ch11 water and the hydrosphere

Environmental Chemistry

Chapter 11:Water and the Hydrosphere

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 electronegativitiesof the atoms in a bond creates a polar bond

A polar covalent bond is acovalent bond in which theelectrons are not equally shared,but rather displaced toward themore electronegative atomPartial charges result

from bond polarization

H-Bonding

• 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 bondsCovalent bonds are intramolecular bonds

Polarized bonds allow hydrogen bonding to occur

• 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…

Unique Properties

There is No Substitute for WaterBox 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.

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 Why H-Bonding is Important

Unique Properties

• Unlike other substances water is less dense in solid form than liquid form

• Water at different temperatures has different densities – leads to layering in lakes

• D ~ 1/V …as ice melts D inc. and V dec.

Becoming less dense

Unique PropertiesIce 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 WatersThe Blue Marble

71 % liquid water

0.001 %water vapor

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).

The HydrosphereHydrologic Cycle

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

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

SourcesWhere Does Potable (fit for consumption) Drinking Water Come From?

Surface water: from lakes, rivers, reservoirs (< 0.01 % of total)Ground water: pumped from wells drilled into underground aquifers (0.3 %)

Less than one third of salt-free water is liquid

Natural Waters Uses of Water

World Resources 1998-99

SourcesThe number of people living in countries facing severe or chronic water shortages is projected to increase more than fourfold over the next 25 years. This will be from an estimated 505 million people today to between 2.4 and 3.2 billion people by 2025.

Engelman et al., 2000

< 1000 m3 per person per year

Access to Water

Uneven distribution of water

Region Total Renewable Water Resources

(km3 yr-1)

Total Water Withdrawals

(m3 yr-1)

Per Capita (m3 person-1 yr-1)

Average % of Renewable Resources

Average % Used by Agriculture

Average % Used by Industry

World 43,249 3,414,000 650 - 71 20

Asia 11,321 1,516,247 1,028 29 79 10

Europe 6,590 367,449 503 9 25 48

Middle East/N. Africa

518 303,977 754 423 80 5

N. America 4,850 512,440 1,720 14 27 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 substances into, within, and out of living organisms– distributes soluble substances (e.g. pesticides, lead, mercury)– reduces 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)

Precipitation CONUS

Fastest growing areas are most water deficient: S. CA, AZ, NV, CO

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

Hoover Dam

Glen Canyon 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

Liquid Medium Water Cycle

Transport medium

volumes, residence times, fluxes

Largest reservoir – oceans

τ = 40,000 yr

Compartments of the Hydrosphere


• Surface waters (watersheds) – streams, lakes reservoirs, wetlands, estuaries– Standing surface water vs. flowing surface water


• Temperature-density relationship leads to layering in lakes

Warmer water floats on colder = thermal stratification


• Groundwater – most from precipitation and infiltration– Composition depends on surrounding rock formations

(porosity and permeability)

Aquatic Chemistry

Aquatic Chemistry

• Algal photosynthesis:

Converts inorganic C (2HCO3

-) to organic form (CH2O, emp. formula for sugars)

CO32- is either converted back

to HCO3-, or ppts as

limestone

Biomass (CH2O) produced

Aquatic Chemistry

• Most redox reactions in water are catalyzed by bacteria

– e.g. N compounds to NH4+ in anoxic conditions

– e.g. N to NO3- 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.

CaCO3 ⇌ Ca2+ + CO32-

CO32- + H2O ⇌ HCO3

- + 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 reactionCO3

2- + H+ ⇌ HCO3-

Bicarbonate neutralization reactionHCO3

- + H+ ⇌ H2O.CO2 ⇌ H2O + CO2

Hydroxide neutralization reactionH+ + OH- ⇌ H2O

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

2-] – [H+]

• Units are mg L-1 CaCO3 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 H3O+

• ‘Master Variable’ – controls parameters e.g. speciation

• Ranges 5.5 - 9

Acidity

• Acidity results from presence of weak acids: H2PO4-, CO2, H2S,

proteins, fatty acids, metal ions (e.g. Al3+, Fe3+)

e.g. [Al(H2O)6]3+ + H2O ⇌ [Al(H2O)5OH]2+ + H3O+

simplifies as [Al(H2O)6]3+ ⇌ [Al(H2O)5OH]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 WaterMetal Ions

• Mn+ 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(H2O)x]n+ coordinate bonded to water molecules or other bases (e- donors)


• 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.


