Environmental Chemistry

Chapter 11:

Water and the Hydrosphere

Copyright © 2011 by DBS

Contents

• • • • • • • • • • • The Fantastic Water Molecule and the Unique Properties of Water The Hydrosphere Compartments of the Hydrosphere Aquatic Chemistry Alkalinity and Acidity Metal Ions Oxidation-Reduction Complexation and Chelation Interactions with Other Phases Aquatic Life Microbially Mediated Elemental Transitions and Cycles

The Fantastic Water Molecule and the Unique Properties of Water

The Fantastic Water Molecule and the Unique Properties of Water Region of partial negative charge Regions of partial positive charge

Polarity

A difference in the electronegativities of the atoms in a bond creates a

polar bond

Partial charges result from bond polarization A

polar covalent bond is

a covalent bond in which the electrons are not equally shared, but rather displaced toward the more electronegative atom

H-Bonding

Polarized bonds allow

hydrogen bonding

to occur • • • A

hydrogen bond

is an electrostatic attraction between an atom bearing a partial positive charge in one molecule and an atom bearing a partial negative charge in a neighboring molecule The H atom must be bonded to an O, N, or F atom Hydrogen bonds typically are only about one-tenth as strong as the covalent bonds that connect atoms together

within

molecules H –bonds are intermolecular bonds Covalent bonds are intramolecular bonds

Unique Properties

• • • • • • • • Water shrinks on melting (ice floats on water) Unusually high melting point Unusually high boiling point Unusually high surface tension Unusually high viscosity Unusually high heat of vaporization Unusually high specific heat capacity And more…

There is No Substitute for Water

Box 1.1 Major Properties of Water

Unique Properties

Unusually high Mpt. and Bpt.

Predicted melting point at 73 ºC and boiling point at 98 ºC.

Unique Properties

Why H-Bonding is Important This increase in the ‘thermal window’ of liquid water from a predicted 25º to its actual 100º allows aquatic life to exist over a broader range of temperatures H-bonding leads to viscosity and surface tension

Unique Properties

• Unlike other substances water is less dense in solid form than liquid form Becoming less dense • Water at different temperatures has different densities – leads to layering in lakes • D ~ 1/V …as ice melts D inc. and V dec.

Unique Properties

Ice shrinks on melting as 15% H-bonds are lost A certain mass of ice occupies more space than the same mass of water

The Hydrosphere

Natural Waters

The Blue Marble 0.001 % water vapor 71 % liquid water

The Blue Marble

is a famous photograph of the Earth taken on 7 December 1972 by the crew of the

Apollo 17

spacecraft at a distance of about 29,000 km or about 18,000 miles. It is one of the most widely distributed photographic images in existence. The image is one of the few to show a fully lit Earth, as the astronauts had the Sun behind them when they took the image. To the astronauts, Earth had the appearance of a child's glass marble (hence the name).

• • • • • Compartments Atmosphere Land Groundwater Rivers lakes Oceans

The Hydrosphere

Hydrologic Cycle Source: http://www.nasa.gov/vision/earth/environment/warm_wetworld.html

Where Does

Potable

Sources

(fit for consumption) Drinking Water Come From?

Less than one third of salt-free water is liquid

Surface water

: from lakes, rivers, reservoirs (< 0.01 % of total)

Ground water

: pumped from wells drilled into underground aquifers (0.3 %)

Natural Waters

Uses of Water World Resources 1998-99

The number of people living in countries facing severe or chronic water shortages is

Sources

estimated 505 million people today to between 2.4 and 3.2 billion people by 2025.

< 1000 m 3 per person per year Engelman

et al

., 2000

Access to Water

Access to Water

Uneven distribution of water

Region

World Asia Europe Middle East/N. Africa N. America

Total Renewable Water Resources (km 3 yr -1 )

43,249 11,321 6,590 518 4,850

Total Water Withdrawals (m 3 yr -1 )

3,414,000 1,516,247 367,449 303,977 512,440

(m 3 Per Capita person -1 yr -1 )

650 1,028 503 754 1,720

Average % of Renewable Resources

29 9 423 14

Average % Used by Agriculture

71 79 25 80 27

Average % Used by Industry

20 10 48 5 58 Subject to contamination Using water at a rate faster than it can be supplied (>100 due to use of sea water)

