Transcript Environmental Chemistry - Robert Morris University

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
Partial charges result
from bond polarization
A difference in the electronegativities
of the atoms in a bond creates a
polar bond
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
• 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 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).
The Hydrosphere
Compartments
• Atmosphere
• Land
• Groundwater
• Rivers lakes
• Oceans
Hydrologic Cycle
Source: http://www.nasa.gov/vision/earth/environment/warm_wetworld.html
Where Does Potable
(fit for consumption) Drinking Water Come From?
Sources
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
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.
Sources
< 1000 m3 per person per year
Engelman et al., 2000
Access to Water
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
518
303,977
754
423
80
5
4,850
512,440
1,720
14
27
58
Middle East/N.
Africa
N. America
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)
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,
residence
times, Water Cycle
fluxes
Liquid Medium
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
(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+ + CO32CO32- + 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 reaction
CO32- + H+ ⇌ HCO3Bicarbonate neutralization reaction
HCO3- + H+ ⇌ H2O.CO2 ⇌ H2O + CO2
Hydroxide neutralization reaction
H+ + OH- ⇌ H2O
Alkalinity = [OH-] + [HCO3-] + 2[CO32-] – [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 Water
Metal 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)
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:
[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
Metal Ions and Calcium in Water
Metal Ions
•
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(H2O)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
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(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+
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(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 Water
Calcium 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)
Metal Ions and Calcium in Water
Calcium and Harness
CO2(g) + H2O(aq) ⇌ H2CO3(aq)
H2CO3(aq) ⇌ H+ + HCO3CaCO3(s) ⇌ Ca2+ + CO32CO32- + H2O ⇌ HCO3- + OHH+ + OH- ⇌ H2O
KH
Ka
Ksp
Kb
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
Metal Ions and Calcium in Water
Calcium and Harness
•
CaCO3(s) + CO2(g) + H2O(aq) ⇌ Ca2+ + 2HCO3K = 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
Metal Ions and Calcium in Water
Calcium and Harness
•
CO32-, H+ and OH- can be derived
Ksp = [Ca2+][CO32-]

[CO32-] = 8.8 x 10-6 mol L-1
Kb = [HCO3-][OH-]
[CO32-]

[OH-] = 1.8 x 10-6 mol L-1
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?
Metal Ions and Calcium in Water
Calcium and Harness
• Simple model does not include CO2 from respiration of MO’s!
• MO’s directly affect conc. of Ca2+ in water
Metal Ions and Calcium in Water
Calcium and Harness
Ca2+
Mg2+
Fe2+
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.
Counter ions of
alkalinity ions
Hardness = [Ca2+] + [Mg2+]
Carbonate minerals:
limestone - CaCO3
dolomite - CaCO3.MgCO3
sulfates - CaSO4
Alkalinity is a
good indicator
of hardness and
vice-versa
(also Al3+, Fe3+, Mn2+ and Zn2+)
Metal Ions and Calcium in Water
Calcium and Harness
•
•
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
Metal Ions and Calcium in Water
Calcium and Harness
•
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 solution
O2 + 4H+ + 4e- ⇌ 2H2O
O2 is reduced from 0 to -2
state in H2O or OH−
Basic solution
O2 + 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
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) + 2OH4H2O + 4e- ⇌ 4H2(g) + 4OHAdd 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-Reduction
Oxygen Demand
•
The most common substance oxidized by DO in water is organic matter
(plant debris, dead animals etc.)
0 to -2
CH2O(aq) + O2(aq) → CO2(g) + H2O(aq)
0 to +4
•
•
•
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
Oxidation-Reduction
Oxygen Demand
Half reactions
Oxidation:
Reduction:
CH2O + H2O → CO2 + 4e- + 4H+
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 + 4OHO2 + 4H2O + 4e-  2H2O + 4OHO2 + 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-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 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-Reduction
Chemical 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 oxidation
Back 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-Reduction
The 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 =
E
2.303 RT/F
Oxidation-Reduction
The pE Scale
•
When a significant amount of O2 is dissolved, the reduction of O2 is the
dominant reaction determining e- availability:
¼ O2 + H+ + e- ⇌ ½ H2O
•
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 )
Oxidation-Reduction
The pE Scale
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 H2O
E0 = 1.23 V
pE0 = 1.23/0.0591
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 (PO2¼)
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
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
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
pE
A similar analysis gives boundary
below which water will be reduced
pE = - pH
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)
ax = x log a
log(1/b) = -log b
= 14.15 - 5/4pH -1/8log([NH4+]/[NO3-])
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 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
Zone dominance of various oxidation states
pE independent
pH
independent
pE independent
•
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
Complexation and Chelation
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)
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 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 Chelation
Occurrence and Importance
•
•
Chelating agents are common potential pollutants
Occur in sewage and industrial wastes
e.g. EDTA (ethylenediaminetetraacetic acid)
+ Mn+
•
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
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
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 CO2 via
photosynthesis
hv
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)
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 CO2
• 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
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 SO42-
•
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
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 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
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)