, Ksp (CaCO3) ~ 10-9, Ca(OH2) ~ 10-6 and Ca3(PO4)2 ~ 10-33

Metals in the Hydrosphere

• Type B metal ions (“soft sphere metals”)• Typically transition metals (with residual d-electrons)

• More polarizable than class A ions

• “softness” correlated with Χ2𝑟 (where Χ = electronegativity)

• A preference for I-, N- and S-containing ligands

• Complex stability follows trend

Mn2+ < Fe2+ < Co2+ < Ni2+ < Cu2+ > Zn2+

Remember, generally• Χ - increases to right and top of periodic table.• r - decreases to right of periodic table• r - increases down periodic table

Classification of Metals

[Stumm and Morgan, Aquatic Chemistry, John Wiley and Sons, 1996, Fig. 6.12]

Ionic bonds

Covalent bonds

• Of special environmental concern• Great diversity of compounds possible (see As example below)

• Increased compound (i.e., metal) volatility

• Enhanced bioaccumulation

• Enhanced toxicity

Organometallic (Metalloidal) Compounds

• Owing to their dramatic toxicity increase upon chemical transformation in the environment, spotlight is on

Hg, Pb, Cd and As

Heavy Metals of “Special” Concern

• Global emission estimates (Mg y-1)

Mercury in the Environment

1990 2000 2010

Hg0 2,150 1,540 1,480

ASGM 278 366 378

Products 829 576 398

Combustion 1040 600 699

Hg (II) 739 617 807

Total2,890 (2,120–

4,180)2,160 (1,490–

3,470)2,280 (1,520–

3,730)

410 𝑚𝑖21600𝑚

𝑚𝑖

2

= 1.25 𝑥 109𝑚2

1.25 𝑥 109𝑚210−5𝑔 𝐻𝑔(𝐼𝐼)

𝑚2 𝑦𝑟=

= 12.5 𝑘𝑔𝐻𝑔 𝐼𝐼

𝑦𝑟‼!

In Lake Champlain:

• Hg(0) is volatile (Gaseous elemental Mercury, GEM)• At room temp and pressure, pHg ~ 8 Torr!

• Compare to pH2O ~ 55 Torr

• Because it is neutral, it can cross the blood-brain barrier• In contrast to other heavy metals that are toxic primarily as ions

• Hg readily forms amalgams• Liquid mixtures of Hg with other metals (usually, to soften them)

• Dental fillings still predominant (non-energy) use

• Dentists collectively release about the same amount of Hg metal as emitted by coal-fired power plants!

• Gold and silver production

• 1 g Hg(0) lost to environment per g of Ag produced

• 15 g Hg(0) lost to environment per g of Au produced!

• People living in mining regions subjected to >50 ug m-3 of Hg(0) vapor

• More that 50x the WHO public exposure guideline

• Hg(II) is predominant environmental form• Present as insoluble salts (such as HgS)

• Modern levels (even in remote areas) are 2-5x greater than pre-industrial

• In natural waters, Hg(II) is attached to suspended particulates• Eventually deposits into sediments (maybe more later)

• Major sources include• Waste incineration

• Coal and fuel oil combustion

• Associated predominantly with atmospheric particulate matter• Particulate mercury (TPM)

• Oxidation of vaporized mercury results in rain-out of the pollutant• Produces a regional problem rather than just a local one

• Reactive gaseous mercury• Likely gaseous mercuric chloride (HgCl2)

𝐻𝑔2+ 𝑐𝑜𝑎𝑙flame

𝐻𝑔𝑜 𝑔𝑎𝑠𝐶𝑙𝑥

𝐻𝑔𝐶𝑙2 (𝑔)• Reaction may be heterogeneous, likely (?) catalyzed by atmospheric aerosol

• Oxidized fraction depends upon Cl content of the coal combusted

• Lifetime depends on Hg speciation

• TPM and RGM can be scrubbed efficiently to reduce emissions• Virtually impossible to keep most of the GEM from being emitted

