Ecosystems, energetics, and
chemical driving principles
In the last class, we reviewed
definitions of ecology, systems, and
ecosystems.
Watershed Ecology
Today we’ll flesh these out and think
more about the functional properties
of ecosystems.
Lecture notes, K. Limburg
In a watershed context, we’re very interested in
ecological function as well as in structure and spatial
context. How well did we capture this in the room
description exercise?
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Number of observations
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Systems analysis of classroom
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Positive vs negative feedbacks
Negative feedbacks: an increase in
process Y (e.g., predation) causes a
decrease in the status of X (e.g., the
prey population size) – called a
damping effect
Positive feedbacks do the opposite –
called amplifying effects. Can sometimes
be destabilizing in a system. But they can
also be powerful catalysts.
The relationships among parts of the food web
are important. This diagram reflects energy flow,
but also implies the negative feedbacks
The main sources of energy available to
ecosystems:
• solar energy
• chemically-bound energy
• (energy of position, geothermal)
Organisms that fix solar energy are
called photosynthetic autotrophs
(or photoautotrophs)
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The overall stoichiometric equation for
photosynthesis is (simplest form)
6CO2 + 6H2O + light Æ C6H12O6 + 6O2
Organisms that derive energy from breaking
chemical bonds are called chemoautotrophs (note:
our cells and organelles can be defined as
chemoautotrophs). The reaction of oxidizing
glucose during respiration runs the above reaction
in the opposite direction:
C6H12O6 + 6O2 Æ 6CO2 + 6H2O
At the ecosystem level, we often are
concerned with the sources of energy and
matter, and refer to sources as either
Autochthonous – generated within the
system (in-situ energy fixation)
Allochthonous – generated elsewhere,
transported into the system
More often, we define as
chemoautotrophs only those organisms
that obtain energy directly from chemical
(instead of biochemical) sources.
We call consumption of biochemically
fixed energy heterotrophy (“feeding on
others”)
Energy flows, but matter cycles,
between autotrophs and heterotrophs
The catabolic processes of primary
energy fixation are collectively referred
to as production (P), and the anabolic
processes of metabolism are called
respiration (R).
An index of system state is then the
ratio of P to R. This P:R ratio often
corresponds to succession or disturbance
within the system
We can also call a
system “net
autotrophic” or “net
heterotrophic”
Source: Odum 1971
steady state!
Source: Hynes 1975
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Carbon budget for the Hudson River
We’ve recognized that energy flow is critical
to material cycling in ecosystems. But what
makes a reaction happen?
The potential between the energy in chemical
(or biogeochemical) reactants, and products of a
reaction involving those precursors, is important
in determining whether or not the process will
occur
J. Willard Gibbs: a reaction will take place
spontaneously at constant temperature &
pressure if it produces useful work
Source:
Howarth et al.
1996
Gibbs free energy change, ∆G°’ = the maximum
amount of useful work that can be accomplished
at const. T, P
By definition, when ∆G°’ is negative (-), the
reaction occurs spontaneously
We can quantify ∆G°’ with an equation from
chemistry:
∆G°’ = -RT * ln Ke ,
where
R = universal gas constant = 8.314 J °K-1 mol-1
T = temperature in degrees Kelvin (25 °C = 298 °K)
Ke = the chemical equilibrium constant between
products and reactants
Gibbs free energy can be thought of as a measure of
the potential of a reaction to proceed:
A+B↔C+D
Like potential energy of mass at the top of a hill, the Gibbs free
energy depends on the gradient of activities (reactivities) between the
source chemical compounds (“reactants”) and the sink chemical
compounds (“products”)
Initial free energy level {A}i {B}i {C}i {D}i
rate determined
by chemical
kinetics
Let’s look back at the basic functions of the
metabolic cycle, in terms of Gibbs free energy:
Respiration:
C6H12O6 + 6O2 Æ 6CO2 + 6H2O
∆G°’ = - 686 kcal/mol (net amt of useful energy
yield)
Photosynthesis:
Final free energy level {A}f {B}f {C}f {D}f
6CO2 + 6H2O + light Æ C6H12O6 + 6O2
∆G°’ = + 686 kcal/mol (net amt of light energy
req.)
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Example of Gibbs free energy calculation:
C6H12O6 + 6O2
We calculate the Gibbs free energy of a
reaction as the difference between all the
G.f. energy in the products and all the G.f.
energy in the reactants:
Æ
6CO2 + 6H2O
Chemical
“species”
Role in reaction
∆G°’ (kcal/mol)
-------------------------------------------------------C6H12O6
∆G°’ = Σ (∆G°’products) - Σ (∆G°’reactants)
Let’s look at the energetics of the oxidation of
glucose in respiration:
Calculating Gibbs free energy in glucose
metabolism:
[ 1 x (-219.2) + 6 x (0) ]
product
- 94.2
0.0
H2O
product
- 56.7
Minerals that formed first also tend to be
the ones that chemically weather first, too.
= [ -565.2 – 340 ] – [ - 219.2 ] = - 686 kcal/mol
Weathering
proceeds with the
“easiest”
(spontaneous)
reactions first.
CO2
It is responsible for much of the chemical
weathering of rock to soil. Another way to
put this is that, as molecules react with each
other, they constantly “seek” a lower energy
state, losing energy along the way.
= [ 6 x (-94.2) + 6 x (-56.7) ] –
(Mg-based)
- 219.2
reactant
Gibbs free energy is the driving force in
many, if not most biogeochemical reactions.
