Chapter 14: Aquatic Life and Oxidation/Reduction Preview

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Chapter 14: Aquatic Life and Oxidation/Reduction
Preview
We now turn to water as the medium that supports life. All organisms require
water, and a large fraction of them make their home in rivers, lakes, and the oceans.
Life started in the ocean and occupied dry land only later. Moreover, biological processes
have a profound influence on the chemistry of natural waters, and indeed of the entire
globe. Were it not for the evolution of photosynthetic organisms, first in the ocean, and
then on land, the atmosphere would be devoid of oxygen. The profound influence of
oxygen on the chemistry of the atmosphere was considered at length in Part II. O2 is
also the dominant actor in the chemistry and biochemistry of the hydrosphere. The
limited availability of O2 in water sets the boundary between aerobic and anaerobic life,
with crucial consequences for water quality and the health of ecosystems. In this chapter
we consider
• Redox energy and dissolved oxygen
• Biological redox and the reduction potential
• Linkage of redox with acid/base chemistry
• Earth’s redox evolution
• Biological CO2 pump
• Overfertilization of surface waters: eutrophication
• Redox and metal pollution
• Ocean fertilization with iron
14.1 Redox Reactions and Energy
Life is powered by redox reactions, chemical processes in which electrons are
transferred from one molecule to another, with the release of energy. Organisms have
evolved machinery, made up of proteins and membranes, which channels this energy into
the biochemical pathways that support vital functions.
In an aerobic environment, the most important biological redox process is
respiraton,
(CH2O)
+
O2
=
CO2
+
H2O
[14‐1]
which we encountered previously as part of the global carbon cycle [p.??]. In this case
carbohydrate molecules provide electrons for the reduction of dioxygen. All higher life
forms obtain their energy via respiration. However, many other redox processes are
utilized by bacteria. Indeed, bacteria have evolved to exploit just about any redox
process that is available in nature. Anyplace where a supply of oxidizable molecules
coexist with molecules capable of oxidizing them, it is a good bet that bacteria are
present which can utilize the potential redox reaction. The oxidation of FeS2 by
thiobacillus ferrooxidans in the above discussion of acid mine drainage is a good example.
1
14.2 Biological oxygen demand
Wherever oxygen is present, respiration provides life-supporting redox energy,
but in liquid water oxygen can easily become depleted. The solubility of O2 in water is
only 9 mg/L (about 0.3 millimoles) at 20 °C, and less at higher temperatures. The
oxygen supply can be replenished by contact with the air, as in rapidly flowing streams.
But in standing water or in waterlogged soils, the diffusion of oxygen from the
atmosphere is slow relative to the speed of microbial metabolism, and the oxygen is
used up.
Given the centrality of oxygen to metabolism, a parameter called biological
oxygen demand (BOD) has been defined to measure the reducing power of water
containing organic carbon. BOD is the number of milligrams of O2 required to carry out
the oxidation of organic carbon in one liter of water. Values for various industrial wastes
and municipal sewage are given in Table 14.1.
Worked
Problem
14.1:
BOD
Q:
What
is
the
BOD
of
water
in
which
10
mg
of
sugar
[empirical
formula
CH2O]
is
dissolved
in
a
liter?
How
does
this
compare
with
the
O2
solubility
at
20
oC?
A:
Since
each
mole
of
CH2O
requires
one
mole
of
O2
[equation
14.1],
we
divide
10
mg
by
the
molecular
weight
of
CH2O
[30
g],
to
obtain
the
required
number
of
moles
of
O2
[32
g/mol],
and
then
multiply
by
the
molecular
weight
of
O2
to
obtain
the
number
of
mg:
BOD
=
10
mg
x
32
g/30g
=
10.7
mg/l
This
exceeds
the
O2
solubility
[9
mg/l]
by
about
20%.
14S.1
Oxidation
Levels
and
Water
Many
elements
can
exist
in
multiple
oxidation
states,
depending
on
the
number
of
electrons
added
to
or
removed
from
the
valence
shell
of
the
atoms.
In
an
aqueous
world,
the
stability
of
these
different
oxidation
states
depends
on
the
properties
of
water.
Thus
we
are
familiar
with
Na+,
and
Mg2+
ions,
because
sodium
and
magnesium
have
one
and
two
electrons,
respectively,
in
their
valence
shells,
which
are
easily
removed
when
water
molecules
are
available
to
stabilize
the
2
resulting
ions
[Figure
12.5].
All
metals
form
positive
ions
in
water,
and
in
the
case
of
transition
metals
multiple
oxidation
states
are
available;
for
example,
iron
can
exist
in
water
as
Fe3+
or
Fe2+.
Nonmetals,
being
electronegative
elements,
readily
attain
negative
oxidation
levels,
depending
on
the
number
of
electrons
that
their
valence
shells
can
accommodate.
Thus
the
lowest
oxidation
levels
attainable
by
F,
O,
N
and
C
are
–I,
‐II,
‐III,
‐IV
respectively;
we
use
Roman
numerals
to
denote
the
oxidation
number
to
distinguish
them
from
the
actual
charge.
Thus,
although
Cl‐
ions
exist
as
such
in
water,
O2‐
ions
do
not.
Their
proton
affinities
are
high
enough
that
they
are
completely
converted
to
OH‐
[or
to
H2O,
depending
on
the
pH].
The
lowest
oxidation
levels
for
N
and
C,
are
represented
by
NH3
[or
NH4+]
and
CH4.
Positive
oxidation
levels
are
also
accessible
to
the
non‐metals
because
of
the
stabilization
available
through
bonding
to
oxide
ions.
Thus
C,
N,
S
and
Cl
are
in
their
maximum
oxidation
states,
+IV,
+V,
+VI
and
+VII
respectively
when
surrounded
by
oxide:
CO2
[or
CO32‐],
NO32‐,
SO42‐,
and
ClO4‐.
The
actual
charges
on
the
central
atoms
in
these
molecules
are
much
less
than
+4,
+5,
+6
or
+7,
since
electrons
are
shared
in
the
polar
but
covalent
bonds
with
the
O
atoms.
Nevertheless,
the
oxidation
state
is
crucial
in
determining
the
possibilities
for
redox
chemistry.
For
example,
eight
electrons
must
be
removed
from
N
in
order
to
convert
NH3
to
NO32‐.
In
the
case
of
the
respiration
reaction,
[14.1],
carbon
in
(CH2O)
is
in
the
oxidation
state
O
[the
rules
are
that
O
counts
for
–2,
and
H
counts
for
+1
in
determining
the
‘effective’
charge,
i.e.
the
oxidation
state,
of
the
remaining
atoms];
four
electrons
are
transferred
to
O2
in
converting
(CH2O)
to
CO2.
Worked
Problem
14.2:
Calculating
the
Oxidation
State
and
Balancing
Redox
Equations
Q:
What
is
the
oxidation
state
of
N
in
the
nitrite
ion,
NO2‐?
A:
Since
O
counts
as
–2,
and
there
is
an
overall
–1
charge,
N
must
have
an
effective
charge
of
+3.
The
oxidation
level
is
III.
Q:
Write
a
balanced
chemical
equation
for
the
reduction
of
NO2‐
to
NH3
by
H2.
A:
First
balance
the
number
of
electrons
transferred
from
oxidant
to
reductant.
Since
N
changes
from
III
to
–III,
six
electrons
are
transferred.
H
changes
from
O
to
I,
so
six
H
atoms,
or
three
H2
molecules,
are
required
to
receive
the
electrons.
NO2‐
+
3H2
=
NH3
Since
the
reaction
is
in
water,
it
is
permissible
to
add
H2O
or
H+
or
OH‐
to
either
side
of
the
reaction,
as
needed.
Seeing
that
nitrite
had
two
O
atoms,
we
balance
these
by
adding
two
water
molecules
to
the
right
hand
side.
NO2‐
+
3H2
=
NH3
+
2H2O
The
total
H
count
on
the
right
hand
side
is
now
seven,
which
we
balance
by
adding
one
H+
to
the
left
hand
side.
This
also
balances
the
charge.
NO2‐
+
3H2
+
H+
=
NH3
+
2H2O
3
14.3 Natural sequence of biological reduction
When water is depleted of oxygen, organisms that depend upon aerobic
respiration cannot survive, and anaerobic bacteria take over. These bacteria utilize
oxidants other than O2. These alternative oxidants are less powerful than O2, and cannot
produce as much energy. Nevertheless, bacteria are quite capable of surviving on lower
energy processes; in doing so, they can fill ecological niches that are not available to
aerobic organisms. The oxidizing power of anaerobic environments in the biosphere is
mainly controlled by five molecules. In decreasing order of energy produced, they are
nitrate (NO3–), manganese dioxide (MnO2), ferric hydroxide (Fe(OH)3), sulfate (SO42–) and,
under extreme conditions, carbohydrate (CH2O) itself. The biological oxidation processes
supported by these oxidants are described in Table 14.2. Microbial populations first
use the oxidant that produces the most energy until it is depleted; only then does
another agent become the dominant oxidant.
Table
14.2
Redox
Reactions,
Products,
and
Consequences
Redox
reaction
Reaction
products/consequences
1.
O2
+
CH2O
→
CO2
+
H2O
The
aerobic
condition,
characterized
by
the
highest
redox
potential,
occurs
when
there
is
an
abundance
of
O2,
and
the
relative
absence
of
organic
matter
owing
to
oxic
sewage
wastes,
and
the
decomposition
of
organic
matter
near
the
surface
of
well‐aerated
soils.