• Allows for loss of H+, reactions:

Acid base: [Fe(H2O)6]3+ ⇌ [FeOH(H2O)5]2+ + H+

Ppt: [Fe(H2O)6]3+ Fe(OH)⇌ 3(s) + 3H2O + 3H+ (results from A-B)

Redox: [Fe(H2O)6]2+ Fe(OH)⇌ 3(s) + 3H2O + e- + 3H+

• Due to these reactions conc. of the hydrated cation, e.g. [Fe(H2O)6]3+ is very small


• Acid-base reaction is more completely shown by H+ ion is being pulled off by a water molecule in soln:

[Fe(H2O)6]3+ + H2O [FeOH(H⇌ 2O)5]2+ + H3O+

• Successive deprotonations:

[FeOH(H2O)5]2+ +H2O ⇌ [Fe(OH)2(H2O)4]+ + H3O+

[Fe(OH)2(H2O)4]+ +H2O [⇌ Fe(OH)3(H2O)3](s) + H3O+

• Forms a neutral complex which does not dissolve and precipitates, Fe(OH)3


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(H2O)6]3+ ⇌ [FeOH(H2O)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(H2O)6]3+ Fe(OH)⇌ 3(s) + 3H2O + 3H+


• Properties of metals dissolved in water depend upon the nature of metal species dissolved in water, called speciation

• In addition to hydrated [M(H2O)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 WaterCalcium and Harness

• Ca2+ generally has highest conc. and most influence on aquatic chemistry

• Why?

• Calcium is a key element in many geochemical processes

• Primary minerals: gypsum (CaSO4.2H2O), anhydrite (CaSO4), dolomite (CaMg(CO3)2, calcite and aragonite (CaCO3)


CO2(g) + H2O(aq) ⇌ H2CO3(aq) KH

H2CO3(aq) ⇌ H+ + HCO3- Ka

CaCO3(s) ⇌ Ca2+ + CO32- Ksp

CO32- + H2O ⇌ HCO3

- + OH- Kb

H+ + OH- ⇌ H2O 1/Kw

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

• Giant titration of acid from atmospheric CO2 with base from carbonate ion in rocks

K = KspKbKHKa/Kw = 1.5 x 10-6 = [Ca2+][HCO3-]2

PCO2


• CaCO3(s) + CO2(g) + H2O(aq) ⇌ Ca2+ + 2HCO3-

K = KspKbKHKa/Kw = 1.5 x 10-6 = [Ca2+][HCO3-]2

PCO2

If [Ca2+] = S, [HCO3-] = 2S

1.5 x 10-6 = [Ca2+][HCO3-]2 = S (2S)2

PCO2 0.00037 atm

S = [CO2] = 5.2 x 10-4 mol L-1 (34 x amount calculated from Henry’s law)

S = [Ca2+] = 5.2 x 10-4 mol L-1 (this is 4x closed system)

[HCO3-] = 2S = 1.0 x 10-3 mol L-1

Acid-base reaction increases the solubility of both the gas and he solid – water that contains CO2 more readily dissolves calcium carbonate


• CO32-, H+ and OH- can be derived

Ksp = [Ca2+][CO32-] [CO3

2-] = 8.8 x 10-6 mol L-1

Kb = [HCO3-][OH-] [OH-] = 1.8 x 10-6 mol L-1

[CO32-]

Kw = [H+][OH-] [H+] = 5.6 x 10-9 mol L-1

Conclude natural water at 25 °C with a pH determined by saturation with CO2 and CaCO3 should be alkaline (pH = 8.3)

Actual value of calcereous waters is around 7-9 …why the difference?


• Simple model does not include CO2 from respiration of MO’s!

• MO’s directly affect conc. of Ca2+ in water


Ca2+ Mg2+

Fe2+

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


Hard water contains high concentrations of dissolved calcium and magnesium ions

Soft water contains few of these dissolved ions.

