Lecture Notes - Aqueous and Environmental Geochemistry

Environmental Geochemistry
DM Sherman, University of Bristol
2011/2012
Introduction to Environmental
Geochemistry
Environmental Geochemistry
DM Sherman, University of Bristol
What is Environmental Geochemistry?
The chemical processes that couple the lithosphere,
hydrosphere, atmosphere and biosphere.
Applications:
•  Predicting the fate and transport of chemical
pollutants in rivers, lakes, soil and groundwater.
• Understanding controls on atmospheric CO2
• Identification of paleochemical proxies to determine
the history of environmental change.
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Environmental Geochemistry
DM Sherman, University of Bristol
2011/2012
Tools we will use…
•  Equilibrium thermodynamics. We will be
especially interested in chemical equilibria between
minerals and aqueous solutions.
•  Kinetic models (rates of dissolution and oxidation/
reduction and biological processes).
• Mass balance or “Box models” of open systems
(assessment of input and output fluxes).
How the unit is run…
•  Lecture notes, suggested reading and some selfassessment quizzes are at
http://mineral.gly.bris.ac.uk/envgeochem
• Practicals will focus on solving problems using
thermodynamics and box models. Practicals are not
assessed; worked solutions are posted on
“http://mineral.gly.bris.ac.uk/envgeochem”
•  Assessment is only via a end-of-year exam.
• This is a difficult unit, but it is (potentially) one of the
most valuable for your career.
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Environmental Geochemistry
DM Sherman, University of Bristol
2011/2012
Open Systems vs. Closed Systems
•  Open Systems are those which can exchange
matter (via “fluxes”) with their surroundings. With
time, the composition of an open system may
eventually reach steady state (when flux in = flux
out).
•  Closed systems cannot exchange matter with the
surroundings; with time, closed systems will reach
thermodynamic equilibrium (chemical equilibrium).
Reversible vs. Irreversible Processes
•  Reversible Processes can be reversed by a small
infinitesimal change without producing entropy.
•  Irreversible processes
• Respiration CH2O + O2 ➞ CO2 + H2O
• Photosynthesis CO2 + H2O ➞ CH2O + O2
Biological processes are irreversible; biological systems
are never in chemical equilibrium with their surroundings
(they can be in “steady state’, however).
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Environmental Geochemistry
DM Sherman, University of Bristol
2011/2012
Types of chemical equilibria"
• Acid base:
CO2 + H2O = H2CO3
H2CO3 = H+ + HCO3• Solubility
PbCO3(cerrussite) = Pb+2 + CO3-2
• Oxidation-Reduction
MnO2 + 2e- + 4H+ = Mn+2 + 2H2O
Types of chemical equilibria (cont.)"
• Complexation:
Zn+2 + 3Cl- = ZnCl3• Adsorption (complexation by mineral surfaces):
2>FeOH-0.5 + Pb+2 = (>FeOH)2Pb+
• Ion Exchange
0.165Ca+2 + Na0.33Mg3(Al0.33Si3.67)O10(OH)2
= 0.33Na+ + Ca0.165Mg3(Al0.33Si3.67)O10(OH)2
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Environmental Geochemistry
DM Sherman, University of Bristol
2011/2012
The Equilibrium Constant"
Consider a chemical reaction where A and B reversibly
transform to C and D:
αA + βB = γC + δD
Eventually, this reaction will go to equilibrium so that
ΔG = 0. At equilibrium, we will find that
€
& (a )γ (a )δ )
K=( Cα D β+
' (aA ) (aB ) *
with lnK = −ΔG 0 / RT
ΔG0 = (δµD0 + γµC0 ) − (αµA0 + βµB0 )
€
€
The Equilibrium Constant (cont.)"
& (a )γ (a )δ )
K=( Cα D β+
' (aA ) (aB ) *
€
with lnK = −ΔG 0 / RT
where ai is the activity of species i.
ΔG0 is the change in free energy when all species
are in their standard state.
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Environmental Geochemistry
DM Sherman, University of Bristol
2011/2012
Standard States and activities"
• For solid solutions, we can define a standard state of a
component as being the solid phase made of the pure
component. The activity of a species that is a pure solid (or
liquid) will be 1.
• For species in aqueous solution, we usually define the
standard state as being a 1 m solution with the properties
of infinite dilution (!). The activity of a species in an ideal
(dilute) aqueous solution will be its molal concentration.
• For a species in a gas phase, the standard state is a
pure gas (ideal) at 1 bar pressure. The activity of a species
in an ideal gas will be its partial pressure.
The p-function"
It is very convenient to linearize mass-action expressions
by taking the –log (or “p”) function:
αA + βB = γC + δD
€
% (a )γ (a )δ (
K = '' C α D β **
& (aA ) (aB ) )
pK = γ pC + δ pD − α pA − β pB
Be sure to practice writing out p expressions for reactions.
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Environmental Geochemistry
DM Sherman, University of Bristol
2011/2012
Example:Solubility of FeOOH"
The concentration of iron(III) in aqueous solution is
FeOOH(goethite) + 3H+ = Fe+3 + 2H2O pK = -0.53
However, we need to account for the hydrolysis of Fe3+:
Fe3+ + H2O = FeOH+2 + H+
pK = 2.19
Fe3+ + 2H2O = Fe(OH)2+ + 2H+
pK = 5.67
Fe3+ + 3H2O = Fe(OH)30 + 3H+
pK = 12.56
Fe3+ + 4H2O = Fe(OH)4- + 4H+
pK = 21.60
Solubility of FeOOH (cont.)"
By taking the “p-function” (i.e., -log) of the equilibrium
expressions, we get:
p[Fe+3] = pK1 + 3pH = -0.53 + 3pH
(1)
p[FeOH+2] = pK2 - pH + p[Fe+3] = 1.66 + 2pH
(2)
p[Fe(OH)2+] = pK3 - 2pH + p[Fe+3] = 5.14 + pH
(3)
p[Fe(OH)3] = pK4 - 3pH + p[Fe+3] = 12.03
(4)
p[Fe(OH)4-] = pK5 - 4pH + p[Fe+3] = 21.1 - pH
(5)
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Environmental Geochemistry
DM Sherman, University of Bristol
2011/2012
Solubility of FeOOH
(cont.)"
Plotting the p[Fe(OH)n]
(equations 1-5) gives a
map of the saturation
field of FeOOH and the
dominant aqueous Fe
complexes at each pH.
3
2
1
5
4
Notice that [Fe]tot
reaches a minimum at
the pH of seawater.
What is expected of you…"
After this unit, you should be able to:
• Identify processes that will be controlling the chemistry
of an environmental system (esp. aquatic environments).
• Use thermodynamics to predict aqueous
concentrations.
• Use kinetic and mass balance constraints to set up
dynamic models of open systems.
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Environmental Geochemistry
DM Sherman, University of Bristol
2011/2012
Right now you must…"
• Find the unit website!
• Review L1 Chemistry material
• Review the unit factor method!
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