The
end
products,
CO2
and
water,
are
nontoxic.
2.
4NO3‐
+
5CH2O
+
4H+
When
molecular
oxygen
is
depleted
from
the
soil
or
→
5CO2
+
2N2
+
7H2O
water
column,
as
would
be
the
case,
for
example,
in
waterlogged
soils
and
wetlands,
available
nitrate
is
the
most
efficient
oxidant.
Denitrifying
bacteria
consume
nitrate
and
release
N2.
N2O,
a
greenhouse
gas,
is
also
released
as
a
side‐product.
In
agricultural
soils,
denitrification
can
lead
to
losses
of
nitrogen
fertilizer
amounting
to
as
much
as
20%
of
inputs.
Denitrifying
bacteria
are
also
very
active
in
heavily
polluted
rivers
or
in
stratified
estuaries
where
organic
matter
accumulates.
In
some
estuary
systems,
denitrification
may
significantly
affect
the
transfer
of
nitrogen
to
the
adjacent
coastal
waters
and
atmosphere.
+
3a.
2MnO2
+
CH2O
+
4H In
aerobic
environments
where
nitrates
are
in
low
→
2Mn2+
+
3H2O
+
CO2
concentration
and
manganese
and
ferric
oxides
are
abundant,
the
metal
oxides
are
a
source
of
oxidant
3b.
4Fe(OH)3
+
CH2O
+
8H+
for
microbial
oxidation.
This
may
be
the
case
in
2+
→
4Fe +
11H2O
+
CO2
natural
soils,
and
in
the
sediments
of
lakes
and
rivers.
The
environmental
significance
of
these
metal
oxides
is
that
they
serve
a
dual
role.
Not
only
are
they
a
source
of
oxidants
to
microorganisms,
they
are
also
important
for
their
capacity
to
bind
4
4a.
½
SO42‐
+
CH2O
+
H+
→
½
H2S
+
H2O
+
CO2
4b.
MS2
+
7
O2
+
H2O
→
M2+
+
2SO42‐
+
2H+
5.
CH2O
+
CH2O
→
CH4
+
CO2
toxic
heavy
metals,
deleterious
organic
compounds,
phosphates,
and
gases.
When
the
metal
oxides
are
reduced,
they
become
water‐soluble
and
lose
their
binding
ability.
This
loss
may
result
in
the
release
of
toxic
materials.
Sulfidic
conditions
are
brought
about
almost
entirely
by
the
bacterial
reduction
of
sulfate
to
H2S
and
HS‐
accompanying
organic
matter
decomposition.
Sulfate
reduction
is
very
common
in
marine
sediments
because
of
the
ubiquity
of
organic
matter
and
the
abundance
of
dissolved
sulfate
in
seawater.
In
freshwater,
such
reactions
are
important
in
areas
affected
by
acidic
deposition
in
the
form
of
sulfuric
acid.
H2S
is
an
extremely
toxic
gas.
Sulfides
are
also
important
in
scavenging
heavy
metals
in
bottom
sediments.
Conversion
of
a
heavy‐metal
sulfide
(MS2)
to
sulfate
may
also
occur
when
anaerobic
sediments
are
exposed
to
the
atmosphere,
as
in
the
case
of
the
raising
of
dredge
soils.
It
may
also
occur
when
wetlands
containing
pyrites
(FeS2)
are
drained
for
agriculture
or
in
coal‐mining
areas
as
acid
mind
drainage.
One
consequence
may
be
an
increase
in
acidification
from
the
generation
of
sulfuric
acid;
another
might
be
the
release
of
toxic
metals.
Under
anaerobic
conditions
at
a
redox
potential
of
about
‐200
mV,
and
in
the
presence
of
methogenic
bacteria
as
may
be
found
in
swamps,
flooded
areas,
rice
paddies,
and
the
sediments
of
enclosed
bays
and
lakes,
partially
reduced
carbon
compounds
can
disproportionately
produce
methane
as
well
as
CO2.
This
reaction
is
more
typical
in
freshwater
systems
because
sulfate
concentrations
are
much
lower
than
in
marine
environments,
averaging
about
one
one‐hundreth
the
concentration
in
seawater.
Methane
is
a
critical
gas
in
the
determination
of
global
climate.
Since
the
early
1970s,
global
atmospheric
methane
levels
have
been
increasing
at
a
rate
of
1%
per
year.
Although
the
reasons
for
this
increase
are
still
under
investigation,
the
expansion
of
rice
paddy
cultivation
in
southeast
Asia
has
been
cited
as
a
contributing
cause.
See
chapter
3,
pp.
?.
5
Source:
W.M.
Stigliani
(1988).
Changes
in
valued
capacities
of
soils
and
sediments
as
indicators
of
nonlinear
and
time‐delayed
environmental
effects.
Environmental
Monitoring
and
Assessment
10:245‐307.
The oxidizing power of a molecule depends on the specific reaction being carried
out, and is measured as the reduction potential associated with the reduction of the
oxidant. These are listed in Table 14.3 for the environmental oxidants we are
considering. Microbial populations first use the oxidant that produces the most energy
until it is depleted; only then does another agent come the dominant oxidant. Thus, the
redox potential of a body of water tends to fall in a stepwise pattern as BOD increases
(Figure 14.1).
As oxidants are consumed in the conversion of reduced carbon to CO2, the
reduction potential falls to successively lower plateaus, corresponding to the
successively lower potential redox couples O2/H2O, NO3–/N2, MnO2/Mn2+, Fe(OH)3/Fe2+,
SO42–/HS–, and CO2/CH4 . These couples do not give reversible potentials at electrodes,
but the metabolic activity of the vast array of microbes in soils and in water ensure that
electron transfer does occur on a time-scale of hours or days (Table 14.3).
Consequently, all redox-active materials respond to the reduction potential established
by the microbial activity.
Note, however, that, while there is a general correspondence with the Eh values
of the half-reactions, the plateau potentials in Figure 14.1 deviate substantially from the
numbers listed in Table 14.3. This is because conditions in the environment are far
from the standard conditions which establish the Eh values. While the pH may be close
to 7, the concentrations of other reactants and products are unlikely to be 1.0 M [or 1
atm, in the case of a gas].
6
Figure
14.1
Sequence
of
redox
reactions
in
aqueous
environments.
O2
in
natural
waters
at
20
oC
is
sufficient
to
oxidize
about
3.4
mg
of
organic
carbon
(shown
here
as
CH2O)
per
liter
of
water.
When
the
rate
of
replenishment
of
O2
from
the
atmosphere
is
slower
than
the
rate
of
oxidation
of
CH2O,
oxygen
is
depleted
and
microbes
will
select
the
next
most
energetic
oxidant
in
the
sequence
shown.
For
simplicity,
only
major
products
and
their
valence
states
are
shown.
See
Table
14.2
for
balanced
equations.
Source:
W.M.
Stigliani
(1988).
Changes
in
valued
capacities
of
soils
and
sediments
as
indicators
of
nonlinear
and
time‐delayed
environmental
effects.
Environmental
Monitoring
and
Assessment
10:245‐307.
14S.2
Reduction
Potentials
All
redox
reactions
can
be
divided,
at
least
conceptually,
into
two
reduction
half­reactions,
one
proceeding
forward
and
the
other
in
reverse.
For
example,
the
oxidation
of
hydrogen
by
oxygen,
2H2
+
O2
=
2H2O
[14‐2]
can
be
divided
into
O2
+
4e–
+
4H+
=
2H2O
[14‐3]
+
–
and
4H +
4e =
2H2
[14‐4]
Subtracting
half‐reaction
[14‐4]
from
[14‐3]gives
the
overall
reaction
[14‐2].
These
half‐reactions
can
actually
be
carried
out
at
the
electrodes
of
a
hydrogen‐
oxygen
fuel
cell,
as
discussed
in
Chapter
10
(pp.
??).
A
potential
difference
is
developed
between
the
oxygen
electrode
and
the
hydrogen
electrode,
allowing
a
current
to
flow
through
the
external
circuit.
For
the
hydrogen‐oxygen
fuel
cell,
this
potential
difference
approaches
1.24
volts
(V),
at
the
standard
temperature
of
25
oC,
when
the
gases
are
at
1
atmosphere
pressure,
and
the
electrodes
behave
reversibly,
7
that
is,
when
the
reactants
and
products
are
at
equilibrium
with
the
electrodes
(implying
rapid
electron
transfer
rates).
The
potential
difference,
ΔE,
is
the
energy
of
the
electrochemical
cell
per
unit
of
charge
delivered.
[Specifically
1
V
=
1
J/C,
where
V
=
volt,
J
=
joule
and
C
(coulomb)
is
the
unit
of
charge].
ΔE
is
related
to
the
free
energy
of
the
cell
reaction
by
the
relation
ΔG
=
–nFΔE
[I4‐5]
where
F
(the
Faraday)
is
the
amount
of
charge
in
a
mole
of
electrons,
[96,500
C]
and
n
is
the
number
of
electrons
transferred
in
the
reaction.
Thus
in
reaction
[I4‐2],
4
electrons
are
transferred
from
2H2
to
O2,
and
ΔG
=
–4
e‐
x
96,500
C/mol
e‐
x
1.24
J/C
=
–479,000
J,
or
–479
kJ.
