1. Environmental Geochemistry
January 26, 2007

2. What is geochemistry? The study of
-chemical composition of the Earth and other planets
-chemical processes and reactions that govern the composition of rocks and soils
-the cycles of matter and energy that transport the Earth's chemical components in time and space
-and their interaction with the hydrosphere and the atmosphere.
wikipedia

3. Outline of Topics Formation of the elements Composition of Earth
Aqueous Solutions
Chemical Equilibrium
Acid-Base Equilibria
Redox
Biogeochemistry
Stable Isotopes
(with comments on weathering, sorption, pollution…)

4. Formation of the Elements
Fig 1
-Universe produced by Big Bang 13.7 billion years ago - based on measurements of the cosmic microwave background and red shift
-protons and neutrons combined to form light nuclei (H and He)
-fusion of H to form He nuclei in stars
-once H in star used up, He nuclei fusion to form C, O
-C fusion to form Si (also Na, Mg, Ne)…
-Si burning to form elements near Fe
-can’t get elements heavier than Fe (mass 56) with direct fusion - formation proceeds by supernovae and neutron capture
Fig 2
-H and He are most abundant in solar system
-abundances of Li, Be, and B anomalously low
-abundances of elements whose atomic numbers greater than 6 (C) decrease exponentially with increasing Z
-Fe exceptionally abundant, abundance of F lower than expected
-Tc and Pm do not occur naturally in solar system (all isotopes unstable)
-abundance of Pb somewhat greater than expected (stable end product of decay of U and Th)
-U has lowest abundance of any element in solar system (except as noted above)
-abundances of elements between Bi through Th low because these elements are unstable daughters of decay of naturally occurring isotopes of U and Th

5. Composition of Earth Accretion Models for Earth Formation
Homogenous Accretion (COLD):
Accretion began after nebula was cool enough to have both
metallic and silicate particles. Accumulation of planetesimals
formed a well-mixed earth.
Gravitational collapse released heat causing partial melting.
Fe sank into the core (the “iron catastrophe”; LIL’s rose to
the crust
Heterogenous Accretion (HOT):
Proto-Earth accreted from dense Fe-rich particles that first
condensed from solar nebula.
After cooling of the nebula, Si-O rich particles formed which
then accreted later to form the mantle and crust.

6. Chalcophile elements
- high bonding affinity—usually in the form of covalent bonds—with sulfur, and are, accordingly, usually abundant in sulfides.
- exhibit a bonding affinity with selenium, tellurium, arsenic, and antimony
-When sulfur is abundant, chalcophile elements readily form sulfide minerals as they precipitate from the magma. This process partially explains the formation of extensive deposits of iron-nickel-copper sulfides.
Lithophiles
- high bonding affinity with oxygen.
- affinity to form ionic bonds and are represented by silicates (silicon and oxygen) in the crust and mantle
-Other lithophile elements include magnesium, aluminum, sodium, potassium, iron, and calcium.
Siderophiles
-exhibit a weak affinity to both oxygen and sulphur.
-affinity for iron
-high solubility in molten iron.
-generally have a low reactivity
-affinity to form metallic bonds.
-most often found in their native state.
-Not abundant in the core or mantle
-most siderophiles are thought to be richest at Earth's core.
-Platinum (Pt) group metals, including Ruthium (Ru), Rhodium (Rd), Palladium (Pd), Osmium (Os), and Iridium (Ir), show exhibit a strong siderophile tendency.
Atmophiles are a related fourth class of elements characterized by their ability to form van der Waals bonds. Atmophiles are also highly volatile.

7. Aqueous Solutions Water is special
Compared with hydrides of other elements in same column on periodic table water is odd. It has a larger liquid temperature range and higher melting and boiling point than expected. Dipole moment, H bonding. Because of H bonding has a much higher boiling point than liquids of a similar mass (CH4).
Max density at 4 degrees C. - influences stratification and mixing of lakes.

8. Like dissolves like
Water able to dissolve huge quantities of salts, therefore and important factor in transporting substances in nature
More than 300g NaCl can be dissolved in 1kg of water

9. TDS - total dissolved solids - combined content of all inorganic and organic substances contained in a liquid which are present in a molecular, ionized or micro-granular (colloidal sol) suspended form, go through 2um filter
In solutions with high ionic strength you start to see deviations - Have to use activities (effective concentration) instead of concentrations. Non polar solutes lead to less deviation from ideal behavior than ions (for ions the electrostatic interactions are greater).
Activity of water in dilute aqueous solutions is 1. Activity in sea water 0.98.
Ionic Strength
I = 1/2 ∑mz2

10. Chemical Equilibrium
Exists when a system is in a state of minimum energy (G)
- Often not completely attained in nature (e.g., photosynthesis leaves products out of chemical equilibrium)
- A good approximation of real world
Gives direction in which changes can take place (in the absence of energy input.)
Systems, including biological systems, can only move toward equilibrium.
-Gives a rough approximation for calculating rates of processes because, in
general, the farther a system is from equilibrium, the more rapidly it will move
toward equilibrium; however, it is generally not possible to calculate reaction rates from thermodynamic data.

