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Environmental Geochemistry

Environmental Geochemistry. January 26, 2007. 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

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Environmental Geochemistry

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  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.

  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

  5. Composition of Earth

  6. Aqueous Solutions Water is special

  7. Ionic Strength I = 1/2 ∑mz2

  8. 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.

  9. Reaction Rates/Equilibrium

  10. 2 3 3 3 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-

  11. Acid-Base Many reactions influence pH 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 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)

  12. Acid-Base Alkalinity ≈ ANC Alkalinity = ∑(base cations) - ∑(strong acid anions) Any process that affects the balance between base cations and acid anions must affect alkalinity.

  13. 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.

  14. 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 Fe was oxidized, Mn was reduced

  15. 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.

  16. 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

  17. Biogeochemistry

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

  19. 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

  20. 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)

  21. 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

  22. Stable Isotope Examples • Rayleigh fractionation: light isotopes evaporate more easily, and heavy isotopes rain out more quickly d= {(Rsample – Rstandard) / Rstandard} x 103

  23. 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

  24. Stable Isotopes • C inorganic 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)

  25. 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

  26. 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

  27. References http://mineral.gly.bris.ac.uk/Geochemistry/ http://mineral.gly.bris.ac.uk/envgeochem/ http://www.soest.hawaii.edu/krubin/gg425-sched.html http://geoweb.tamu.edu/courses/geol641/notes.html http://www.imwa.info/Geochemie/Chapters.HTML (WM White Geochemistry Ch9 - Stable Isotopes) Isotopes: Principles and Applications - Faure & Mensing How to Build a Habitable Planet - Wally Broecker

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