[Recall
that
this
value,
in
combination
with
the
entropy
of
the
reaction
gives
a
theoretical
energy
conversion
efficiency
of
80%
for
the
H2/O2
fuel
cell
–
P.
?]
Numerous
electrode
combinations
are
possible
in
electrochemical
cells,
and
it
is
convenient
to
specify
a
standard
potential,
Eo,
for
each
electrode
by
referencing
it
to
the
hydrogen
electrode,
whose
standard
potential
is
defined
as
zero.
Thus
Eo
=
1.24
V
for
the
oxygen
electrode,
represented
by
half‐reaction
[I4‐3].
The
standard
conditions
for
Eo
are
unit
activities
(partial
pressure
or
molar
concentration)
of
the
reactants
and
products,
at
25°C.
There
are
many
half‐reactions
whose
electrode
potential
cannot
actually
be
measured,
because
the
electron
transfer
reaction
at
an
electrode
is
too
slow.
These
potentials
can
nevertheless
be
calculated
from
the
free
energy
of
appropriate
redox
reactions.
For
example,
the
formation
of
NO
from
N2
and
O2,
whose
thermodynamics
was
considered
in
Chapter
4
(p.
??),
is
a
redox
reaction:
O2
+
N2
=
2NO
[14‐6]
which
can
be
divided
into
the
half‐reactions
and
O2
+
4e–
+
4H+=
2H2O
–
+
2NO
+
4e +
4H =
N2
+
2H2O
[14‐7]
[14‐8]
From
the
free
energy
of
the
overall
reaction
(p.
?),
173.4
kJ,
we
obtain
a
cell
potential
of
–0.45V
(using
equation
[14‐5]).
Then,
knowing
that
the
standard
potential
of
the
oxygen
electrode
is
1.24
V,
we
can
readily
calculate
that
the
standard
potential
for
half‐reaction
[14‐8]
is
1.69
V
(1.24
V
–
[‐0.45
V]),
even
though
it
is
impossible
to
measure
this
potential
directly
because
the
electron
transfer
between
the
electrode
and
the
NO
and
N2
molecules
is
too
slow
to
establish
a
reversible
potential.
14S.3
Concentration
Dependence
of
the
Potential:
pH
and
E0[w]
What
happens
to
the
reduction
potential
when
conditions
are
not
standard?
As
in
all
chemical
reactions,
the
driving
force
for
electrochemical
processes
depends
on
the
concentrations
of
reactants
and
products.
This
dependence
is
given
by
the
Nernst
equation:
8
E
=
Eo
–
[RT/nF]lnQ
[14‐9]
where
Eo
is
the
standard
potential,
R
is
the
gas
constant,
n
is
the
number
of
electrons
transferred
in
the
reaction,
and
Q
is
the
equilibrium
quotient,
i.e.
the
concentration
expression
for
the
equilibrium
constant.
In
the
fuel
cell
reaction
[14‐2],
for
example,
Q
=
1/PO PH 2
(the
water
activity
being
defined
as
unity),
and
n
=
4.
Therefore
E=1.24–[RT/4F]x[–lnPO –2lnPH ]
2
2
2
2
A
convenient
form
of
the
Nernst
equation
is
E=Eo–[0.059/n]logQ
[14‐10]
where
0.059
is
the
value
of
RT/F
at
25
°C,
multiplied
by
the
conversion
factor
from
natural
to
base‐ten
logarithms
[ln10
=
2.303].
For
temperatures
other
than
25
°C,
the
factor
0.059
must
be
raised
or
lowered
accordingly.
The
Nernst
equation
applies
equally
to
whole
cell
reactions
or
half‐reactions.
Thus
the
potential
of
the
hydrogen
electrode
(half‐reaction
[14‐4])
at
25
°C
is
(after
dividing
through
by
n
=
4]
E=0
–
0.059{log
PH /[H+]}
[14‐11]
From
this
we
see
that
the
hydrogen
electrode
potential
becomes
more
negative
as
[H+]
diminishes.
Thus
H2
gas
is
a
more
powerful
reductant
in
alkaline
solution
than
in
acid.
E
falls
by
–0.059
V
for
every
unit
rise
in
pH.
At
pH
7,
the
hydrogen
electrode
potential
is
–0.42
[when
all
other
conditions
are
standard].
Likewise
O2
is
a
less
powerful
oxidant
in
alkali
than
in
acid,
because
protons
are
consumed
in
the
reduction
half‐reaction,
[14‐3].
The
oxygen
potential
(again
after
dividing
by
n
=
4)
is
E
=
1.24
–
0.059log{1/PO21/4[H+]}
[14‐12]
Again
the
potential
drops
0.059
V
for
every
unit
rise
in
pH
and
is
0.82
V
at
pH
7.
Because
pH
7
is
closer
to
most
biologically
and
environmentally
relevant
conditions
than
is
pH
0,
electrode
potentials
are
often
cited
for
pH
7,
as
they
are
in
Table
14.3.
The
Eo[w]
values
are
Eo
values
recalculated
for
pH
7.
Even
if
no
protons
appear
explicitly
in
a
half‐reaction,
the
potential
may
be
pH‐dependent
because
of
secondary
acid‐base
reactions.
For
example,
the
potential
of
the
Fe3 + reduction
half‐reaction
Fe3+
+
e‐
=
Fe2+
[14‐13]
has
no
proton
dependence
per
se,
but
the
equilibrium
quotient,
[Fe2+]/[Fe3+],
is
highly
dependent
on
pH
because
of
the
acidic
character
of
Fe3+.
At
quite
low
pH,
it
forms
a
series
of
hydroxide
complexes,
and
precipitates
as
the
highly
insoluble
Fe(OH)3
(Ksp
=
10‐37).
In
contrast,
Fe2+
forms
hydroxide
complexes
only
at
high
pH,
and
Fe(OH)2
(Ksp
=
1015)
is
more
soluble
than
Fe(OH)3.
Consequently,
the
reduction
potential
falls
with
increasing
pH,
because
[Fe3+]
declines
more
rapidly
than
[Fe2+].
2
1/2
9
Worked
Problem
14.3:
Eo[w]
and
Ksp
of
Fe[OH]3
Q:
The
Fe3+/2+
standard
potential
[equation
[14‐13]
is
0.77
V.
From
this
value
and
the
Ksp
calculate
Eo[w]
for
the
reduction
of
Fe[OH]3
to
Fe2+
[see
Table
14.3].
A:
Eo[w]
is
the
Fe3+/2+
potential
at
pH
7,
when
Fe[OH]3
is
certainly
precipitated.
This
potential
can
be
calculated
from
the
Nernst
equation
E
=
0.77
–
0.059{log([Fe2+]/[Fe3+])}
and
[Fe3+]
can
be
calculated
from
Ksp
=
[Fe3+][OH‐]3.
Substitution
gives
E
=
0.77
–
0.059{log[Fe2+]
‐
log(Ksp)
+
3log[OH‐]}
At
pH
7,
[OH‐]
=
1.0
x
10‐7
M
E
=
0.77
‐
0.059{log[Fe2+]
+
37
‐
21}
=
‐0.17–
0.059{log[Fe2+]}
which
is
the
Nernst
equation
for
Fe[OH]3
reduction,
with
Eo[w]
=
‐0.17
V.
Worked
Problem
14.4:
Effective
Oxygen
Potential
Q.
The
first
plateau
in
Figure
14.1,
corresponding
to
O2
reduction,
is
at
0.5
V,
whereas
the
Eo[w]
value
[Table
14.3]
is
0.812
V.
What
might
account
for
this
difference?
A:
Assuming
that
the
environmental
pH
is
7,
the
difference
must
arise
from
the
O2
concentration
dependence.
The
potential
diminishes
with
decreasing
O2
concentration.
Recall
that
E
=
1.24
–
0.059log{1/PO21/4[H+]}
[14‐14]
or,
at
pH
7,
E
=
0.812
–
0.059log{1/PO21/4}
[14‐15]
If
E
=
0.50,
then
substituting
into
equation
[14‐15]
gives:
log{1/PO21/4}
=
[‐logPO2]/4
=
‐0.312/(‐0.059)
=
5.28
or
PO2
=
10‐21.1
atm.
This
may
seem
a
bizarrely
low
value,
but
it
reflects
the
fact
that
when
microbes
are
actively
respiring
in
an
aqueous
medium,
they
draw
down
the
O2
to
very
low
levels
in
their
immediate
vicinity.
14S.4
Electron
and
Proton
Affinities
Are
Linked:
pE
versus
pH
Most
reduction
reactions
are
accompanied
by
proton
uptake,
and
oxidations
generally
lead
to
proton
release.
Since
adding
an
electron
increases
negative
charge
while
adding
a
proton
decreases
it,
the
coupling
of
electron
and
proton
transfers
is
a
simple
consequence
of
the
tendency
to
lower
the
energy
of
the
molecule
by
neutralizing
charge.
This
coupling
leads
to
a
strong
dependence
on
the
solution
pH
for
most
half‐reaction
potentials.
10
The
hydrogen
electrode
potential
is
pH‐dependent
because
two
protons
combine
with
two
electrons
in
producing
H2.