Form of Hg Chemical formula τatmos

GEM Hg(0) (g) Months - year

TPM Hg2+ (adsorbed) Few weeks

RGM HgCl2 (g) Days - week

• Methylmercury, (CH3)nHgXn-1 – special case• Covalent compounds (like HgCl2), not ionic solids

• Formed under anaerobic conditions (eg, muddy sediments of rivers and lakes)

• Biomethylation occurs in anaerobic bacteria and microorganisms vis methylcobalamin

• Dimethyl mercury, (CH3)2Hg• Volatile molecular liquid

• Evaporates from water relatively quickly

• Mixed methyl mercury compounds, CH3HgX, • X = Cl or OH typically

• Even more readily formed that dimethyl mercury (at the surface of sediments)

• Less volatile than dimethyl mercury

• Methylmercury production predominates over dimethylmercury formation in acidic or neutral aqueous systems

• Methylation of inorganic mercury occurs especially near the interface of the epilimnion and the hypolimnion

• Methylation also occurs at the interface of the hypolimnion with sediments (but not in aerobic water)

Production and Fate of Dimethyl Mercury

Special concern for organic mercury in the environment?

Hg2+ is not readily transported across biological membranes

Organo-Hg is soluble in fatty tissue and accumulates there

Binds with S-containing amino acids, allowing it to cross blood-brain barrier

Organo-Hg half-life inhuman body ~ 70 days (so, has time to accumulate!)

Organo-Hg is more hazardous than even Hg(0) vapor!

Lead in the Environment

• Effects of Pb poisoning known to ancient Greeks!

• Established limits often exceeded in modern times, due to:• Lead solder used in joints of domestic copper pipes

• Lead pipes used historically

• Leaching of Pb is exacerbated by• Slightly acidic water

• Softened water (i.e., lowered carbonate levels)

Lead in Drinking Water

Lead in Drinking-Water Systems

• Maximum allowed levels of Pb in drinking water is 10 (US) to 15 (Canadian) ppb

• Pb used in solder at joints of Cu water pipes, and Pb used in previous decades to

construct pipes, can dissolve into water during transport to point of consumption,

especially in soft or acidic waters

• Pb2+ in insoluble carbonates deposited as scale can be oxidized further to Pb4+ by

action of residual hypochlorous acid (HOCl) from disinfection process in water

treatment plants:

PbCO3(s) + HOCl + H2O ➔ PbO2(s) + HCl + H2CO3

• If change is made to a less powerful oxidizing agent (e.g., chloramines) to purify water,

Pb4+ in some of the PbO2 present in the scale of pipes can be reduced to Pb2+ and

contaminate the purified drinking water (this happened in Greenville, NC, in 2005):

PbO2(s) + 4 H+ + 2 e-➔ Pb2+ + 2H2O

• Pb2+ can also be released at junction of Pb and Cu piping by oxidation of Pb by Cu:

Pb(s) + Cu2+ ➔ Pb2+ + Cu(s); Cu2+ formed by Cu(s) reacting with Cl or NH3 in H2O

• Lead forms insoluble carbonate salt (Ksp = 1.5 x 10-13)• Lower pH limit for drinking water is 6.5

• At this pH, solubility of PbCO3(s) is

𝐾𝑠𝑝 = 𝑃𝑏2+ 𝐶𝑂32− = 𝑃𝑏2+ 𝛼2 𝐶𝑂2 𝑑𝑖𝑠𝑠 = 1.5 𝑥 10−13

But, CO2 also comes from dissolved PbCO3, so

𝐶𝑂3 𝑡𝑜𝑡𝑎𝑙 = 𝐶𝑂32−

𝑃𝑏𝐶𝑂3+ 𝐶𝑂3 𝑎𝑡𝑚𝑜𝑠 = 𝐶𝑂3

2−𝑃𝑏𝐶𝑂3

+ 𝛼2 𝐶𝑂2 𝑑𝑖𝑠𝑠𝑜𝑙𝑣𝑒𝑑

We can then recast the solubility equation as

1.5 𝑥 10−13 = [𝑃𝑏2+] 𝑃𝑏2+ + 𝛼2𝑑𝐶𝑂2

Solving for α2 at pH 6.5, yields 8.8 x 10-5 and dCO2 = 1.4 x 10-5 M for a ξCO2 = 400 ppm. Plugging all values in and solving using quadratic equation yields [Pb2+] = 3.9 x 10-7 M or 79 µg/L (i.e., ppb)

[Note: This is a case where dissolution of PbCO3 and atmospheric CO2 both contribute to the dCO2!]