Natural Waters

Role of Water in the Environment • Water is an important constituent in our body and our survival depends on natural waters – – –

transports distributes reduces

substances into, within, and out of living organisms soluble substances (e.g. pesticides, lead, mercury) concentrations via dilution and dispersion e.g. rainwater carries substances (e.g. acids) down to earth’s surface, washes out (cleanses) the air but pollutes waterways

Withdrawls (2000) Fastest growing areas are most water deficient: S. CA, AZ, NV, CO Precipitation CONUS

Provides electricity from hydroelectric plants for 30 million people (1/10 th of the U.S. population) Glen Canyon Dam Hoover Dam

Ogallawa (High Plains) Aquifer

• • • • • World’s largest aquifer Composed of fossil water from last ice age Rapidly dropping water table supports $32 billion agriculture most areas water withdrawn much faster than recharge

Transport medium volumes,

Liquid Medium

times, Water Cycle fluxes Largest reservoir – oceans τ = 40,000 yr

Compartments of the Hydrosphere

Compartments of the Hydrosphere • Surface waters (watersheds) – streams, lakes reservoirs, wetlands, estuaries

– Standing surface water vs. flowing surface water

Compartments of the Hydrosphere

• Temperature-density relationship leads to layering in lakes Warmer water floats on colder = thermal stratification

Compartments of the Hydrosphere • Groundwater – most from precipitation and infiltration

– Composition depends on surrounding rock formations (porosity and permeability)

Aquatic Chemistry

Aquatic Chemistry

• Algal photosynthesis: Converts inorganic C (2HCO 3 ) to organic form (CH 2 O, emp. formula for sugars) CO 3 2 is either converted back to HCO 3 , or ppts as limestone Biomass (CH 2 O) produced

Aquatic Chemistry

• Most redox reactions in water are catalyzed by bacteria – e.g. N compounds to NH 4 + – e.g. N to NO 3 in anoxic conditions in oxic conditions • Chelation of metals • Gas exchange with atmosphere • Solute exchange between aquesous and solid phases (sediments)

Alkalinity and Acidity

Alkalinity and Acidity

• Alkalinity – the capacity of water to accept H + – Measure of the ability of a water body to neutralize acidity – Serves as a pH buffer and reservoir for inorganic C – Helps determine ability of water to support algal growth and aquatic life, used as a measure of water fertility

• Dissolution of limestone and other minerals produces alkalinity e.g.

CaCO 3 CO 3 2 ⇌ Ca 2+ + H 2 O ⇌ + CO 3 2 HCO 3 + OH • Water supply with high total alkalinity is resistant to pH change (>> buffering capacity) • Two samples with identical pH but different alkalinity behave differently on addition of acid – Different capacity to neutralize acid – pH is an

intensity

factor whilst alkalinity is a

capacity

factor

Alkalinity

• Measurement of the buffer capacity (resistance to pH change) • • • e.g. Carbonate neutralization reaction CO 3 2 + H + ⇌ HCO 3 Bicarbonate neutralization reaction HCO 3 + H + ⇌ H 2 O.CO

2 ⇌ H 2 O + CO 2 Hydroxide neutralization reaction H + + OH ⇌ H 2 O Alkalinity = [OH ] + [HCO 3 ] + 2[CO 3 2 ] – [H + ] Units are mg L -1 CaCO 3 or mEq L -1 (regardless of species) Acid titration to change the pH to 4.5 (methyl orange end-point) If pH < 4.5 there is no acid neutralizing capacity i.e. no need to measure alkalinity

Acidity

pH = - log [H + ] • H + usually surrounded by water of hydration, written H 3 O + • ‘Master Variable’ – controls parameters e.g. speciation • Ranges 5.5 - 9

Acidity

• Acidity results from presence of weak acids: H 2 PO 4 , CO 2 , H 2 S, proteins, fatty acids, metal ions (e.g. Al 3+ , Fe 3+ ) e.g. [Al(H 2 O) 6 ] 3+ + H 2 O ⇌ [Al(H 2 O) 5 OH] 2+ + H 3 O + simplifies as [Al(H 2 O) 6 ] 3+ ⇌ [Al(H 2 O) 5 OH] 2+ + H + • Difficult to measure due to volatility of gases • Total acidity is determined by titration with base to pH 8.2

Metal Ions and Calcium in Water

Metal Ions and Calcium in Water

Metal Ions • • • M n+ exists in various forms in water (species) Cannot exist as free ion, seeks max stability of outer e shells, does this by accepting lone pairs from donor molecules Exist as hydrated cations [M(H 2 O) x ] n+ other bases (e donors) coordinate bonded to water molecules or

Metal Ions and Calcium in Water

Metal Ions • The hydrogen atoms attached to the water ligands are sufficiently positive that they can be pulled off in a reaction involving water molecules in the solution.