From
reaction
[14‐4]:
pE
=
0
–
[1/2]logPH2
+
log[H+]
[14‐16]
If
pH2
is
maintained
at
1
atm,
then
pE
=
–pH.
Thus,
hydrogen
gas
is
far
more
reducing
in
alkaline
than
in
acid
solution.
For
example,
at
pH
8,
pE
=
–8
V,
and
the
hydrogen
electrode
potential
is
–0.47
V,
nearly
half
a
volt
more
negative
(reducing)
than
at
pH
=
0.
At
the
same
time,
oxygen
is
harder
to
reduce
in
alkaline
solution
(or,
conversely,
O2
is
more
oxidizing).
From
reaction
[14‐3]:
pE=pEo
+
[1/4]logPO2
+
log[H+]
=
20.75
–
pH,
at
PO2=1
atm
[4‐17]
and
the
electrode
potential
at
pH
7
is
0.83
V.
Because
pH
7
is
closer
to
most
biologically
and
environmentally
relevant
conditions
than
is
pH
0,
electrode
potentials
and
pE
values
are
often
cited
for
pH
7,
as
in
Table
14.3.
The
pEo[w]
values
are
pEo
values
recalculated
for
pH
7.
Since
both
the
hydrogen
and
oxygen
electrodes
have
the
same
pH
dependence,
the
difference
between
them
is
pH‐independent,
reflecting
the
fact
that
there
is
no
gain
or
loss
of
protons
in
the
overall
reaction
for
hydrogen
oxidation
by
oxygen
[reaction
[14.2]].
Thus,
the
potential
of
the
hydrogen‐oxygen
fuel
cell
is
independent
of
the
pH
of
the
cell
compartments,
even
though
the
individual
electrode
potentials
are
strongly
affected.
Even
if
no
protons
appear
explicitly
in
a
half‐reaction,
the
potential
may
be
pH‐dependent
because
of
secondary
acid‐base
reactions.
For
example,
the
potential
of
the
Fe3+
reduction
half‐reaction
[14‐13]
potential
has
no
proton
dependence
per
se,
but
the
equilibrium
quotient
[Fe2+]/[Fe3+]
is
highly
dependent
on
pH
because
of
the
acidic
character
of
Fe3+.
At
quite
low
pH,
it
forms
a
series
of
hydroxide
complexes,
and
precipitates
as
the
highly
insoluble
Fe(OH)3
(pKsp
=
38).
In
contrast,
Fe2+
forms
hydroxide
complexes
only
at
high
pH,
and
Fe(OH)2
(pKsp
=
15)
is
more
soluble
than
Fe(OH)3.
Consequently,
the
reduction
potential
falls
with
increasing
pH,
because
[Fe3+]
declines
more
rapidly
than
[Fe2+].
The
relationship
between
pE
and
pH
is
conveniently
illustrated
in
a
diagram
such
as
that
shown
for
the
Fe3+/2+
couple
in
Figure
14.2.
The
regions
of
the
diagram
are
labeled
according
to
the
dominant
chemical
species
present,
and
the
lines
show
the
pE/pH
dependence
at
the
edges
of
these
stability
fields.
Thus,
the
horizontal
line
at
the
top
left
of
the
diagram
represents
pE
=
13.2,
the
value
expected
for
an
equimolar
solution
of
Fe3+
and
Fe2+
in
the
absence
of
hydroxide
reactions.
Fe3+
predominates
above
this
line,
while
Fe2+
predominates
below
the
line.
The
vertical
line
at
pH
=
3.0
arises
because
of
the
precipitation
of
Fe(OH)3.
This
happens
when
the
Ksp
is
exceeded,
which
depends
on
the
pH
and
on
[Fe3+].
For
the
purposes
of
illustration,
the
iron
concentration
was
set
at
10–5
M
in
drawing
Figure
14.2.
From
Ksp
=
10
–38
=
[Fe3+][OH–]3
[4‐18]
11
we
calculate
[OH–]
=
{10–38/[Fe3+]}1/3
=
10–11,
giving
pH
=
3.
Above
this
pH,
Fe(OH)3
precipitates,
and
[Fe3+]
declines
in
conformity
to
the
Ksp
and
the
pH.
The
effect
of
this
decline
on
pE
is
seen
in
the
line
sloping
downward
from
pH
=
3.
This
line
has
a
slope
of
3.0
because
of
the
three
hydroxide
ions
per
iron
in
Fe(OH)3.
Rearranging
equation
[4‐18],
we
have
log[Fe3+]
=
–38
+
3
pOH
[4‐19]
and
since
pOH
=
14
–
pH
and
pE
=
13.2
–
log{[Fe2+]/[Fe3+]}
[4‐20]
the
dependence
of
pE
on
pH
is
given
by
pE
=13.2
–
log[Fe2+]
+
log[Fe3+]
=
22.2
–
3
pH
(with[Fe2+]=10–5
M)
Above
this
line,
Fe2+
is
oxidized
and
precipitates
as
Fe(OH)3,
while
below
the
line
Fe(OH)3
dissolves
by
reduction
to
Fe2+.
Figure
14.2
pE/pH
diagram
for
a
Fe‐O‐H
system.
12
The
second
vertical
line,
at
pH
=
9.0,
arises
from
the
precipitation
of
Fe(OH)2.
From
Ksp
=
10–15
=
[Fe2+][OH–]2
[4‐21]
we
calculate
[OH–]
=
{10–15/[Fe2+]}1/2
=
10–5,
giving
pH
=
9.
Above
this
pH,
Fe(OH)2
precipitates
and
[Fe2+]
declines.
The
line
sloping
downward
above
this
pH
represents
the
phase
boundary
between
Fe(OH)3
and
Fe(OH)2.
Its
slope,
minus
one,
is
the
difference
between
the
two
hydroxides
of
Fe(OH)2
and
the
three
hydroxides
of
Fe(OH)3.
Rearranging
equation
[4‐21],
we
have
log[Fe2+]
=
–15
+
2
pOH
[4‐22]
When
both
hydroxides
are
present,
equations
[4‐19]
and
[4‐22]
can
be
substituted
into
equation
[4‐20]
to
obtain
the
dependence
of
pE
on
pH:
pE
=
13.2
–
log[Fe2+]
+
log[Fe3+]
=
–9.8
+
pOH
=
4.2
–
pH.
The
top
and
bottom
diagonal
lines
in
Figure
14.2
represent
the
pE/pH‐
dependence
of
the
hydrogen
and
oxygen
reduction
reactions,
equations
[14‐3]
and
[14‐4].
They
represent
the
stability
limits
for
aqueous
solutions.
Below
the
bottom
diagonal
line,
water
is
reduced
to
hydrogen,
while
above
the
upper
diagonal
line,
water
is
oxidized
to
oxygen.
Figure
14.2
is
not
a
complete
diagram
of
the
Fe3+/2+
system
because
soluble
complexes
of
the
ions
have
been
omitted
from
consideration.
Hydroxide
complexes,
already
mentioned
above,
are
always
present
in
aqueous
solutions,
but
they
predominate
only
in
narrow
regions
of
pH
and
do
not
greatly
affect
the
appearance
of
the
diagram.
Other
complexing
agents
can
have
significant
effects.
The
naturally
occurring
anions
chloride,
carbonate,
and
phosphate
bind
Fe3+
and
can
lower
[Fe3+]
and
therefore
pE,
as
can
organic
constituents
of
soils,
especially
the
humic
acids,
which
can
either
bind
the
Fe3+
to
soil
particles
or
form
soluble
chelates
with
the
Fe3+.
Despite
these
complexities,
Figure
14.2
presents
the
main
features
of
the
Fe3+/2+
system,
which
is
dominated
by
the
species
Fe2+
and
Fe(OH)3.
Over
most
of
the
available
pH
range,
3–9,
these
are
the
only
significant
species.
14.4 Biological oxidations
Bacteria also catalyze oxidation of reduced substances by molecular oxygen,
even though such reactions can occur spontaneously in an aerobic environment. Thus
HS – oxidation to sulfate is catalyzed by sulfide oxidizers. These bacteria manage to
extract energy from the HS–/SO42– and O2/H2O redox couples. Another important
oxidation process is nitrification, the conversion of NH4+ to nitrate ion. Since plants take
up and utilize nitrogen mainly in the form of nitrate, this is a key reaction in nature,
especially in connection with the use of ammonium salts in fertilizers [see p. ?]. The
process actually occurs in two steps, ammonium to nitrite, NO2–, and nitrite to nitrate:
NH 4+
+
2H2O
=
NO2–
+
8H+
+
6e–
NO2–
+
H2O
=
NO3–
+
2H+
+
2e–
[14‐23]
[14‐24]
13
These half reactions are catalyzed by two separate groups of bacteria,
Nitrosomonas and Nitrobacter, each utilizing the oxidizing power of O2 to extract energy
from the process.
In summary, redox potential can be considered as a kind of chemical switch in the
aqueous environment, one that determines the sequence by which oxidants are utilized
by microorganisms. Changes in redox potential can have important consequences for
environmental pollution (Table 14.2).
14.5 Aerobic Earth
O2 was not always a constituent of the atmosphere; it arose from the evolution
of life itself. The primitive Earth had an atmosphere derived from outgassing of the
minerals in the interior. Once the surface cooled sufficiently to condense water, and with
it acidic gases like HCl and SO2, the main atmospheric constituents would have been N2
and CO2.