∆G°’ = Σ (∆G°’products) - Σ (∆G°’reactants)
This diagram
illustrates how the
main classes of
rock minerals
weather, leaving
the most resistant
minerals.
reactant
O2
(Al-based)
Spontaneous weathering:
Oxidation of ferrous (iron II) to ferric
(iron III) :
2 FeO + ½ O2 Æ Fe2O3
∆G°’ = ∆G°’Fe2O3 – (2∆G°’FeO + ½ ∆G°’O2)
= - 742 – (-2*251 + 0) (units: kJ/mol)
= - 742 + 502
= -240 kJ/mol
Others require coreactions
Source: Schlesinger 1997
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A not-so-spontaneous weathering reaction:
sodium feldspar weathering to kaolinite by
contact with water saturated with CO2:
2NaAlSi3O8 + 9H2O + 2H2CO3 ⇔
Al2Si2O5(OH)4 + 2 Na+ + 2HCO3- + 4Si(OH)4
∆G°’ (kJ/mol)
NaAlSi3O8 (s)
sodium feldspar
H2O (l)
water
H2CO3 (aq)
carbonic acid
Al2Si2O5(OH)4 (s)
kaolinite
- 3698.7
- 237.2
- 623.4
- 3776.9
Na+ (aq)
sodium ion
- 261.9
HCO3- (aq)
bicarbonate ion
- 587.0
Si(OH)4 (aq)
silicic acid
- 1308.8
Q: why does this reaction proceed, even if the
thermodynamics are unfavorable?
This reaction wouldn’t get very far, if it weren’t
for the fact that the CO2-saturated water is
flowing past the rock.
The reaction proceeds a little bit, products are
borne away, and new water comes into contact
with the rock, permitting more product to form.
2NaAlSi3O8 + 9H2O + 2H2CO3 ⇔
Al2Si2O5(OH)4 + 2 Na+ + 2HCO3- + 4Si(OH)4
∆G°’ = (-3776.9 – 2x 261.9 - 2x 587.0 - 4x 1308.8)
Σ (∆G°’products)
- (-2x 3698.7 - 9x237.2 - 2x 623.4)
Σ (∆G°reactants)
= - 10,709.9 + 10,779 = + 69.1 kJ/mol
Re-dox and chemical potential.
Chemical reactions proceed from one equilibrial
state to another. As part of the interactions
of chemical species, electrons are rapidly
passed around amongst them.
Recall from chemistry that ions have different
valences (degree of negative or positive charge),
which is determined by whether or not they
have enough electrons (-) to balance off their
protons (+).
( thought experiment – visualize molecules
vibrating, electrons zinging between them)
At any given instant in time, there are very few free
electrons or free protons in a solution.
However, elements (and the chemical compounds they
form) have characteristic tendencies to associate
and dissociate with electrons or protons (we denote a
free proton as a hydrogen ion, H+ and a free electron
as e- )
The concept of reduction comes
from the idea that some compounds
have a tendency to lower their
valence (become more “electronegative”). How? By gaining
electrons.
+
Valence
scale
more
e--
Oxidation, on the other hand, is the
opposite tendency (named after the
element oxygen): oxidation is the
loss of electrons.
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Here is an example:
The re-dox “actors”:
A compound that gives away (donates) electrons is
called a reductant (think of it like a “perpetrator” – a
reductant perpetrates a reduction).
A compound that takes in (accepts) electrons is
called an oxidant.
Chemists (and therefore watershed ecologists) look
at the reduction and oxidation going on in chemical
reactions by dissecting the reaction equations into
“half reactions” and keeping track of the
dissociations of e- and H+ .
In aqueous solutions (like rivers, groundwater,
wetland porewaters, lakes, cells of organisms, etc),
redox reactions determine the distribution of
different forms taken by elements including
chemical compounds and various dissociations and
associations.
Iron in oxygenated waters:
1st half-reaction: 4 Fe2+ = 4 Fe3+ + 4 e(oxidation)
2nd ---- “ ---- : O2 + 4 H+ + 4 e- = 2 H2O
(reduction)
Entire redox reaction:
O2 + 4 H+ + 4 Fe2+ = 4 Fe3+ + 2 H2O
How does this relate to Gibbs free energy?
Another way to write the equation for ∆G°’ is
∆G°’ = -nF∆E , where
n = number of e- involved in reaction
F = Faraday const. = 23 kcal/mol
Another example -- sulfate reduction:
SO42- + 10H+ + 8e- = H2S + 4H2O
reduction
2CH2O + 2H2O = 2CO2 + 8H+ + 8e-
oxidation
SO42- + 2CH2O + 2H+ = H2S + 2CO2 + 2H2O
redox
Redox potential, qu’est-ce que c’est ???
Redox potential is an arbitrary measure of
how oxidized or reduced a compound is when
compared to a standard.
The standard reaction used is the
dissociation of hydrogen into proton and
electron:
H+ + e- ↔ ½ H2 (gas)
∆E = difference in redox potential in the
reaction
The difference between a
compound and the standard
can be measured with a
voltmeter that measures
the electrical potential in
volts (or more realistically,
in millivolts mv). The units
of measurement are called
Eh (similar to pH).
As Eh ↓, the solution has
more e- to give (is more
reduced) and as Eh↑, it will
accept more e- (is more
oxidized).
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