• So, the PbCO3 layer that forms on the surface of the lead prevents dissolution of the metal underneath• Much like aluminum oxide layer prevents oxidation of bulk Al

• In normal water treatment processes, hypochlorous (HOCl) acid is used for disinfection, ant the PbCO3 is oxidized further to PbO2

𝑃𝑏𝐶𝑂3 𝑠 + 𝐻𝑂𝐶𝑙 + 𝐻2𝑂 ⇌ 𝑃𝑏𝑂2 𝑠 + 𝐻𝐶𝑙 + 𝐻2𝐶𝑂3(𝑎𝑞)

PbO2 is less soluble than even PbCO3 ([Pb2+] = 10-8.6[H+]4

• Release of Pb from scale, therefore, does not normally occur, unless…chemistry happens…

• If more easily oxidized species are introduced into the water, they can be oxidized by PbO2, producing Pb2+

• Two examples include chloramines (NH2Cl) and Cu(s)

• Washington, DC – 2002• Water disinfection treatment was changed from chlorine to chloramine

• By 2004, [Pb2+] had increased from average of few ppb to > 60 ppb, with maximum concentrations of > 300 ppb

• Problem not unique to Washington, DC or Flint, MI (few ppm!) or Greensville, NC (hundreds ppb)

• Approximately 18 million Americans live in communities where the water systems are in violation of the law, with many systems exceeding levels measured in Flint!

• Some cities add phosphate to drinking water to produce insoluble Pb3(PO4)2 scale to protect lead pipes

• Whereas compounds of Pb2+ are ionic, most of Pb4+ compounds are covalent (with exception of PbO2)

• Most important covalent compounds of Pb4+ are tetraalkyl compounds, such as PbR4; especially tetramethyllead [Pb(CH3)4)] and tetraethyllead [Pb(C2H5)4]

• These compounds found widespread use as additives to gasoline; tetraalkylPb compounds are volatile, and so entered the environment as gaseous form (from gasoline)

• PbR4 compounds are not water soluble, but are readily absorbed through skin

• In human liver, PbR4 molecules are converted to more toxic PbR3+ ions, which

are neurotoxins (i.e., can cross the blood-brain barrier)

• In contrast to Hg, Pb-methylation does not occur naturally (so, all PbR4 still in environment likely originated from leaded gasoline• Which became prohibited in 1991!

• When PbR4 additives are used in gasoline, atomic Pb is liberated by combustion• Must be removed to prevent metallic deposits and damage to engine

• Ethylene dibromide and ethylene dichloride are also added to gasoline to oxidize metallic Pb, forming lead halides

• Lead halides photo-oxidize to PbO, which exists in particulate form (i.e., as an aerosol)• Lifetimes of hours to days, so long-range transport possible

• Similar to Montreal Protocol limiting CHCs, elimination of Pb from gasoline is one of the few environmental success stories

• In many countries (especially developing), the use of leaded gasoline continues; in these areas the air is the major source of Pb ingested by humans

• Most ingested Pb in humans is initially present in blood, but that amount eventually reaches a tipping point where any excess Pb enters soft tissues (e.g., organs) and bones

• Toxicity of Pb is proportional to amount present in soft tissues

• Health Effects of Pb:• dysfunctional sperm in males

• inability of to bring the fetus to term in females

• deterioration of bones in adults

• interfere with normal development of brains in children; can affect children’s behavior and attentiveness and possibly lower IQs

Toxicity of Pb

Pb in Blood in Children

1976 – 1980

1988 – 1991

Pb in Blood by Age

• Lingering effects of Pb pollution

• Dramatic decrease with phasing out of leaded gasoline

• Seem to be converging on non-zero minimum concentration

• On atom-for-atom basis, Pb is not as dangerous as Hg

• But, general population is exposed to Pb from a greater variety of sources and at higher levels than for Hg