Metal Ions and Calcium in Water

Metal Ions • Allows for loss of H + , reactions: Acid base: Ppt: [Fe(H 2 O) 6 ] 3+ ⇌ [FeOH(H 2 O) 5 ] 2+ + H + [Fe(H 2 O) 6 ] 3+ ⇌ Fe(OH) 3 (s) + 3H 2 O + 3H + (results from A-B) Redox: [Fe(H 2 O) 6 ] 2+ ⇌ Fe(OH) 3 (s) + 3H 2 O + e + 3H + • Due to these reactions conc. of the hydrated cation, e.g. [Fe(H 2 O) 6 ] 3+ is very small

Metal Ions and Calcium in Water

Metal Ions • Acid-base reaction is more completely shown by H + water molecule in soln: ion is being pulled off by a [Fe(H 2 O) 6 ] 3+ + H 2 O ⇌ [FeOH(H 2 O) 5 ] 2+ + H 3 O + • Successive deprotonations: [FeOH(H 2 O) 5 ] 2+ +H 2 O ⇌ [Fe(OH) 2 (H 2 O) 4 ] + + H 3 O + [Fe(OH) 2 (H 2 O) 4 ] + +H 2 O ⇌ [Fe(OH) 3 (H 2 O) 3 ](s) + H 3 O + • Forms a neutral complex which does not dissolve and precipitates, Fe(OH) 3

Metal Ions and Calcium in Water

Metal Ions Hydrated Metal Ions as Acids • Hydrated metals with +3 charge or more act as Bronsted acids (inc with charge, dec with radius) e.g. [Fe(H 2 O) 6 ] 3+ ⇌ [FeOH(H 2 O) 5 ] 2+ + H + • Solutions containing +3 hexaaqua ions tend to have pH's in the range from 1 to 3. Solutions containing +2 ions have higher pH's - typically around 5 - 6, although they can go down to about 3.

• Tendency of hydrated metal ions to act as acids leads to acid mine water [Fe(H 2 O) 6 ] 3+ ⇌ Fe(OH) 3 (s) + 3H 2 O + 3H +

Metal Ions and Calcium in Water

Metal Ions • Properties of metals dissolved in water depend upon the nature of metal species dissolved in water, called speciation • In addition to hydrated [M(H 2 O) x ] n+ and the associated hydroxo species, metals may exist as complexes (reversibly bound to inorganic anions, organic compounds) or organometallic compounds

Metal Ions and Calcium in Water

Calcium and Harness • Ca 2+ generally has highest conc. And most influence on aquatic chemistry • Why?

• Calcium is a key element in many geochemical processes • Primary minerals: gypsum (CaSO 4 .2H

2 O), anhydrite (CaSO 4 ), dolomite (CaMg(CO 3 ) 2 , calcite and aragonite (CaCO 3 )

Metal Ions and Calcium in Water

Calcium and Harness CO 2 (g) + H 2 O(aq) ⇌ H 2 CO 3 (aq) H 2 CO 3 (aq) ⇌ H + + HCO 3 CaCO 3 (s) ⇌ CO 3 2 + H 2 O Ca 2+ ⇌ + HCO CO 3 3 + 2 OH H + + OH ⇌ H 2 O CaCO 3 (s) + CO 2 (g) + H 2 O(aq) ⇌ Ca 2+ + 2HCO 3 -

K

H

K

a

K

sp

K

b 1/

K

w • Giant titration of acid from atmospheric CO 2 with base from carbonate ion in rocks

K

=

K

sp

K

b

K

H

K

a /

K

w = 1.5 x 10 -6 = [Ca 2+ ][HCO 3 ] 2

P

CO2

Metal Ions and Calcium in Water

Calcium and Harness • CaCO 3 (s) + CO 2 (g) + H 2 O(aq) ⇌ Ca 2+ + 2HCO 3 -

K

=

K

sp

K

b

K

H

K

a /

K

w = 1.5 x 10 -6 = [Ca 2+ ][HCO 3 ] 2

P

CO2 If [Ca 2+ ] =

S

, [HCO 3 ] = 2

S

1.5 x 10 -6 = [Ca 2+ ][HCO 3 ] 2

P

CO2 =

S

(2S) 2 0.00037 atm

S

= [CO 2 ] = 5.2 x 10 -4

S

= [Ca 2+ ] = 5.2 x 10 -4 mol L -1 (34 x amount calculated from Henry’s law) mol L -1 (this is 4x closed system) [HCO 3 ] = 2