Life arose quite early in the Earth’s history; microfossils resembling modern
cyanobacteria have been found in 3.5 billion year-old rocks. How life started is unknown,
and remains one of the great scientific issues of our time. It is known that simple
organic molecules are common in the universe, and are present in meteorites, which
would have bombarded the young Earth. Laboratory experiments show that they could
also have been formed from inorganic precursors when subjected to electric discharges
from lightning, or to ultraviolet irradiation. The ultraviolet flux would have been intense,
since, in the absence of an oxygen atmosphere, the Earth would have lacked an ozone
shield. Many of the organic building blocks of organisms could have been produced in
this way. Alternatively, the building blocks might have been formed on the surfaces of
sulfide minerals under the high pressures and temperatures found in hydrothermal vents
on the sea floor. [These vents are found in regions where the crustal plates are being
formed through upwelling from the Earth’s mantle]. Recent experiments show that
complex organic molecules can be formed in this way. Still a third possibility is that the
building blocks came from outer space, since complex organic molecules are found in
some meteorites. How the building blocks were assembled into the first self-replicating
organisms remains an unanswered question, although many ingenious proposals have
been put forward.
The first organisms were heterotrophic, assimilating organic compounds from
their environment. Since there was no O2, they must have obtained their energy from
redox reactions other than respiration, similar to the modern anaerobic processes
discussed in the preceding section. The splitting of simple organic molecules, such as
acetic acid:
CH3COOH
=
CH4
+
CO2
[14‐25]
may have been the first of such processes; this reaction still provides energy for modern
acetogenic bacteria.
However, photosynthesis evolved quite early on, probably in the cyanobacteria
mentioned above, which survive today as photosynthetic organisms in the oceans.
Photosynthesis made these organisms autotrophic, capable of synthesizing their own
organic molecules from CO2. They had a strong selective advantage over heterotrophs.
In addition to the fossil evidence mentioned above, carbon isotope measurements on the
fossil organic carbon show photosynthesis to be at least 3.5 billion years old. The fossil
14
carbon is found to be depleted in the stable 13C isotope, relative to 12C, as a result of the
slightly slower diffusion of 13CO2 and its slower rate of capture by the CO2-fixing enzyme
ribulose bisphosphate carboxylase.
O2 was a byproduct of the rise of autotrophic organisms. Because of the
reactivity of O2, it would have been a toxic byproduct; most anaerobes are very sensitive
to O2, and cannot survive in an aerobic environment. However, O2 did not become a
significant constituent of the atmosphere for a long time after the advent of
photosynthesis, because it was first consumed by oxidizable elements in the ocean and
in the Earth’s crust, particularly iron and sulfur. The early ocean would have had a high
concentration of Fe2+, which is abundant in silicate minerals of the mantle, and is quite
soluble, in contrast to Fe3+. Photosynthetic O2 would initially have been used up by
reaction with Fe2+ to produce precipitates of Fe[OH]3. Indeed ferric oxide begins to be
seen in sedimentary rock that is about 3.5 billion years old, occurring in banded iron
formations, in which Fe2O3 is interbedded with siliceous sediment. These formations reach
a peak occurrence in rock which is 2.5 to 3 billion years old.
Once the oceanic Fe2+ was used up, the accumulating O2 attacked oxidizable
minerals on land, principally FeS2 [pyrite], producing Fe[OH]3 and H2SO4 [the same
chemistry that still produces acid mine drainage (p. ?)]. Evidence for this transition is
found in the occurrence of red beds, deposits of Fe2O3 found in geologic layers of
terrestrial origin, starting about 2 billion years ago, after the last of the banded iron
formations were formed.
Finally, when the rate of O2 production exceeded its rate of consumption by
exposed oxidizable material, the O2 concentration in the atmosphere began to rise,
permitting the evolution of respiring organisms. Fossil evidence of eukaryotic organisms
has been found in rocks that are 1.3-2 billion years old. Eukaryotes [in contrast to the
more primitive prokaryotes] have mitochondria, organelles which are specialized for
respiration. Some eukaryotes can survive on O2 at only 1% of the present
concentration, suggesting that this level was attained over 1 billion years ago. O2
production would have accelerated with the evolution of chloroplasts in the eukaryotes,
organelles which are specialized for photosynthesis. The rising O2 was also accompanied
by the production of stratospheric ozone, which permitted life to colonize the
continents, freed from the destructive effects of UV radiation. Fossils of multicellular
organisms have been found in sedimentary rocks that are 680 million years old, but the
rise of green plants, and with them the modern O2 atmosphere, dates to 400 million
years ago.
The time-line for the course of O2 production is shown schematically in Figure
14.3. The present atmospheric reservoir accounts for only about 2% of the estimated
cumulative production of O2, the rest having been used up in the oxidation of minerals.
Interestingly, the O2 concentration seems to have stayed at about 20% of the
atmospheric gases over the last 400 million years; this constancy suggests some sort of
feedback control. As with any reservoir (see p. ?), the amount of O2 reflects the
balance between the rate of production and the rate of consumption. Over geologic
time, O2 consumption results from exposure and weathering of reduced carbon-bearing
rock; this rate is set largely by the earth’s tectonic movements. O2 production results
from the burial of reduced carbon, whose rate depends [among other things] on the
total biomass. The biomass is limited, at least in part, by forest fires, and it is possible
that feedback control arises from the dependence of fires on the O2 concentration. It is
15
known that fires cannot be maintained when the O2 concentration is less than 15%, while
even wet organic matter burns freely at a concentration greater than 25%.*
-----
J.E. Lovelock (1974). Gaia: A New Look at Life on Earth. Oxford University Press:
Oxford, U.K.
----*
Figure
14.3
Cumulative
history
of
O2
released
by
photosynthesis
through
geologic
time.
Of
more
than
5.1
x
1022
g
of
O2
released,
about
98
percent
is
contained
in
seawater
and
sedimentary
rocks,
beginning
with
the
occurrence
of
banded
iron
formations
at
least
3.5
billion
years
ago
(bya).
Although
O2
was
released
to
the
atmosphere
beginning
about
2.0
bya,
it
was
consumed
in
terrestrial
weathering
processes
to
form
red
beds,
so
that
the
accumulation
of
O2
to
present
levels
in
the
atmosphere
was
delayed
to
400
mya.
Source:
W.H.
Schlesinger
(1997).
Origins.
Biogeochemistry:
An
Analysis
of
Global
Change
(2nd
ed.)
(San
Diego:Academic
Press).
If carbon burial has balanced O2 accumulation for the last 400 million years, what
accounts for the rising O2 level starting 4 billion years ago? A much larger carbon burial
rate sees unlikely. It has recently been suggested† that UV photolysis of CH4 could have
provided the driving force. Methane production would have been much higher when O2
levels were low; methane-producing anaerobes would have been abundant, and the
methane would have escaped to the atmosphere without oxidation. In the absence of
the ozone UV shield, the methane would have been exposed to photons energetic
enough to break the C-H bonds. At the top of the atmosphere, the light H atoms would
have escaped Earth’s gravitational field, and would have been lost to space. This
removal of oxidizable H atoms from the earth-atmosphere system would provide a
mechanism for O2 accumulation.
16
--†
D.C.
Catling
et
al.
Biogenic
methane,
hydrogen
escape,
and
the
irreversible
oxidation
of
early
Earth
(2001).
Science
293:839‐843.
‐‐‐
14.6 Water as Ecological Medium
14.6a The euphotic zone and the biological pump
Biological productivity depends on primary producers, organisms that fix carbon
via photosynthesis, and provide the food for the animal food chain. In water, the primary
produces are cyanobacteria, phytoplankton and algae. Because of their dependence on
sunlight, they are limited to the region near the surface, where sunlight can penetrate.
This is the euphotic zone. Its depth depends on the clarity of the water.
Most biological activity takes place in the euphotic zone. The primary producers
are eaten by animals or decomposed by bacteria, in a continuing cycle of photosynthesis
and respiration. However, because of gravity, some dead organisms fall below the
euphotic zone. In the deeper layers bacterial decomposition continues and the waters
are enriched in carbon and the other elements of life. Because of thermal stratification,
there is little physical mixing between the warmer surface layer and the cold deep layer.
Consequently, there is a kind of ‘biological pump’, which transfers carbon, nitrogen,
phosphorus, sulfur, etc., from the surface to the deep layers and the sediments. Figure
14.4 shows the effect of biological production on the depth profiles of nitrate and iron,
as well as oxygen, in the north Pacific. O2 is high at the surface and diminishes sharply
over the first few hundred meters. Nitrate and iron are drawn down at the surface, due
to uptake by organisms, but increase sharply with depth as the organisms are
decomposed; below the surface layer their concentrations remain at elevated levels.
In the oceans, the biological pump is responsible for increasing the carbonate
concentration of the deep layers with respect to the surface layers. This drawing down
of carbonate from the surface increases the rate of transfer of CO2 from the
atmosphere. This is an important contribution to the global carbon cycle. It has been
calculated that the atmospheric CO2 level would double in the absence of the biological
pump.
17
Figure
14.4
Vertical
distribution
of
Fe,
NO3,
and
O2
in
the
central
North
Pacific
Ocean.
Source:
J.H.
Martin
et
al.
(1989).
VERTEX:
Phytoplankton/iron
studies
in
the
Gulf
of
Alaska.
Deep
Sea
Research
36:649‐680.