Hardness = [Ca2+] + [Mg2+]

(also Al3+, Fe3+, Mn2+ and Zn2+)

Carbonate minerals: limestone - CaCO3 dolomite - CaCO3.MgCO3 sulfates - CaSO4

Counter ions of alkalinity ions

Alkalinity is a good indicator of hardness and vice-versa


• Deposition of white solid CaCO3 or MgCO3 when water is heated

– ‘furring-up blocks pipes and lowers efficiency of industrial processes

• Formation of scum (insoluble ppt) with soap and water

Ca2+(aq) + 2Na(C17H33COO-

)(aq) 2Na+ + Ca(C17H33COO- )2(s)

– detergent action is blocked

• Staining (due to transition metals)

A pipe with hard-water scale build up


• Solid deposit = carbonate hardness or temporary hardness

Ca2+ + 2HCO3- ⇌ CaCO3(s) + CO2(g) + H2O(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 O2 (atmospheric)

Acidic solutionO2 + 4H+ + 4e- ⇌ 2H2O

Basic solutionO2 + 2H2O + 4e- ⇌ 4OH−

Concentration of O2 in water is low (10 ppm average), governed by Henry’s law:

O2(g) ⇌ O2(aq)

KH = [O2 (aq)]

PO2

At 25 °C, KH = 1.3 x10-3 mol L-1 atm-1

O2 is reduced from 0 to -2 state in H2O or OH−

Dissolved O2 influences chemical speciation of elements in natural and polluted waters

Oxidation-Reduction

• Show that (a) O2 + 2H2O + 4e- ⇌ 4OH− (from above)

Is the same as (b) 2H2O + 2e- ⇌ H2(g) + 2OH- (from Manahan)

Double (b):

2(2H2O + 2e- ⇌ H2(g) + 2OH-

4H2O + 4e- 4⇌ H2(g) + 4OH-

Add O2 + 2H2 ⇌ 2H2O

O2 + 2H2O + 4e- ⇌ 4OH−

Question

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

At 25 °C, KH = 1.3 x10-3 mol L-1 atm-1

KH = [O2 (aq)] / PO2

[O2 (aq)] = KH x PO2

[O2 (aq)] = (1.3 x10-3 mol L-1 atm-1 ) x 0.21 atm = 2.7 x 10-4 mol L-1

[O2 (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 O2

• Temperature (inc)

• Pressure (dec)

• Salts (inc)

• Organic matter (inc)

Dissolved O2 decreases with increasing temperature

Oxidation-ReductionOxygen Demand

• The most common substance oxidized by DO in water is organic matter (plant debris, dead animals etc.)

CH2O(aq) + O2(aq) → CO2(g) + H2O(aq)

• Similarly DO is consumed by NH3 and NH4+ in the nitrification process

• Water in streams and rivers are constantly replenished with oxygen• Stagnant water and deep lakes can have depleted oxygen

0 to +4

0 to -2

Oxidation-ReductionOxygen Demand

Half reactions

Oxidation: CH2O + H2O → CO2 + 4e- + 4H+

Reduction: 4H+ + O2 + 4e- → 2H2O

CH2O(aq) + O2(aq) → CO2(g) + H2O(aq)

In basic conditions?

O2 + 4H+ + 4e- 2 H2O

React with hydroxide O2 + 4H+ + 4OH- + 4e- 2H2O + 4OH-

O2 + 4H2O + 4e- 2H2O + 4OH-

O2 + 2H2O + 4e- 4OH- Same overall

Question

P9-4: Determine the balanced redox reaction for the oxidation of ammonia to nitrate ion by O2 in alkaline solution (basic)

Does this reaction make the water more basic or less?