S

= 1.0 x 10 -3 mol L -1 Acid-base reaction increases the solubility of both the gas and he solid – water that contains CO 2 more readily dissolves calcium carbonate

Metal Ions and Calcium in Water

Calcium and Harness • CO 3 2 , H + and OH can be derived

K K

sp b = [Ca 2+ ][CO [CO 3 2 ] 3 2 = [HCO 3 ][OH ] ]   [CO 3 2 ] = 8.8 x 10 -6 mol L -1 [OH ] = 1.8 x 10 -6 mol L -1

K

w = [H + ][OH ]  [H + ] = 5.6 x 10 -9 mol L -1 Conclude natural water at 25 °C with a pH determined by saturation with CO 2 CaCO 3 should be alkaline (pH = 8.3) and Actual value of calcereous waters is around 7 9 …why the difference?

Metal Ions and Calcium in Water

Calcium and Harness • Simple model does not include CO 2 • MO’s directly affect conc. of Ca 2+ from respiration of MO’s!

in water

Metal Ions and Calcium in Water

Calcium and Harness Ca

2+

Mg

2+

Fe

2+

Common cations of high enough concentration to be readily monitored are good indicators of pollution events

Metal Ions and Calcium in Water

Calcium and Harness

Hard water

contains high concentrations of dissolved calcium and magnesium ions

Soft water

contains few of these dissolved ions.

Hardness = [Ca

2+

] + [Mg

2+

]

Counter ions of alkalinity ions Alkalinity is a good indicator of hardness and vice-versa Carbonate minerals: limestone - CaCO 3 dolomite - CaCO 3 .MgCO

3 sulfates - CaSO 4 (also Al 3+ , Fe 3+ , Mn 2+ and Zn 2+ )

Metal Ions and Calcium in Water

Calcium and Harness • • • Deposition of white solid CaCO 3 or MgCO 3 when water is heated – ‘furring-up blocks pipes and lowers efficiency of industrial processes Formation of scum (insoluble ppt) with soap and water Ca 2+ (aq) + 2Na(C 17 H 33 COO ) (aq) – detergent action is blocked Staining (due to transition metals) 2Na + + Ca(C 17 H 33 COO ) 2(s) A pipe with hard-water scale build up

Metal Ions and Calcium in Water

Calcium and Harness • Solid deposit = carbonate hardness or temporary hardness Ca 2+ + 2HCO 3 ⇌ CaCO 3 (s) + CO 2 (g) + H 2 O(aq) (removed via boiling) – Causes deposit in pipes and scales in boilers – Temporary hard water has to be softened before it enters the boiler, hot-water tank, or a cooling system • No solid = non-carbonate or permanent hardness – Amount of metal ions that can not be removed by boiling Total hardness = temporary hardness + permanent hardness

Oxidation-Reduction

Oxidation-Reduction

Most important oxidizing agent is dissolved O 2 (atmospheric)

Acidic solution

O 2 + 4H + + 4e ⇌ 2H 2 O

Basic solution

O 2 + 2H 2 O + 4e ⇌ 4OH − O 2 is reduced from 0 to -2 state in H 2 O or OH − Concentration of O 2 in water is low (10 ppm average), governed by Henry’s law: O 2 (g) ⇌ O 2 (aq)

K

H = [O 2 (aq) ]

P

O2 At 25 °C, K H = 1.3 x10 -3 mol L -1 atm -1 Dissolved O 2 influences chemical speciation of elements in natural and polluted waters

Oxidation-Reduction

• Show that (a) O 2 + 2H 2 O + 4e Is the same as (b) 2H 2 O + 2e ⇌ 4OH − (from above) ⇌ H 2 (g) + 2OH (from Manahan) Double (b): 2(2H 4H 2 2 O + 2e O + 4e ⇌ ⇌ H 2 (g) + 2OH 4 H 2 (g) + 4OH Add O 2 + 2H 2 ⇌ 2H 2 O O 2 + 2H 2 O + 4e ⇌ 4OH −

Question

P9-1: Confirm by calculation the value of 8.7 mg L -1 water at 25 °C for the solubility of oxygen in At 25 °C,