• Both metals are more toxic in their organometallic form than as the simple inorganic cations

• In terms of concentration, Pb is much closer – within a factor of 10 – to levels at which overt signs of poisoning become manifest than is any other substance, including Hg

• It is still appropriate for society to continue taking steps to further reduce human exposure to Pb• Dust, lead paint, lead pipes, industrial manufacturing, mining/smelting

Pb vs Hg

• Most common forms in environment are

• As3+ (arsenite) is more toxic• Retained in body longer than As5+ as it becomes bound to sulfhydryl groups in several

enzymes

• As5+ reduced to As3+ by (for example)• Microbes

• Dissolved sulfide: 𝐻2𝐴𝑠𝑂4− + 𝐻2𝑆 + 𝐻+ ⇌ 𝐻3𝐴𝑠𝑂3 +

1

8𝑆8 + 𝐻2𝑂 Eo = +0.521 V

• Interestingly, as As becomes methylated in the liver, it does not bind tightly to enzymes and is largely detoxified

Arsenic

Arsenite Arsenate

• Anthropogenic sources of As include:• Critical component in herbicides

• Na3(AsO3 and Cu3(AsO3)2

• Also present in insecticides• Pb3(AsO4)2 and Ca3(AsO4)2

• Release during mining/smelting of Au, Pb, Cu and Ni

• Coal combustion

• As-contaminated groundwater

• Note: Natural levels of As in water can be high• Commonly more of a health problem than anthropogenic sources

• More than 200 As-containing minerals exist

As in Drinking Water

• For once, it’s not our fault…

• Groundwater As contamination due to geology

• For example,• Bengal Delta has very high levels of As in the groundwater

• In 1970s-80s, UNICEF funded building of tens of millions of tube wells in effort to

• reduce epidemics (diarrhea, cholera and other water-borne diseases)

• Reduce child mortality due to microbially unsafe water from streams, ponds and shallowwells

• Wells affected about 50 million persons living in Bangladesh

• As levels in water from these wells can approach 1 ppm!

• Compared to guideline levels of 10 ppb.

• WHO has called this the “largest mass poisoning of a population in history.”

• Where does the As come from?

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• Inorganic arsenate ion (𝐴𝑠𝑂43−) co-precipitates or is adsorbed onto iron oxides

(𝐹𝑒2𝑂3) in soil

• If DOC consumes O2, anerobic conditions result in reduction of Fe3+ to Fe3+

• Reduced iron is much more soluble than Fe3+

• As(V) previously bound to iron oxide dissolves in the water along with Fe2+

• Reduction of As(V) to As(III) also enhances its water solubility

• So, the question is “Is the enhanced mobilization of As due to natural or human activity?”

• Actually, human impact is likely accelerating the natural mechanism of As release into the groundwater• As more water is drawn from the deeper wells, shallower sources replenish it• Replenishment from shallower sources carries more DOM to the deeper aquifers• More O2 is consumed by oxidation of DOM• Water environment becomes more reducing

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Deep aquiferLow dO2, low pE, high pH

As(V) As(III)Fe(III) (Fe(II)

Shallow aquiferHigh DOC

As(V)Fe(III)

Irrigation, drinkingAs(III) Fe(II)

• Problem is not limited to Bengal Delta

• Almost 10% of deaths among Chilean adults is attributable to As poisoning

• About 60 million Americans ingest water with [As] > 1 ppb

• West coast counties drinking water often exceeds guidelines

• Most common means of As remediation is to• Pump deep water to shallow ponds or over a well aerated surface

• Fe2+ rapidly oxidized by surface water to make FeO3 precipitate

• As rapidly adsorbs to FeO3 particle surface

• Sometimes Fe is added, further oxidizing As(III) to As(V), which is more readily adsorbed to surfaces

• LD50 values for As in humans• Amount of material that causes death of 50% of group of test animals

Water Standards for Heavy Metals

• Pesticides – synthetic chemicals that kill or otherwise control unwanted organisms• Insecticides and pesticides are common classes

• 109 kg used annually in North America alone!• About half used for agriculture

• Majority use of insecticides is for growing cotton (GMO cotton reducing need for insecticides)

• Majority use of herbicides is for growing corn and soybeans

• Half of foods eaten in US contain measurable levels of at least one pesticide!