NH3 + O2 NO3- + H2O

Using standard redox balancing techniques:

NH3 + 2O2 + OH- NO3- + 2H2O

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-ReductionBiological 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 O2 in the stream or lake. Take a sample and store at 25oC for five days and remeasure O2 content. The difference is the BOD

– BOD5 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 – significantly lower than the solubility of O2 in water (8.7 mg L-1)

– Sewage has BOD of ~100 mg L-1

Oxidation-ReductionChemical Oxygen Demand

O2 + 4H+ + 4e- → 2H2O

• Dichromate ion, Cr2O72- dissolved in sulfuric acid is a powerful oxidizing agent. It

is used as an oxidant to determine COD

Cr2O72- + 14H+ + 6e- → 2Cr3+ + 7 H2O

• Excess dichromate is added to achieve complete oxidationBack titration with Fe2+ gives the desired endpoint value

# moles of O2 consumed = 6/4 x (#moles Cr2O7 consumed)

Note: Cr2O72- is a powerful oxidizing agent and can oxidize species that are

not usually oxidized by O2 - hence gives an upper limit

Question

P9-5: A 25 mL sample of river water was titrated with 0.0010 M Na2Cr2O7 and required 8.7 mL to reach the endpoint. What is the COD (mg O2/L)?

No. moles Cr2O72- = 0.0010 mol L-1 x (8.7 x 10-3 L) = 8.7 x 10-6 mols

No. moles O2 = 1.5 moles Cr2O72- = 1.5 x (8.7 x 10-6 mols)

= 1.3 x 10-5 mols O2

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-ReductionThe pE Scale

• Oxidation and reduction are controlled by the concentrations of electrons which are present:

pE = - log10[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 = E2.303 RT/F


• When a significant amount of O2 is dissolved, the reduction of O2 is the dominant reaction determining e- availability:

¼ O2 + H+ + e- ½ H⇌ 2O

• 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+] PO2¼)

OR

pE = 20.75 – pH + ¼ log(PO2 )


A convenient approach is to use Nernst Equation of electrochemistry

E = E0 – (RT/F) (log [products] / [reactants])…for 1 electron redox process

E = E0 - 0.0591(log [products] / [reactants])where E0 is the standard electrode potential for a one electron reduction

One can equate pE to the Electrode Potential E

pE = E/0.0591 or pE0= E0/0.0591

• Dividing throughout by 0.0591:

pE = pE0 - (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 = pE0 - (log [products] / [reactants]) to an equation relating pE to pH

Redox Chemistry in Natural Waters The pE Scale

¼ O2 + H+ + e- ½ H⇌ 2O

pE = pE0 - (log [products] / [reactants]) pE = 20.75 - log 1/ [reactants] = 20.75 + log ([reactants])

pE = 20.75 + log([H+] PO2¼) = = 20.75 + log([H+] + log (PO2

¼)

pE = 20.75 – pH + ¼ log(PO2 )

pE = 20.75 – pH + ¼ log(PO2 )

For a neutral sample of water that is saturated with oxygen from air (PO2 = 0.21 atm) that is free from CO2 (pH = 7) the pE value corresponds to 13.9

…pE value decreases with decrease in O2 and increase in pH

E0 = 1.23 VpE0 = 1.23/0.0591

Dominant redox equilibrium reaction determines pE of water (O2 may not be dominant redox species!)

Question

What is the most oxidizing conditions possible in water?

PO2 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 waterwill be oxidized

A similar analysis gives boundary below which water will be reduced

pE = 20.75 - pH

pH

pE

pE = - pH

Redox Chemistry in Natural Waters The pE Scale

Example

1/8NO3− + 5/4H+ + e- ⇌ 1/8NH4

+ + 3/8H2O E0 = +0.836 V

pE0 = E0/0.0591 = 0.836/0.0591 = +14.15

pE = pE0 – log [NH4+]1/8

[NO3-]1/8[H+]5/4)

= 14.15 - 5/4pH -1/8log([NH4+]/[NO3

-])

Note: Express the reactions as one electron reduction process….. Follow the examples given on page 435

ax = x log alog(1/b) = -log b

Question

9-7: Deduce the equilibrium ratio of concentrations of NH4+ to NO3

- 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([NH4+] / [NO3

-])

11 = 14.15 – (5/4) x 6 – (1/8)log([NH4+] / [NO3

-])

log([NH4+] / [NO3

-]) = -8(4.35) = -34.8

[NH4+] / [NO3

-] = 1.6 x 10-35

pE = 14.15 – (5/4)pH – (1/8)log([NH4+] / [NO3

-])