K

H = 1.3 x10 -3 mol L -1 atm -1

K

H = [O 2 (aq) ] /

P

O2 [O 2 (aq) ] = K H x

P

O2 [O 2 (aq) ] = (1.3 x10 -3 mol L -1 atm -1 ) x 0.21 atm = 2.7 x 10 -4 mol L -1 [O 2 (aq) ] = 2.7 x 10 -4 mol L -1 x 32.00 g mol -1 = 8.7x 10 -3 g L -1 = 8.7 mg L -1 = 8.7 ppm

Oxidation-Reduction

• • • • Depletion of O 2 Temperature (inc) Pressure (dec) Salts (inc) Organic matter (inc) Dissolved O 2 decreases with increasing temperature

Oxidation-Reduction

Oxygen Demand • The most common substance oxidized by DO in water is organic matter (plant debris, dead animals etc.) 0 to -2 CH 2 O (aq) + O 2(aq) → CO 2(g) + H 2 O (aq) • • • 0 to +4 Similarly DO is consumed by NH 3 and NH 4 + in the nitrification process Water in streams and rivers are constantly replenished with oxygen Stagnant water and deep lakes can have depleted oxygen

Oxidation-Reduction

Oxygen Demand Half reactions Oxidation: Reduction: CH 2 O + H 2 O → CO 2 4H + + O 2 + 4e + 4e → 2H 2 O CH 2 O (aq) + O 2(aq) + 4H → CO 2(g) + + H 2 O (aq) In basic conditions?

O 2 + 4H + + 4e  2 H 2 O React with hydroxide O O O 2 2 2 + 4H + 4H + 2H + 2 2 + 4OH O + 4e O + 4e + 4e    2H 2 O + 4OH 4OH 2H 2 O + 4OH Same overall

Question

P9-4: Determine the balanced redox reaction for the oxidation of ammonia to nitrate ion by O 2 in alkaline solution (basic) Does this reaction make the water more basic or less?

NH 3 + O 2  NO 3 + H 2 O Using standard redox balancing techniques: NH 3 + 2O 2 + OH  NO 3 + 2H 2 O The water becomes less basic since OH is removed

Measures of amount of organics/biological species in water

• • • • • Biochemical Oxygen Demand (BOD) Chemical Oxygen Demand (COD) Total Organic Carbon (TOC) Dissolved Organic Carbon (DOC) (TOC)-(DOC) = Suspended carbon in water

Oxidation-Reduction

Biological Oxygen Demand • • The capacity of the organic and biological matter in a sample of natural water to consume oxygen, a process usually catalyzed by bacteria, is called BOD Procedure: measure O 2 in the stream or lake. Take a sample and store at 25 o C for five days and remeasure O 2 content. The difference is the BOD – BOD 5 corresponds to about 80% of the actual value. It is not practical to measure the BOD for an infinite period of time – Surface waters have a BOD of about 0.7 mg L -1 the solubility of O 2 in water (8.7 mg L – Sewage has BOD of ~100 mg L -1 -1 ) – significantly lower than

Oxidation-Reduction

Chemical Oxygen Demand O 2 + 4H + + 4e → 2H 2 O • Dichromate ion, Cr 2 O 7 2 dissolved in sulfuric acid is a powerful oxidizing agent. It is used as an oxidant to determine COD Cr 2 O 7 2 + 14H + + 6e → 2Cr 3+ + 7 H 2 O • Excess dichromate is added to achieve complete oxidation Back titration with Fe 2+ gives the desired endpoint value # moles of O 2 consumed = 6/4 x (#moles Cr 2 O 7 consumed)

Note: Cr 2 O 7 2 is a powerful oxidizing agent and can oxidize species that are not usually oxidized by O 2 - hence gives an upper limit

Question

P9-5: A 25 mL sample of river water was titrated with 0.0010 M Na 2 Cr 2 O 7 required 8.7 mL to reach the endpoint. What is the COD (mg O 2 /L)?

and No. moles Cr 2 O 7 2 = 0.0010 mol L -1 x (8.7 x 10 -3 L) = 8.7 x 10 -6 mols No. moles O 2 = 1.5 moles Cr 2 O 7 2 = 1.5 x (8.7 x 10 -6 mols) = 1.3 x 10 -5 mols O 2 1.3 x 10 -5 mol x 32.00 g mol -1 = 4.2 x 10 -4 g 0.42 mg / 0.025 L = 17 mg L -1