11. Reaction Rates/Equilibrium

12. Acid-Base Equilibria
Bronsted-Lowry definition: acid donates H+; base accepts H+
In aqueous systems, all acids stronger that H2O
generate excess H+ ions (or H3O+); all bases stronger
than H2O generate excess OH- 2 3
pH - measure of proton activity
HA --> A- + H+
pH = pKa + Log [A-]/[HA]
pH = -Log[H+]

13. Acid-Base Many reactions influence pH pH influences many processes
Photosynthesis and respiration are acid base reactions.
aCO2(g) + bNO3- + cHPO42- + dSO42- + f Na+ + gCa2+ + hMg2+ iK+ + mH2O + (b + 2c + 2d -f -2g - 2h - i)H+<-----> {CaNbPcSdNafCagMghKiH2Om}biomass + (a + 2b)O2
Oxidation reactions often produce acidity.
Reduction reactions consume acidity
Oxidation often produces acidity, reduction consumes acidity.
Photosynthesis and respiration - also acid base reactions - photosynthesis consumes acidity, respiration produces acidity.
Decrease in pH also accelerates weathering processes - some of the most common elements in rocks (Al3+ and Fe3+) become much more soluble at lower pH.
In forest ecosystems etc lower pH accelerates leaching of base cations from soil which can negatively impact forest growth. H+ will exchange for other cations - Na+, K+, etc.
pH can also influence sorption properties of species - above isoelectric point for minerals there is net negative surface charge, below there is net positive surface charge. H2PO4- sorbs less as pH increases because of electrostatic interactions with surface.
pH influences many processes
-weathering (Fe and Al more soluble at lower pH)
-cation exchange (leaching of base cations from soil due to acid rain)
-sorption (influences surface charge on minerals and therefore what sticks to them)

14. Acid-Base Alkalinity ≈ ANC
People often only consider carbonate alkalinity but total alkalinity also includes NH3, [B(OH)4]-, and [HPO4]2-
Most natural waters have positive values for ANC.
Waters impacted by acid mine drainage may have negative ANC.
if a process consumes Ca2+ without affecting the concentration of any other base cation or acid anion, then charge balance must be maintained through release of H+ or through consumption of OH-
Alkalinity = ∑(base cations) - ∑(strong acid anions)
Any process that affects the balance between base cations and acid anions must affect alkalinity.

15. Redox The oxidation state of an atom is defined with the following
convention:
The oxidation state of an atom in an elemental form is 0.
In O2, O is in the 0 oxidation state.
When bonded to something else, oxygen is in oxidation
state -2 and hydrogen is in oxidation state of +1 (except for
peroxide and superoxide).
In CO32-, O is in -2 state, C is in +4 state.
The oxidation state of a single-atom ion is the charge on
the ion.
For Fe2+, Fe is in +2 oxidation state.

16. Redox
Redox reactions tend to be slow and are often out of thermodynamic equilibrium - but life exploits redox disequilibrium.
Oxidation - lose electrons
Reduction - gain electrons
Rxns occur as redox couples - don’t have free floating electrons
Oxidation reaction examples
Fossil fuel --> CO2 (C as 0 to C as +4)
Fe --> rust (Fe as 0 to Fe as +3)
Reduction rxn examples
SO > H2S (gas that smells like rotten eggs, S as +6 to S as -2)
Fe was oxidized, Mn was reduced

17. Why do we care about redox rxns?
Oxidation state can impact
Sorption/desorption
Solubility
Toxicity
Biological uptake
etc.
Measure of oxidation-reduction potential gives us info about chemical species present and microbes we may find.
(relative strengths of oxidant or reductant depend on pH)
1- (related to solubility) As oxidation state influence sorption to Fe oxides.
Presence or absence of Fe oxides influences sorption of Zn2+ and Pb2+.
2- As III more soluble than As V.
Oxidized U more soluble than reduced U.
Cr VI more soluble than Cr III.
Under reducing conditions Hg forms organometallic complexes which are mobile (and toxic).
Cr VI more toxic. As III more toxic.
4. NH3 (reduced N) easier to incorporate into bacterial biomass. Can take up NO3- but have to expend energy to convert it to reduced N.