14.6b Eutrophication in freshwater lakes.
Because the supply of oxygen is restricted, the species that inhabit an aquatic
ecosystem are in a dynamic balance, one that is easily disturbed by humans. In water,
the O2 concentration falls with increasing distance from the air-water interface. Thus,
aerated soils support oxygen-utilizing microbes as well as higher life forms, while deeper
in the soil, in the saturated zone where the soil pores are filled with water, anaerobic
bacteria dominate and utilize progressively lower E0[w] redox couples. Likewise in lakes,
the sediments are generally oxygen-starved and rich in anaerobic microorganisms, while
in the water column above, the O2 concentration increases towards the surface. The
concentration of O2 at the surface is increased not only because the surface is in
contact with air, but because the surface waters support the growth of vegetation and
algae, which release O2 as a product of photosynthesis.
The biological productivity of a temperate lake varies annually in a cycle (Figures
14.5 and 14.6). The onset of winter diminishes the solar heating of the surface. The
thermal stratification disappears and the water’s density becomes uniform, allowing easy
mixing by wind and waves, which brings nutrient- rich waters to the surface. In winter,
the nutrient supply is high, but productivity is inhibited by low temperatures and light
18
levels. Spring brings sunlight and warming, leading to a bloom of phytoplankton and
other water plants. As plant growth increases, the nutrient supply diminishes and
phytoplankton activity falls. Bacteria decompose the dead plant matter, gradually
replenishing the nutrient supply, and a secondary peak of phytoplankton activity is
observed in the autumn. Because the nutrient supply is limited in unpolluted waters, the
BOD in the surface waters rarely outstrips the available oxygen.
Figure
14.5
Seasonal
cycling
of
nutrients
in
lakes.
EZ
=
thermocline
and
end
of
the
euphotic
zone;
stipple
represents
phytoplankton
growth;
N
→
signifies
direction
of
nutrient
flow;
enclosed
arrows
indicate
circulation
of
waters.
The
solid
line
at
the
right
is
the
temperature
profile
of
the
water
column.
19
Figure
14.6
Seasonal
phytoplankton
productivity
as
a
function
of
sunlight
and
nutrient
concentration.
Source:
Adapted
from
W.D.
Russel‐Hunter
(1970).
Aquatic
Productivity
(New
York:
Macmillan
Publishing
Co.,
Inc.).
This natural cycle can be disrupted, however, by excessive nutrient loading from
human sources such as wastewaters or agricultural runoff. The added nutrients can
support a higher population of phytoplankton, producing “algal blooms” (Figure 14.7).
When masses of algae die off, their decomposition can deplete the oxygen supply, killing
fish and other life forms. If the oxygen supply is exhausted, the bacterial population may
switch from predominantly aerobic bacteria to mainly anaerobic microorganisms that
generate the noxious products (NH3, CH4, H2S) of anaerobic metabolism.
This process is called eutrophication or, more accurately, cultural eutrophication.
Eutrophication is the natural process whereby lakes are gradually filled in. Over time, an
initially clear (oligotrophic) lake eutrophies, filling with sediment and becoming a marsh,
and then dry land. This process normally proceeds over thousands of years because
biological growth and decomposition in the euphotic zone are closely balanced — the
surface layers remain well oxygenated, and only a small fraction of biological production
is deposited as sediment. When this balance is upset by overfertilization of the water,
the eutrophication process accelerates greatly.
20
Figure 14.7 False-color LandSat image of cyanobacteria surface blooms in lakes
Mendota and Monona, Madison, Wisconsin. (Image courtesy NorthTemperate Lakes Long
Term Ecological Research Program, http://lter.limnology.wisc.edu.) (From Carpenter, S.R.
PNAS (2008)105,11039.)
14.6c Nitrogen and phosphorus: the limiting nutrients.
The slow pace of natural eutrophication reflects the nutrient dynamics of an
aquatic ecosystem (Figure 14.8).The nutrients are assimilated from the environment
by the primary producers, which serve as food for secondary producers, including fish.
Dead plant and animal tissues are decomposed by bacteria, which restore the nutrients
to the water. The growth of the primary producers is controlled by the limiting nutrient,
the element that is least available in relation to its required abundance in the tissues. If
the supply of the limiting nutrient increases through over fertilization, the water can
produce algal blooms, but not otherwise; conversely, management of the aquatic
ecosystem requires that the supply of the limiting nutrient be restricted.
The major nutrient elements are carbon, nitrogen, and phosphorus, which are
required in the atomic ratios 106:16:1, reflecting the average composition of the
molecules in biological tissues. Numerous other elements are also required, including
sulfur, silicon, chlorine, iodine, and many metallic elements. Because the minor elements
are required in small amounts, they can usually be supplied at adequate rates in natural
waters. On the other hand, carbon, the element required in the largest amounts, is
plentifully supplied to phytoplankton from CO2 in the atmosphere. Phytoplankton outrun
the supply of CO2 only under conditions of very rapid growth such as in some algal
blooms. In these cases, the pH of the water can be driven as high as 9 or 10 through the
required shift of the carbonate equilibrium
HCO3– + H2O = OH- + CO2
[14-26]
The increase in pH can in turn alter the nature of the algal growth, selecting for varieties
that are resistant to high pH.
21
Figure
14.8
Nutrient
cycling
in
an
aquatic
ecosystem.
Normally, the limiting nutrient element is either N or P. Although nitrogen makes
up 80% of the atmosphere, it is unavailable except through the agency of N2-fixing
bacteria, living in symbiotic association with certain species of plants. On land, these
species are rare enough to make nitrogen the limiting nutrient under most conditions. In
water, however, N2-fixing algal species are common, and nitrate ions are often abundant
because of runoff from the land. Consequently, nitrogen is not usually limiting, although
it may be in some regions, especially the oceans, where nitrate concentrations are low.
This leaves phosphorus as the element that is usually limiting to growth, at least
in fresh water. A 37-year study at the Experimental Lakes Area in Canada discovered
‡
that phosphorous is the main cause of lake eutrophication. In this study, the
experimental lake was fertilized with constant inputs of phosphorus and decreasing
amounts of nitrogen, and then during the last 16 years, phosphorous alone was added.
Nitrogen-fixing cyanobacteria were able to provide the nitrogen inputs necessary from
the atmosphere to allow biomass production in proportion to the phosphorus added to
the lake. The lake was highly eutrophic, despite no additional inputs of nitrogen.
‡
D.W. Schindler et al. (2008) Eutrophication of lakes cannot be controlled by reducing nitrogen
input: Results of a 37-year whole-ecosystem experiment. Proceedings of the National Academy
of Sciences 105:11254-11258.
Phosphorus has no atmospheric supply because there is no naturally occurring
gaseous phosphorus compound. Moreover, the input of phosphorus in runoff from
22
unfertilized lands is usually low because phosphate ions, having multiple negative
charges, are bound strongly to mineral particles in soils. In surface waters, most of the
phosphorus is contained in the plankton biomass; the phosphorus availability depends on
recycling of the biomass by bacteria.
Some of the phosphorus is lost to the deeper water and to the sediments when
dead organisms sink. When a lake turns over in winter, the phosphorus in the deep
waters is carried to the surface and supports the plankton bloom in the spring. Whether
this phosphorus is available to the surface waters depends on conditions in the lake. At
the bottom, phosphate ions may be adsorbed onto particles of iron and manganese
oxide. However, when the sediment becomes anoxic, the metal ions are reduced to the
divalent forms, the oxides dissolve, and the phosphate ions are released into solution
(see notes on maganese and iron oxides in Table 14.2). Phosphate solubility is also
increased through acidification since at successively lower pH values, HPO42–, H2PO4–, and
H3PO4 are formed (p. ?).
The most notorious instance of phosphate-induced eutrophication was in Lake
Erie, which “died” in the 1960s. Excessive algal growth and decay killed most of the fish
and fouled the shoreline. A concerted effort by the United States and Canada to reduce
phosphate inputs was put into effect in the 1970s. Over $8 billion was spent in building
sewage treatment plants to remove phosphates from wastewater, and the levels of
phosphate in detergents were restricted. These efforts, along with other pollution
control measures, succeeded in bringing the lake back to life. Commercial fisheries have
revived, and the beaches are once again in use.
14.6d Anoxia and coastal marine ʻdead zonesʼ
Enhanced nutrient loading is also affecting many coastal areas, creating ‘dead
zones’ where marine life is curtailed by oxygen depletion. Over 400 dead zones have
been identified, their global distribution corresponding roughly to the ‘human
footprint’(Figure 14.9).
The progression of marine hypoxia is illustrated in Figure 14.10. As nutrient
input increases, there is an initial pulse of energy up the food chain, but then a steady
decrease as higher animals die off, and microbes take over. Because seawater is rich in
sulfate salts, the favored reaction under anaerobic conditions is sulfate reduction to
hydrogen sulfide (H2S), a chemical that is extremely toxic to fish and humans. Although
H2S is generally confined to the lower layers of seawater, during storms the deeper,
anoxic layers can mix with surface layers, exposing aquatic life to the deadly gas.
23
Figure 14.9 Global distribution of 400-plus systems that have scientifically
reported
accounts
of
being
eutrophication‐associated
dead
zones.
The
map
is
color
coded
according
to
the
normalized
human
influence
Source:
R.J.
Diaz
and
R.
