origin of elements and its distibution in earth...

1. Formation of the Elements

2. Composition of Earth

4. 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.

5. Reaction Rates/Equilibrium

6. 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
3
3

7. 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)

8. 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.

9. 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.

10. 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

11. Why do we care about redox rxns?
O
Measure of oxidation-reduction
potential gives us info about
chemical species present and
microbes we may find.

12. Biogeochemistry

14. QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
Nitrification
ammonia→ nitrite → nitrate
Denitrification
nitrate → nitrite → nitric oxide → nitrous
oxide → N2
N Fixation
N2 →ammonia

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

16. 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)

17. 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

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

19. Stable Isotope Examples
∀ δ18Ocarbonate in forams
depends on δ18Oseawater as
well as T, S
∀ δ18Oseawater depends on how
much glacial ice there is
– Glacial ice is isotopically
light b/c of Rayleigh fract.
– More ice means higher
δ18Oseawater

20. Stable Isotopes
• C in organic matter, fossil fuels, and hydrocarbon gases is depleted in
13
C ==> 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)

21. Stable Isotope Examples
• Stable isotopes can also tell you about
biology
• Organisms take up light isotopes
preferentially
• So, when an organism has higher
δ30Si, it means that it was feeding from a
depleted nutrient pool

22. Stable Isotopes
Boron isotopes measured in forams used for paleo-pH
δ11B depends on pH
(Gary Hemming)
Nitrogen isotopes used for rapid temp. changes in ice cores
δ15N 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