18. Accumulation of O2 in the Atmosphere
Fe2+ = Fe(II) = slightly soluble in sea water with no O2 present
Add O2 - oxidizes Fe(II)-->Fe(III)
Very small [O2] required
Fe3+ = Fe(III) = extremely insoluble in water
Essentially all of the oxygen in the atmosphere came from photosynthesis
Fe(III) many orders of magnitude less soluble than Fe(II)
Large iron deposits formed between 2.5 and 2 billion years ago, require very low but not zero free oxygen in the atmosphere to form
Accumulation of oxygen in atmosphere requires burial of significant amount of organic carbon - stashed as coal, oil, tar, shales, methane
Detailed accumulation history of oxygen not known, these pictures represent a possibility
Evidence for oxygen accumulation also in S isotopes (mass independent fractionation)

19. Biogeochemistry

21. ammonia→ nitrite → nitrate Denitrification N Fixation
nitrate → nitrite → nitric oxide → nitrous oxide → N2
N Fixation
N2 →ammonia

22. 

23. What is an isotope?
Isotope- line of equal Z. It has the same # protons (ie. they are the same element) but a diff. # of neutrons.
14N
15N
12C
13C
14C
10B
11B

24. How did all this stuff get here?
4 types of isotopes, based on how they formed:
Primordial (formed w/ the universe)
Cosmogenic (made in the atmosphere)
Anthropogenic (made in bombs, etc)
Radiogenic (formed as a decay product)

25. Stable Isotopes
Light isotopes are fractionated during chemical reactions, phase changes, and biological reactions, leading to geographical variations in their isotopic compositions
FRACTIONATION: separation between isotopes on the basis of mass (usually), fractionation factor depends on temperature
Bonds between heavier isotopes are harder to break
Heavy molecules have stronger covalent bonds, less reactive

26. Stable Isotope Examples
Rayleigh fractionation: light isotopes evaporate more easily, and heavy isotopes rain out more quickly
As you move inland precip is more depleted in heavy isotopes.
Temp dependence of 18O and D in precipitated snow in polar regions used to interpret the isotope composition of ice cores recovered from ice sheets in Greenland and Antarctica and ice cores from high mountains. Taken together, these ice cores contain a record of climatic conditions during and subsequent to the Wisconsin glaciation, not only in the polar regins but also at midlatitudes in different places of the Earth.- many complications in interpreting these records though.
d = {(Rsample – Rstandard) / Rstandard} x 103

27. Stable Isotope Examples
d18Ocarbonate in forams depends on d18Oseawater as well as T, S
d18Oseawater depends on how much glacial ice there is
Glacial ice is isotopically light b/c of Rayleigh fract.
More ice means higher
d18Oseawater
Also use O isotopes for thermometry of silicates and oxides - rock forming silicates and oxide minerals fractionate O isotopes when they crystallize from silicate melts or form by recrystallization of volcano-sed rocks during regional metamorphism

28. Stable Isotopes
C in organic matter, fossil fuels, and hydrocarbon gases is depleted in 13C ==> photosynthesis
used as an indicator of their biogenic origin and as a sign for the existence of life in Early Archean time (~ 3.8 billion years ago)
N isotopic composition of groundwater strongly affected by isotope fractionation in soils plus agricultural activities (use of N-fertilizer and discharge of animal waste)
Particulate matter in ocean enriched in 15N by oxidative degradation as particles sink through water column
Used for mixing and sedimentation studies
S isotopes fractionated during reduction of SO42- to S2- by bacteria
didn’t become important until after ~2.35 Ga when photosynthetic S-oxidizing bacteria had increased sulfate concentration in the oceans sufficiently for anaerobic S-reducing bacteria to evolve (photosynthesis preceded S-reduction which was followed by O respiration)
C isotope fractionation during photosynthesis - different types of plants have different levels of 13C depletion - C3 plants fractionate more than C4 plants
Fractionation of N isotopes during fixation, nitrification, and denitrification
Bacterial reduction of S enriches H2S in 32S by 50 per mil or more wrt sulfate
In general - can be used to tell us about biological and earth processes

29. Stable Isotope Examples
Stable isotopes can also tell you about biology
Organisms take up light isotopes preferentially
So, when an organism has higher
30Si, it means that it was feeding from a depleted nutrient pool

30. Stable Isotopes Boron isotopes measured in forams used for paleo-pH
d11B depends on pH
(Gary Hemming)
Nitrogen isotopes used for rapid temp. changes in ice cores
d15N depends on temp. gradient in firn
(Jeff Severinghaus)
Stable isotopes are also used to study magmatic processes, water-rock interactions, biological processes and anthropology and various aspects of paleoclimate
B isotopes can also give evidence of wastewater

31. References http://mineral.gly.bris.ac.uk/Geochemistry/
(WM White Geochemistry Ch9 - Stable Isotopes)
Isotopes: Principles and Applications - Faure & Mensing
How to Build a Habitable Planet - Wally Broecker