• Pesticides are not a modern invention• Ancient Greeks used fumigants (such as SO2 from burning solid sulfur)

• In reality, humans did not invent pesticides• Plants produce natural pesticides that could be more toxic to humans than synthetic ones

Toxic Organic Compounds in the Environment

• Original “dirty dozen” listed as Persistent Organic Pollutants (POPs) by the United Nations Environmental Program (UNEP)• Largely developed during WW II and afterward

• Now banned or being phased out (Stockholm Treaty)

• Stability against environmental decomposition or degradation

• Very low solubility in water

• High solubility in fatty material in living matter

• Relatively high toxicity to insects but “low” toxicity to humans

• A few organochlorine insecticides remain in widespread use

Organochlorine Insecticides

POP List

DDT

Aldrin

Dieldrin

Endrin

Chlordane

Heptachlor

Hexachlorobenzene (HCB)

Mirex

Toxaphene

Chlordecone

Endosulfan

Pentachlorobenzene (PCB)

• Generally concerned with what happens to synthetic organic compounds once they are released into the general environment

• Chemical stability – inversely related to how and at what rate the organic degrades

• Mobility – mechanisms and rates at which organics are transported through environmental compartments

• The relative importance of the two fates is inversely related• If chemical stability is poor (i.e., fast degradation), then organic doesn’t have time to be

transported and rates of mobility may be less important

• If transportation is fast, then different degradation pathways may be operative as organics are transported to a new environment

Stability and Mobility of Organic Compounds

• Most prominent airborne organochlorine pesticide• Domestic insect repellant and deodorizer

• Crystalline solid with high vapor pressure

• Animal carcinogen that accumulates in the environment

• Likely responsible for greatest carcinogenic risk of all indoor VOCs

• Present in blood of most US residents

Para-dichlorobenzene (PCB)

• Rachel Carson dubbed DDT an “elixir of death” (1962)• First widespread use during WW II, to kill fleas (typhus carriers) and mosquitos (malaria

carriers)

• Adopted widely for agriculture in U.S. post-WW II

• Metabolite of DDT (DDE) interferes with Ca distribution in birds (thin egg shells)

• DDT is a legacy source (of pollution) because• Extremely low vapor pressure (volatilization half-life ~ 200 years)

• Very low solubility in water

• Low reactivity to light, chemicals and microorganisms

Para-dichlorodiphenyltrichloroethane (DDT)

Bioconcentration vs. Bioaccumulation

• POPs inherently more soluble in organics than water

𝐾𝑜𝑤 =𝑆 𝑜𝑐𝑡𝑎𝑛𝑜𝑙

𝑆 𝑤𝑎𝑡𝑒𝑟

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POP List Log Kow

DDT 6.2

Dieldrin 6.2

HCB 5.5-6.2

Mirex 6.9-7.5

Toxaphene 5.3

Malathion 2.9

Parathion 3.8

Atrazine 2.2-2.7

• Important to keep in mind that chemical breakdown of organics doesn’t necessarily mean loss of toxicity

• For example,• Generally true that dechlorination = detoxification

• Here, toxicity of DDE < DDT (although DDE is still toxic)

• But,

• Toxicity paraoxon > parathion

Metabolites of Organic Pollutants

• Photolytic degradation

• Depends on:• Where the organic resides

• Must be in gas phase or on a surface exposed to the photons

• Energy of photons• Must be sufficient to break chemical bonds

• Photolytic degradation can also take place in the presence of a photosensitizer• Photosensitizers are typically semi-conductors (like TiO2, UO2, VO2)

• This is termed indirect photolysis

General Classes of Degradation Reactions

• Hydrolysis• Can occur with (biotic) or without (abiotic) aid of microorganisms

• Abiotic hydrolysis is often much slower• Sometimes abiotic rate is augmented by a catalysis in the water/soil system

