Formation of the Elements

Composition of Earth

Chemical EquilibriumExists 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, ingeneral, the farther a system is from equilibrium, the more rapidly it will movetoward equilibrium; however, it is generally not possible to calculate reaction rates from thermodynamic data.

Reaction Rates/Equilibrium

Acid-Base EquilibriaBronsted-Lowry definition: acid donates H+; base accepts H+

In aqueous systems, all acids stronger that H2Ogenerate excess H+ ions (or H3O+); all bases strongerthan H2O generate excess OH-

2

3

3

3

Acid-Base

Many reactions influence pHPhotosynthesis 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)

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.

Redox

The oxidation state of an atom is defined with the followingconvention:•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 oxidationstate -2 and hydrogen is in oxidation state of +1 (except forperoxide 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 onthe ion.

For Fe2+, Fe is in +2 oxidation state.

RedoxRedox 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

Why do we care about redox rxns?

Oxidation state can impact1. Sorption/desorption2. Solubility3. Toxicity4. Biological uptakeetc.

Measure of oxidation-reduction potential gives us info about chemical species present and microbes we may find.

Biogeochemistry

QuickTime™ and aTIFF (Uncompressed) decompressor

are needed to see this picture.

Nitrification ammonia→ nitrite → nitrate

Denitrificationnitrate → nitrite → nitric oxide → nitrous oxide → N2

N FixationN2 →ammonia

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.

12C 13C 14C

14N 15N

10B 11B

• 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)

How did all this stuff get here?

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

Stable Isotopes

• Rayleigh fractionation: light isotopes evaporate more easily, and heavy isotopes rain out more quickly

= {(Rsample – Rstandard) / Rstandard} x 103

Stable Isotope Examples

18Ocarbonate in forams depends on 18Oseawater as well as T, S

18Oseawater depends on how much glacial ice there is– Glacial ice is isotopically

light b/c of Rayleigh fract.– More ice means higher

18Oseawater

Stable Isotope Examples

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)

• 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

Stable Isotope Examples

Boron isotopes measured in forams used for paleo-pH 11B depends on pH(Gary Hemming)

Nitrogen isotopes used for rapid temp. changes in ice cores 15N 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

Stable Isotopes