-3 = 14.15 – (5/4) x 6 – (1/8)log([NH4+] / [NO3

-])

log([NH4+] / [NO3

-]) = 8(9.65) = 77.2

[NH4+] / [NO3

-] = 1.6 x 1077

Problem 9-7

pH = 6, pE = 11,

pH = 6, pE = -3,

Redox Chemistry in Natural Waters The pE-pH Diagram

Fe3+ + e- ⇌ Fe2+

• For this reaction, pE0 = 13.2• pE = 13.2 + log([Fe3+] / [Fe2+]) NOT pH DEPENDENT!

e.g. Ratio of Fe3+ to Fe2+ when pE = -4.1 (reducing)

-4.1 = 13.2 + log([Fe3+] / [Fe2+])

log([Fe3+] / [Fe2+]) = -17.3

[Fe3+] / [Fe2+] = 5 x 10-18

(far more Fe2+)

• Transition between dominance of one form over the other occurs at [Fe3+] = [Fe2+], pE = 13.2 + log(1) = 13.2 + 0 = 13.2

Redox Chemistry in Natural Waters pE – pH Stability Field Diagrams

• Fe3+ ion is stable in oxidizing acidic conditions, Insoluble Fe(OH)3 is predominant iron species

• Fe2+/Fe3+ ions can only exist under acidic conditions

• At higher pH Fe3+ is present as Fe(OH)3. Fe(OH)2 does not precipitate until solution becomes significantly basic

• Changes in redox conditions govern whether the iron will be in solution or in the sediments

Zone dominance of various oxidation states

pH independent

pE independent

pE independent

Complexation and Chelation


• Mn+ exists in various forms in water

• Exist as hydrated cations [M(H2O)x]n+ coordinate bonded to water molecules or other bases (e- donors) called ligands

• Ligands - bond to a metal ion to form a complex ion (coordination compound)

e.g. Cd2+ + CN- ⇌ [CdCN]+

[CdCN]+ + CN- ⇌ Cd(CN)2

Cd(CN)2 + CN- ⇌ [Cd(CN)3]-

[Cd(CN)3]- + CN- ⇌ Cd(CN)42-

(CN- is unidentate ligand)


• 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-


• 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 Ni2+ (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 ChelationOccurrence and Importance

• Chelating agents are common potential pollutants

• Occur in sewage and industrial wastes

e.g. EDTA (ethylenediaminetetraacetic acid)

• Tend to solubilize heavy metals from plumbling and from waste deposits

+ Mn+

Complexation and ChelationComplexation 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 ChelationComplexation by Humic Substances

• Binding of metal ions by humic substances:

Complexation and ChelationComplexation by Organometallic Compounds

• Organometallic compounds – metal attaches to organic ligand

Hg2+ (Mercury (II) ion)

CH3Hg+ (Monomethylmercury ion)

(CH3)2Hg (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


• 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


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 CO2 via photosynthesis

CO2 + H2O → {CH2O} + O2(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)

hv

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


• Biogeochemical cycles – microbially mediated transitions between elemental species


Carbon Cycle• Small amount is

atmospheric CO2

• Large amount present in minerals (carbonates)

• Organic fraction as hydrocarbons

• Manufacture of toxic xenobiotic compounds from hydrocarbons


Carbon Cycle – Involvement of MO’s

• Photosynthesis – algae, higher plants, bacteria use light energy to fix inorganic C

CO2 + H2O → {CH2O} + O2(g)

• Respiration:

Aerobic respiration – OM is oxidized

{CH2O} + O2(g) → CO2 + H2O

Anaerobic respiration – uses oxidants other than O2, NO3- or SO4

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{CH2O} → CH4 + CO2

• Bacterial utilization and degradation of HC’s – oxidation of HC’s

• Biodegradation of organic matter – treatment of wastewater


Nitrogen Cycle

• N is interchanged among the atmosphere, OM, and inorganic compounds

• MO’s mediate reactions


• Nitrogen fixation – binding of atmospheric N2

3{CH2O} + 2N2 + 3H2O + 4H+ → 3CO2 + 4NH4+

• Nitrification – converts ammonium to nitrate

2O2 + NH4+ → NO3

- + 2H+ + H2O

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

• Denitrification – produces N2, N2O or NO, returns to atmosphere


• 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)

ch11 water and the hydrosphere

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