• Most effective for thiols, ethers and amide-containing organics

• Oxidation• Most oxidation mechanisms are mediated by microorganisms

• Common abiotic mechanisms involve OH-radical, singlet oxygen (O2[1D]), H2O2, O3

• Reduction• Occurs only under low pE conditions

• Dominant in some wastewater streams, where BOD is high

• Describes the combined aspects of chemical degradation and mobility

• Important because leachable compounds can potentially move to other (perhaps more sensitive) parts of the environment

• One way of describing leachability is the groundwater ubiquity score (GUS) index

𝐺𝑈𝑆 = log 𝑡ൗ1

2

𝑠𝑜𝑖𝑙 𝑥 3.77 − 𝑙𝑜𝑔𝐾𝑑

𝑓𝑂𝑀• Kd is the distribution coefficient between soil and solution phase

• fOM is the mass fraction of the soil that is comprised of organic matter

• 𝑡 1/2 = half-life for chemical degradation in the soil

Leachability

• GUS values ≲1.8 characterizes a species not prone to leaching

• GUS values ≳2.8 characterize highly leachable species

• Organophosphates• All derivatives of phosphoric acid

• O or S atom double bonded to P atom

• Two methoxy (-OCH3) or ethoxy (-OCH2CH3) singly bonded to P atom

• Longer R group singly bonded to P atom differentiates one organophosphate from another

• Toxicity stems from their inhibition of enzymes in the nervous system

• Largest use is in agriculture, although present in some home products (insecticides)

• Generally not persistent in the environment

• Don’t bioaccumulate as much as organochlorines (good!)

• Acutely more toxic to humans and other mammals (well…poop!)

Other Organic Pollutants

• Most organophosphates in the environment decompose by hydrolysis reactions• Yield ester intermediates, ultimately producing non-toxic alcohols, thiols and phosphoric acid

• Carbamates – analogs of organophosphates, where central P-atom has been replaced by carbon

• Triazenes• Used primarily as herbicides

• Typically, R1 = Cl and R2 and R3 are amino groups

• Most common example of triazene herbicides is atrazine

• Although only slightly toxic to mammals, atrazine can kill sensitive plants in water systems

• Atrazine is degraded by microbes and it only persists in soils for a few months

• However, once metabolites enter waterways, their half lives extend to several years

• Atrazine is not removed by typical treatments of drinking water

• Natural pesticides also exist• For example, Pyrethrins

• Obtained from flowers of a species of crysanthemum

• Typically considered by-products of pesticide production and other anthropogenic processes• Only present in trace amounts, but of greater toxicity than pesticides

• For example, Agent Orange is a mixture of the herbicides 2,4-D and 2,4,5-T

• Used extensively as a defoliant during the Vietnam War

• The defoliant mixture contained dioxin levels of ~3 ppm

• The dioxin has contaminated the soil

• Increased cancers observed among soldiers exposed to Agent Orange

• Originally, Agent Orange was tested by U.S. armed forces near Gagetown, New Brunswick, Canada in 196os

• Soil there is still contaminated with dioxins!!!

• Dioxins bioaccumulate and bioamplify

Dioxins

• Used extensively as coolant fluids in power generation because they are• Chemically inert liquids that are difficult to burn

• Have very low vapor pressures

• Are inexpensive to produce

• Are excellent electrical insulators

• PCBs are persistent in the environment

• Since 1950s, over 1 million metric tons have been produced• About 10% of PCBs produced are still in the environment!

• PCBs are listed as another of the UN “dirty dozen” POPs

• Production halted in 1977, although PCBs remain in “closed” use in some electrical transformers still.

Polychlorinated biphenyls (PCBs)

2,3’,4’,5’-tetrachlorobiphenyl

• PCBs persist for many years in the environment

• Water solubility is very low• Predominantly associated as adsorbed species on suspended matter in particles

• The little bit that does dissolve in water, volatilizes over time and subsequently redeposited on land or in water after traveling in air for a few days• There are measurable background levels of PCBs even in the polar regions and at the bottom of

oceans!

• Soluble in fatty tissue, so PCBs subject to bioaccumulation andbioamplification