Oxidation-Reduction

The pE Scale • Oxidation and reduction are controlled by the concentrations of electrons which are present: pE = - log 10 [e ] Low pE means electrons are available (reducing environment) High pE means electrons are unavailable (oxidizing environment) pE is calculated from electrode potential (

E

) by the relationship: pE =

E

2.303 RT/F

Oxidation-Reduction

The pE Scale • • When a significant amount of O 2 dominant reaction determining e is dissolved, the reduction of O 2 availability: is the ¼ O 2 + H + + e ⇌ ½ H 2 O Under such circumstances, the pE of the water is related to its acidity and to the partial pressure as follows: pE = 20.75 + log([H + ]

P

O2 ¼ )

OR

pE = 20.75 – pH + ¼ log(

P

O2 )

Oxidation-Reduction

The pE Scale A convenient approach is to use Nernst Equation of electrochemistry

E = E 0 – (RT/F) (log



[products] /



[reactants])

…for 1 electron redox process

E = E 0 - 0.0591(log



[products] /



[reactants])

where E 0 is the standard electrode potential for a one electron reduction One can equate pE to the Electrode Potential

E

pE =

E

/0.0591 or pE 0 = E 0 /0.0591

• Dividing throughout by 0.0591:

pE = pE 0 - (log



[products] /



[reactants])

Redox Chemistry in Natural Waters

The pE-pH Diagram • Nature of a chemical species Is usually a function of pH and pE • Move from pE = pE 0 - (log relating pE to pH  [products] /  [reactants]) to an equation

Redox Chemistry in Natural Waters

The pE Scale ¼ O 2 + H + + e ⇌ ½ H 2 O E 0 pE 0 = 1.23 V = 1.23/0.0591

pE = pE 0 - (log  [products] / pE = 20.75 - log 1/  [reactants])  [reactants] = 20.75 + log (  [reactants]) pE = 20.75 + log([H + ]

P

O2 ¼ ) = = 20.75 + log([H + ] + log (

P

O2 ¼ ) pE = 20.75 – pH + ¼ log(

P

O2 ) pE = 20.75 – pH + ¼ log(

P

O2 ) For a neutral sample of water that is saturated with oxygen from air (

P

O2 atm) that is free from CO 2 (pH = 7) the pE value corresponds to 13.9

= 0.21

…pE value decreases with decrease in O 2 and increase in pH

Dominant redox equilibrium reaction determines pE of water (O 2 may not be dominant redox species!)

Question

What is the most oxidizing conditions possible in water?

P

O2 cannot exceed 1, log(1) = 0 pE = 20.75 – pH This can be drawn on a pE/pH diagram as a boundary line, When pE > 20.75 – pH water will be oxidized A similar analysis gives boundary below which water will be reduced pE pH pE = 20.75 - pH pE = - pH

Redox Chemistry in Natural Waters

The pE Scale Example 1/8NO 3 − + 5/4H + + e ⇌ 1/8NH 4 + + 3/8H 2 O pE 0 = E 0 /0.0591 = 0.836/0.0591 = +14.15

E 0 = +0.836 V pE = pE 0 – log [NH 4 + ] 1/8 [NO 3 ] 1/8 [H + ] 5/4 ) = 14.15 - 5/4pH -1/8log([NH 4 + ]/[NO 3 ])

a

x =

x

log

a

log(1/b) = -log

b Note: Express the reactions as one electron reduction process

….. Follow the examples given on page 435

Question

9-7: Deduce the equilibrium ratio of concentrations of NH 4 + to NO 3 at a pH of 6.0 (a) for aerobic water having a pE = +11, and (b) for anaerobic water with pE = -3 pE = 14.15 – (5/4)pH – (1/8)log([NH 4 + ] / [NO 3 ]) 11 = 14.15 – (5/4) x 6 – (1/8)log([NH 4 + ] / [NO 3 ]) log([NH 4 + ] / [NO 3 ]) = -8(4.35) = -34.8

[NH 4 + ] / [NO 3 ] = 1.6 x 10 -35 pE = 14.15 – (5/4)pH – (1/8)log([NH 4 + ] / [NO 3 ]) -3 = 14.15 – (5/4) x 6 – (1/8)log([NH 4 + ] / [NO 3 ]) log([NH 4 + ] / [NO 3 ]) = 8(9.65) = 77.2