Rosenberg
(2008).
Spreading
Dead
Zones
and
Consequences
for
Marine
Ecosystems,
Science
321:926‐929.
Figure
14.10
Energy
flow
in
a
marine
ecosystem
as
eutrophication
progresses.
In
healthy
waters
(green
–
‘normoxia’)
mobile
predators
feed
on
the
organisms
that
live
on
the
seabed
(benthic
organisms).
As
oxygen
is
depleted
in
the
water
(orange
‐
‘hypoxia’),
a
short
pulse
of
energy
is
followed
by
a
decline
in
mobile
predators.
When
no
oxygen
is
left
in
the
water
column
(red
–
‘anoxia’),
microbes
process
all
of
the
energy
and
form
H2S.
Source:
R.J.
Diaz
and
R.
Rosenberg
(2008).
Spreading
Dead
Zones
and
Consequences
for
Marine
Ecosystems,
Science
321:926‐929.
In the U.S., the most notorious dead zone, in the Gulf of Mexico (Figure
14.11). In a vast area at the mouth of the Mississippi River (about 18,000 km2 , which
has more than doubled since 1980) the O2 concentration is too low to support aquatic
life during the spring and summer. These are the seasons of great algal blooms,
24
resulting from overfertilization of the Gulf by the nutrients in the river outflow. More
than 40% of U.S. commercial fisheries are located in the Gulf of Mexico, and these have
been hard hit by the annual appearance of the dead zone.
The Mississippi drains the vast mid-continent farmlands, and delivers 1.5 million
tons of dissolved nitrogen annually to the Gulf, agriculture accounting for 80% of the
total. However, there has been great controversy over whether N or P is the main culprit
in producing the dead zone. Because of the transition from fresh to salt water, it is likely
that both are important, and both need to be controlled.
Figure
14.11
The
“dead
zone”
in
the
Gulf
of
Mexico
due
to
nutrient
enrichment
in
the
drainage
basin
of
the
Mississippi
River.
The complexities of nutrient enrichment are illustrated by the Chesepeake Bay, in
the eastern U.S. Here eutrophication has long been evident, and appears to involve both
N and P. Levels of these nutrients in the estuary rise and fall annually in a seasonal
pattern(Figure 14.12). In winter, cold temperatures and lack of biochemical activity
allow the concentration of O2 to reach its annual maximum. At the same time, nitrogen
enters in large amounts because winter is the period of maximum freshwater flow, with
accompanying transport of sediment and runoff. Simultaneously, sedimentation is
removing phosphorus from the water column, mainly through the precipitation of
manganese and iron oxides, which absorb phosphorus efficiently and are insoluble under
aerobic conditions. (Phosphorus is also removed during the settling of organic debris.)
Beginning in the late spring and early summer, the oxygen levels decline due to increased
biological activity. Nitrogen concentrations also decline because 1) nitrogen is
incorporated into biomass and sinks as the organisms die; 2) little new nitrogen is
introduced in runoff; and 3) nitrogen is depleted as increasingly anoxic conditions force a
switch from oxygen to nitrate as oxidant.
25
Figure
14.12
Oxygen
concentration
in
water
overlying
the
sediments
with
major
seasonal
net
fluxes
of
nitrogen
and
phosphorus
(insets)
in
the
Patuxent
River
at
the
estuary
of
Chesapeake
Bay.
Source:
C.F.
D’Elia
(1987).
Too
much
of
a
good
thing:
Nutrient
enrichments
of
the
Chesapeake
Bay.
Environment
29(2):6‐11,
30‐33.
The opposite situation prevails for phosphorus. Under anaerobic conditions,
phosphorus is liberated from the sediments, in large part due to the reduction of
manganese and ferric oxides to Mn2 + and Fe2+. In the 2+ valence states, the metals are
soluble and release the bound phosphorus formerly adsorbed to the insoluble oxides of
the metals. The phosphorus is readily mixed with the surface layers given the mechanical
turbulence of estuarine environments. Thus, as conditions cycle from aerobic to
anaerobic and back, the phosphorus is continuously recycled between the surface waters
and the sediments. During anaerobic periods, phosphates are released to the water
column to be taken up by microorganisms; during aerobic periods, phosphates are
returned to the sediments. The amount of phosphate trapped in this cycle is vast, much
greater than the annual quantities entering the estuaries from sewage effluents or other
sources; it represents the cumulative inputs of many years. Thus, even though Maryland
and Virginia banned detergents with phosphates in the 1980s, phytoplankton
productivity is still excessive. Now the limiting nutrient may well be nitrogen, but
nitrogen inputs are very difficult to control. Chesapeake Bay receives some of the
highest atmospheric NOx emissions in the world, mainly due to the density of traffic in
26
the adjacent areas. Part of the strategy for cleaning up Chesapeake Bay might include
reducing NOx from vehicle exhausts, demonstrating once again the link between the
atmosphere and the hydrosphere.
14.6e Wetlands as chemical sinks.
Wetlands are typically anoxic and have large amounts of organic carbon; they
create a natural buffer zone for nearby fresh or marine waters by trapping nitrates. The
nitrates enter the wetlands in runoff, but are utilized by bacteria to oxidize stored
carbon via the reduction of nitrate to N2 or N2O, which are vented to the atmosphere
(Figure 14.13a). By depleting the nitrates before they can enter the estuary, the
surrounding wetlands limit the excessive growth of biomass and subsequent anoxic
conditions in the estuary. Restoration of wetlands has been proposed in many areas as a
means of reducing overfertilization from runoff.
Figure
14.13
(a)
Ability
of
wetlands
to
buffer
against
nitrate
and
sulfate
inputs
to
water
bodies;
(b)
under
conditions
where
wetlands
become
dry,
none
of
the
protective
reducing
reactions
occur.
In
addition,
accumulated
sulfides
may
oxidize
to
sulfate
as
sulfuric
acid,
and
leach
into
adjacent
rivers
or
lakes.
Source:
W.M.
Stigliani
(1988).
Changes
in
valued
capacities
of
soils
and
sediments
as
indicators
of
nonlinear
and
time‐delayed
environmental
effects.
Environmental
Monitoring
and
Assessment
10:245‐307.
If the original wetlands are of marine origin, they are likely to contain high
concentrations of sulfur in the form of reduced sulfide minerals such as pyrite. Under the
redox/pH conditions prevalent in wetlands, these sulfides are highly insoluble and
27
immobilized (Figure 14.13a). Draining the wetlands (Figure 14.13b) exposes these
compounds to oxidizing conditions, producing a situation similar to acid mine drainage
(pp. ?).
One example of this phenomenon occurred in a coastal area of Sweden near the
Gulf of Bothnia, where wetlands were drained in the early 1900s for use as agricultural
lands. As shown in Figure 14.14, draining the wetland shifted the E[w]/pH conditions
diagonally to the upper left, from the values typical of waterlogged soils to conditions
close to those of acid mine drainage. The draining exposed sulfides to the atmosphere,
and their oxidation to sulfuric acid acidified the soil and nearby lakes. The pH in one of
these lakes, Lake Blamissusjon, dropped from 5.5 or higher in the last century to a
current value of 3. Even though agricultural activities ceased in the 1960s, the lake has
not recovered; it is widely known as the most acidic lake in Sweden.
Figure
14.14
Eh/pH
as
a
function
of
different
aquatic
environments.
Oval
enclosed
by
dashed
line
indicates
region
of
highest
solubility
of
heavy
metals.
Source:
Adapted
from
W.
Salomons
(1995).
Long‐term
strategies
for
handling
contaminated
sites
and
large‐scale
areas.
In
Biogeodynamics
of
Pollutants
in
Soils
and
Sediments,
W.
Salomons
and
W.M.
Stigliani,
eds.
(Berlin:
Springer‐Verlag).
Recently,
it
was
discovered
that
wetlands
actually
store
more
carbon
than
does
reforested
agricultural
land,
about
3000
vs
100
grams
of
carbon
per
square
meter
per
year.*
Wetlands
capture
carbon
by
absorbing
CO2
from
the
atmosphere
into
new
plant
growth,
but
once
the
plant
dies,
the
material
is
covered
by
water
and
mud
that
slows
reaction
with
O2
and
slows
decomposition.
Organic
peat
soils
formed
from
this
process
have
been
found
that
are
60
ft.
deep
and
7,000‐10,000
years
old.
Microbes
can
use
iron
oxides,
sulfate,
or
CO2
instead
of
oxygen
to
form
energy
from
redox
reactions
(p.
?),
but
when
CO2
is
used,
methane
is
produced.
Methane
is
a
more
potent
greenhouse
gas
than
CO2,
and
in
freshwater
marshes
the
amount
of
methane
produced
cancels
out
any
cooling
effects
of
CO2
absorbed
by
the
28
plant
material.
In
salt
water
marshes,
though,
the
concentrations
of
sulfate
are
so
high
that
microbes
do
not
have
to
use
CO2
as
an
electron
acceptor,
so
negligible
amounts
of
methane
are
produced.
CO2
absorption
by
salt
water
marshes
could
be
one
method
for
reducing
CO2
concentrations
in
the
atmosphere.
‐‐‐
*
J.
Pelley
(2008).
Can
wetland
restoration
cool
the
planet?
Environmental
Science
&
Technology,
8994.
‐‐‐
14.6f Redox effects on metals pollution.