[NH 4 + ] / [NO 3 ] = 1.6 x 10 77

Problem 9-7

pH = 6, pE = 11, pH = 6, pE = -3,

Redox Chemistry in Natural Waters

The pE-pH Diagram Fe 3+ + e ⇌ Fe 2+ • • For this reaction, pE 0 = 13.2

pE = 13.2 + log([Fe 3+ ] / [Fe 2+ ]) NOT pH DEPENDENT!

e.g. Ratio of Fe 3+ to Fe 2+ when pE = -4.1 (reducing) -4.1 = 13.2 + log([Fe 3+ ] / [Fe 2+ ]) log([Fe 3+ ] / [Fe 2+ ]) = -17.3

[Fe 3+ ] / [Fe 2+ ] = 5 x 10 -18 (far more Fe 2+ ) • Transition between dominance of one form over the other occurs at [Fe 3+ ] = [Fe 2+ ], pE = 13.2 + log(1) = 13.2 + 0 = 13.2

pH independent

Redox Chemistry in Natural Waters

pE – pH Stability Field Diagrams Zone dominance of various oxidation states pE independent • Fe 3+ ion is stable in oxidizing acidic conditions, Insoluble Fe(OH) 3 is predominant iron species • Fe 2+ /Fe 3+ ions can only exist under acidic conditions • At higher pH Fe 3+ is present as Fe(OH) 3 . Fe(OH) significantly basic 2 does not precipitate until solution becomes • Changes in redox conditions govern whether the iron will be in solution or in the sediments pE independent

Complexation and Chelation

Complexation and Chelation

• • • M n+ exists in various forms in water Exist as hydrated cations [M(H 2 O) x ] n+ molecules or other bases (e coordinate bonded to water donors) called ligands Ligands - bond to a metal ion to form a complex ion (coordination compound) e.g. Cd 2+ + CN ⇌ [CdCN] + [CdCN] + + CN ⇌ Cd(CN) 2 Cd(CN) 2 + CN ⇌ [Cd(CN) 3 ] + CN [Cd(CN) 3 ] ⇌ Cd(CN) 4 2 (CN- is unidentate ligand)

Complexation and Chelation

• Complexes with chelating agents are more important, can be more than one bonding group on a ligand e.g. nitrilotriacetate (NTA) ligand • Has 4 binding sites, stability inc. with no. of binding sites

Ligands found in natural waters contain a variety of functional groups that can donate e -

Complexation and Chelation

• • • • Ligands may undergo redox, decarboxylation, and hydrolysis Complexation may change the oxidation state of the metal, may become: (i) solubilized from an insoluble compound and enter solution, or (ii) insoluble and removed from solution e.g. complexation with negative species can convert soluble Ni 2+ (cation) into [Ni(CN) 4 ] 2 (anion). Cations are readily bound by ion exchange processes in soils (exchange of H + with another cation), whilst anionic species are not.

Complexation and Chelation

Occurrence and Importance • • Chelating agents are common potential pollutants Occur in sewage and industrial wastes e.g. EDTA (ethylenediaminetetraacetic acid) + M n+ • Tend to solubilize heavy metals from plumbling and from waste deposits

Complexation and Chelation

Complexation by Humic Substances • • • Humic substances - Most important class of complexing agents Formed from decomposition of vegetation Classified based on extraction with strong base: (a) Humin – nonextractable plant residue (b) Humic acid – precipitates after addition of acid (c) Fulvic acid – organic material remaining in acidified solution • High molecular mass, polyelectrolytic macromolecules, e.g. fulvic acid

Complexation and Chelation

Complexation by Humic Substances • Binding of metal ions by humic substances:

Complexation and Chelation

Complexation by Organometallic Compounds • Organometallic compounds – metal attaches to organic ligand Hg 2+ (Mercury (II) ion) CH 3 Hg + (Monomethylmercury ion) (CH 3 ) 2 Hg (Dimethylmercuty) • May enter direct as pollutants or be synthesized biologically by bacteria • Common to find organometallic Hg, Sn, Se and As compounds, all highly toxic

Interactions with Other Phases

Interactions with Other Phases

• Most of the important chemical phenomena do not occur in solution, but rather through interaction of solutes in water with other phases e.g. redox reactions catalyzed by bacteria, solute-particle interactions

Interactions with Other Phases

1. Organic compounds may be present as films on the surface of water, may undergo photolysis 2. Gases are exchange with the atmosphere 3. Photosynthesis and other biological processes (e.g. biodegradation of organics) in bacterial cells 4. Particles introduced by eroding streams or precipitation of insoluble salts

Interactions with Other Phases

• Lipophillic pollutants in aquatic environment are associated with: – Particles; and – Colloidal organic carbon (natural organic matter) • Partition coefficients are used to model particle – water exchange

Aquatic Life

Aquatic Life

• Autotrophic biota – utilize solar or chemical energy to fix elements into complex molecules – Producers – autotrophs that utilize solar energy to synthesize organic matter • Heterotrophic biota – utilize organic substances produced by autotrophs for energy and as raw materials for synthesis of own biomass – Decomposers – a subclass of heterotrophs (bacteria and fungi) which break down material to form simple compounds

Aquatic Life

• Microorganisms – exist as single cell organisms – Bacteria, fungi, algae • Algae and photosynthetic bacteria: – predominant producers of biomass that supports the rest of the food chain – Catalyze chemical reactions – Break down biomass and mineralize essential elements (N, P) – Play important role in biogeochemical cycles – Breakdown and detoxify many xenobiotic pollutants

Aquatic Life

•

Algae

MO’s that consume inorganic nutrients and produce OM from CO photosynthesis 2 via

hv

CO 2 + H 2 O → {CH 2 O} + O 2 (g) • •

Fungi

Nonphotosynthetic, aerobic organisms Important role in determining composition of natural waters since decomposition products enter water (cellulose from wood and other plant materials including humic substances)

Bacteria

• • • Single celled MO’s (rods, spheres, or spirals) Characteristics – unicellular, semi-rigid cell wall, motility with flagella, multiplication via binary fission Obtain energy needed for metabolism and reproduction by mediating chemical reactions (biogeochemical cycles) • Subclasses: – Heterotrophic bacteria – Aerobic bacteria – Anaerobic bacteria – Facultative bacteria

Bacteria

• Prokaryotic bacterial cell – Enclosed in cell wall – Capsule enclosure (slime layer) – Cell membrane controls material transport – Cytoplasm contains nutrients for metabolism

Bacteria

• • • Bacterial Growth and Metabolism Reproduce rapidly, high surface-volume ratio Metabolic reactions of bacteria are mediated by enzymes

Microbially Mediated Elemental Transitions and Cycles

Microbially Mediated Elemental Transitions and Cycles

• Biogeochemical cycles – microbially mediated transitions between elemental species

Microbially Mediated Elemental Transitions and Cycles

• • • •

Carbon Cycle

Small amount is atmospheric CO 2 Large amount present in minerals (carbonates) Organic fraction as hydrocarbons Manufacture of toxic xenobiotic compounds from hydrocarbons

Microbially Mediated Elemental Transitions and Cycles

Carbon Cycle – Involvement of MO’s

• •

Photosynthesis

– algae, higher plants, bacteria use light energy to fix inorganic C CO 2 + H 2 O → {CH 2 O} + O 2 (g)

Respiration:

Aerobic respiration – OM is oxidized {CH 2 O} + O 2 (g) → CO 2 + H 2 O • Anaerobic respiration – uses oxidants other than O 2 , NO 3 or SO 4 2-

Degradation of biomass

– by bacteria and fungi. Prevents accumulation of wastes, converts organic C, N S, P into inorganic forms for use by plants •

Methane production

– in anoxic sediments 2{CH 2 O} → CH 4 + CO 2 •

Bacterial utilization and degradation of HC’s

– oxidation of HC’s •

Biodegradation of organic matter

– treatment of wastewater

Microbially Mediated Elemental Transitions and Cycles

Nitrogen Cycle

• • N is interchanged among the atmosphere, OM, and inorganic compounds MO’s mediate reactions

Microbially Mediated Elemental Transitions and Cycles

•

Nitrogen fixation

– binding of atmospheric N 2 3{CH 2 O} + 2N 2 + 3H 2 O + 4H + → 3CO 2 + 4NH 4 + •

Nitrification

– converts ammonium to nitrate 2O 2 + NH 4 + → NO 3 + 2H + + H 2 O •

Nitrate reduction

– N in compounds is reduced by MO’s to lower oxidation states •

Denitrification

– produces N 2 , N 2 O or NO, returns to atmosphere

Microbially Mediated Elemental Transitions and Cycles

• Microbial transformations of Sulfur – Reduction of sulfate, oxidation of sulfide, degradation of organis S compounds • Microbial transformations of Phosphorus • Microbial transformations of halogens – Operate on xenobiotic compounds • Microbial transformations of Iron – Oxidize iron (II) to iron (III)