Changes in the redox potential can have important consequences for
environmental pollution, especially with respect to metal ions such as cadmium, lead, and
nickel. In general, the solubility of heavy metals is highest in oxidizing and acidic
environments (Figure 14.14). At neutral to alkaline pHs in oxidizing environments,
these metals often adsorb onto the surface of insoluble Fe(OH)3 and MnO2 particles,
especially when phosphate is present to act as a bridging ion. When the redox potential
shifts to only slightly oxidizing or slightly reducing conditions as a result of microbial
action, and the pH shifts toward the acidic range, Fe(OH)3 and MnO2 in soils and
sediments are reduced and solubilized. The adsorbed metal ions likewise become
solubilized and move into groundwater (or into the water column of lakes when there is
Fe(OH)3 or MnO2 in the sediment). Conversely, if sulfate is reduced microbially to HS–
metal ions are immobilized as insoluble sulfides. But as we have seen, if sulfide rich
sediments are exposed to air through drainage or dredging operations, then HS– is
oxidized back to sulfate, and the heavy metal ions are released.
A particularly important instance of biological redox mediation of heavy-metal
pollution occurs in the case of mercury. Inorganic mercury, in any of its common valence
states, Hg0, Hg22+, and Hg2+, is not toxic when ingested; it tends to pass through the
digestive system, although Hg0 is highly toxic when inhaled. But the methylmercury ion
(CH3)Hg+ is very toxic, regardless of the route of exposure. The environmental route to
toxicity involves sulfate reducing bacetria that live in anaerobic sediments. As pat of
their metabolism these bacteria use methyl groups to produce acetate. When exposed
to Hg2+ the bacteria transfer the methyl groups to the mercury, producing (CH3)Hg+;
because methylmercury is soluble, it enters the aquatic food chain, where it is bioaccumulated in the protein-laden tissue of fish (see pp. ?).
14.6g Fertilizing the ocean with iron
Although nitrogen and phosphorus are the limiting aquatic nutrients near land, it
has become evident that in large areas of open ocean, it is actually iron that limits
biological production. Among the ‘trace metals’ essential for life, iron is required in
largest amounts. Iron is utilized in many enzymes involved in electron transport, and in
processing O2 and N2, as well as their reduction and (for N2) oxidation products. Thus all
organisms require a steady supply of iron. Since iron is abundant in the Earth’s crust,
iron limitation is not a problem for land plants, or for phytoplankton growing near land.
However the concentration of iron in the ocean is extremely low (Table 13.1), because
of the low solubility of Fe(OH)3 [p. ?] in the alkaline (pH = 8) seawater.
29
In much of the oceans the settling of dust from the land provides phytoplankton
with sufficient iron for growth. Prevailing winds blow sands from the Sahara and Gobi
deserts far out over the Atlantic and Pacific oceans. Recent satellite measurements show
a fairly good correlation between patterns of dust in the air and phytoplankton growth in
the oceans below. However, there are large areas which are relatively dust-free,
especially in the equatorial Pacific Ocean, and the waters ringing Antarctica at greater
than 60o south latitude, called the Southern Ocean. These areas have less
phytoplankton than could be supported by the available nitrogen and phosphorus. It has
been known for some time that adding iron to samples of these waters stimulates
phytoplankton growth in the laboratory, and a series of field experiments in the 1990’s
showed that spreading iron over areas of nutrient-rich ocean produced phytoplankton
blooms.
Iron limitation on biological productivity is an important ingredient in the carbon
cycle, because phytoplankton take up CO2 and transport some of it to the deep ocean
when they die. This is the mechanism of the ‘biological pump’ for CO2, discussed on p. ?.
In iron-limited areas, adding iron to the oceans could increase the speed of the biological
pump, drawing down the atmospheric CO2. Indeed it has been suggested that iron
supplementation could offer a ‘geoengineering’ solution to the problem of rising
atmospheric CO2. However, this solution has been set aside for several reasons:
a. The remedy would be very expensive, because the iron stimulation of
phytoplankton blooms is a transient effect. The blooms quickly fade as the excess iron
precipitates out of the photic zone. (The duration depends somewhat on the form of
the added iron. Ferrous salts are soluble, but rapidly oxidize to insoluble Fe(OH)3. Ferric
chelates are longer lived, but the chelating agents [see p. ?] would add to the expense).
Consequently iron would have to be added continuously to have a permanent effect.
b. Modeling indicates that the maximum effect on the atmospheric CO2
concentration would be a ~60 ppm lowering, making a relatively small difference in the
rising level.
c. There could be unforeseeable consequences to the biology of the oceans from
such an intervention.
d. There would have to be international agreement on ocean alteration,
particularly in the region of Antarctica, which is protected by international law.
However, the evidence that iron can fertilize the oceans, and that dust is an
important source of iron, raises the possibility that changes in global dustiness may have
contributed to the temperature changes that produced the ice Age. Data from ice cores
and from deep sea sediments indicate that there was much more iron in ocean water
during the ice ages. Thus the biological pump would have been stimulated; the ~60
ppm lowering in the CO2 level that might have been available from this mechanism
corresponds approximately to the CO2 lowering which is also detected in ice cores [see
p. ?]. The increased iron might have resulted from dust due to drying of the continents
and expansion of deserts. However, as is usual in reconstructing the past, it is difficult
to decide which factor is cause and which is effect.
Problems:
1.
Calculate
the
equilibrium
partial
pressure
of
oxygen
in
a
water
sample
at
pH
=
7.0,
which
contains
equal
concentrations
of
NH4+
and
NO3‐.
[See
Table
3.9
for
reduction
potentials].
30
2.
What
is
the
pE
value
of
an
acid
mine
water
sample
having
[Fe3+]
=
8.0
x
10‐
3
M
and
[Fe2+]
=
4.0
x
10‐4
M?
3.
(a)
What
class
of
molecules
is
responsible
for
most
of
the
reducing
power
in
aqueous
environments?
(b)
What
parameter
is
a
measure
of
reducing
power?
4.
Five
hundred
kg
of
n‐propanol
(CH3CH2CH2OH)
are
accidentally
discharged
into
a
body
of
water
containing
108
liters
of
H2O.
By
how
much
is
the
BOD
(in
milligrams
per
liter)
of
this
water
increased?
Assume
the
following
reaction:
C3H8O+9/2O2 = 3CO2+4H2O
5.
A
lake
with
a
cross‐sectional
area
of
1
km2
and
a
depth
of
50
meters
has
a
euphotic
zone
that
extends
15
meters
below
the
surface.
What
is
the
maximum
weight
of
the
biomass
(in
grams
of
carbon)
that
can
be
decomposed
by
aerobic
bacteria
in
the
water
column
of
the
lake
below
the
euphotic
zone
during
the
summer
when
there
is
no
circulation
with
the
upper
layer?
The
bacterial
decomposition
reaction
is:
(CH2O)n+nO2 = nCO2+nH2O
The
solubility
of
oxygen
in
pure
water
saturated
with
air
at
20°C
is
8.9
mg/l;
1
m3
=
1,000
liters.
6.
Assume
that
algae
need
carbon,
nitrogen,
and
phosphorus
in
the
atomic
ratios
106:16:1.
What
is
the
limiting
nutrient
in
a
lake
that
contains
the
following
concentrations:
total
C
=
20
mg/l,
total
N
=
0.80
mg/l,
and
total
P
=
0.16
mg/l?
If
it
is
known
that
half
the
phosphorus
in
the
lake
originates
from
the
use
of
phosphate
detergents,
will
banning
phosphate
builders
slow
down
eutrophication?
7.
Name
the
six
most
important
oxidants
in
the
aquatic
environment,
and
how
the
redox
potential
regulates
their
reactivity.
8.
(a)
If
a
lake
contains
high
concentrations
of
dissolved
Mn2+
and
Fe2+,
what
would
be
the
concentration
of
dissolved
NO3–
and
why?
(b)
What
environmental
effect
may
accompany
reduction
of
MnO2
and
Fe(OH)3?
9.
In
anaerobic
marine
environments,
what
toxic
gas
can
be
generated
and
by
which
reaction
(name
reactants
and
products)?
10.
Explain
the
“phosphate
trap”
in
the
estuary
of
Chesapeake
Bay.
Why
was
a
local
ban
on
phosphorus
in
detergents
not
particularly
helpful
in
mitigating
eutrophication
in
the
estuary?
31
11.
(a)
Explain
why
anaerobic
freshwater
wetlands
with
high
concentrations
of
organic
carbon
can
serve
as
natural
buffers
against
sulfates
and
nitrogen
oxides
(give
reactions).
(b)
When
other
oxidants
are
absent
from
such
wetlands,
which
redox
reaction
is
likely
to
predominate,
and
which
products
will
be
emitted?
12.
An
estuarine
creek
in
New
Jersey
contains
large
amounts
of
mercury
bound
as
sulfide
(with
K
=
10–52)
under
the
prevailing
environmental
conditions
(pH
=
6.8;
Eh
=
–230
mV).
Environmental
scientists
have
been
asked
to
assess
the
potential
impacts
of
the
polluted
sediments.
They
conclude
that
the
mercury
poses
no
danger
in
its
current
state.
However,
they
caution
against
any
action
that
would
expose
it
to
air
and
increase
its
redox
potential.
Explain
why
the
scientists
come
to
this
conclusion?
32

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