organisms CaCO 3 + H 2 O + CO 2 shallow water

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1 Weathering and Reverse weathering Step I:Weathering of igneous rocks 1. Igneous rocks are mainly composed of Al, Si and O 2 with minor and varying quantities of Na, K, Ca and Mg composing pheldspar minerals 2. these minerals are weathered according to the following equations: CaAl 2 Si 2 O 8 Anorthite Ca 2+ 2KAlSi 3 O 8 + 9H 2 O + 6CO 2 2K + + 8SiO 2 (aq) + 3Al 2 Si 2 O 5 (OH) 4 Potassium feldspar 2Na + (Kaolinite) 2NaAlSi 3 O 8-6HCO 3 Sodium feldspar Igneous rock Rain water Seawater Detritus Step II: Equilibration in ocean 3Al 2 Si 2 O 5 (OH) 4 + 2K + + 2HCO 3-2K(AlSiO 4 )(OH 2 )O 2 (Si 2 O 4 ) + 5H 2 O + 2CO 2 (deep water) kaolinite + seawater illite Ca HCO 3 - organisms CaCO 3 + H 2 O + CO 2 shallow water 2HCl + 2HCO 3 - Volcanisms 3Cl - + 2H 2 O + 2CO 2

2 Step III: Metamorphosis of Clay KAlSi 3 O 8 Potassium pheldspar 2K(AlSiO 4 )(OH 2 )O 2 (Si 2 O 4 ) + Na + + Cl - + 8SiO 2 Heat NaAlSi 3 O 8 Sodium Pheldspar + HCl + 2SiO 2 + AlSi 2 O 5 (OH) Pressure KAl 2 (AlSi 3 O 10 )(OH) 2 Potassium Mica SiO 2 CLAY + interstitial water Granite + Volcanic gases + Quartz + Pyrophyllite: Step IV: left behind in Ocean Na + + Cl -

3 Characteristics of ph of seawater: Average ph of seawater 8.1±0.2 Buffering capacity of a separate liter of seawater is very limited, addition of only 3mmol of hydrochloric acid will lower the ph to less than 3 Marine system is globally buffered and resists any changes due to the addition of natural acids or bases. Three types of acids exist in seawater: 1. Oxiacids: H 2 CO 3 ; HCO ; H 3 BO 3 ; H 2 BO 3 2. Hydrated metallic ions that react with water (doubly or more charged cations) M(H 2 O) n+ x + H 2 O M(H 2 O) (n-1)+ x-1 + H 3 O + 3. Cations of very weak acidity (alkali and alkaline earth elements)

4 ph is therefore buffered by the action of the buffering system: carbonic acid bicarbonate carbonate And to a lesser extent, the action of the buffering system: Boric acid borate This explains the stability of the ph of seawater at 8.1 ± 0.2 (observation 1) The low buffering capacity of an isolated quantity of seawater is due to the fact that bicarbonate concentration of 1l of seawater is only 2.5x10-3 M; addition of 3 mmole of HCl would bring the ph lower than 6

5 How can we explain the contradiction between the limited buffering capacity of seawater and the very limited ph variability? Short term processes 1. Continuous mixing of seawater 2. Biogeochemical cycle of carbon Long term processes 1. Ion exchange between water and aluminosilicate minerals 3 Al 2 Si 2 O 5 (OH) 4(solid) + 4SiO 2 + 2K + + 2Ca 2+ +9H 2 O 2KCaAl 3 Si 5 O 16 (H 2 O) 6(s) + 6H + log k = 6log [H + ] 2log [K + ] - 2log [Ca + ] Buffering capacity of silicates is 2000 times the carbonate system buffering capacity Conclusion: 1. Carbonate system is a short term buffer, 2. In the absence of carbonate, boric acid/borate is the short term buffer 3. Silica is the long term buffering system 4. Ion exchange regulates the concentration of the major cations Na +, K +, Ca 2+ and Mg Silicate minerals control the concentration of dissolved silica in seawater

6 The iron paradox!! Dissolved iron concentration in seawater is higher than that expected from the dissolution of iron oxides! Seawater is an oxidizing environment as oxygen was added to the model ocean (0.027 mol) Most of the oxygen stays in the gaseous form and only 2x10-4 mol l -1 is soluble Dissolved oxygen assures an oxidizing potential expressed by the equation: E 0 = standard Redox potential R = constant = (8314 mv coulomb deg -1 mol -1 ) T = absolute temperature F = Faradays constant (96500 coulomb mol -1 ) RTF -1 ln10 = mv at 25 O C = mv at 0 O C pe = - log e - = E 0 /RTF -1 ln10

7 Calculation of pe The following equilibrium determines the pe of seawater or an ideal system: 0.5 O 2 + 2H + + 2e - H 2 O log k = log k = log (H 2 O) 0.5 log po 2 2log (H + ) 2log (e - ) = log (0.21) + 2 ph + 2 pe = (2 x 8.1) + 2 pe pe = 0.5 (25.5) = 12.5 Any increase of ph is accompanied with a decrease of pe po 2 has weak influence on the pe

8 pe is the master variable in any oxidation-reduction system Influence of pe on the ferric/ferrous system:

9 Concentration of dissolved iron (Fe 2+, Fe 3+ ) is controlled by the pe or the ph of the medium This does not appear true in the presence of solid iron oxides Three equilibria control the oxidation/reduction reactions of iron Fe(OOH) s + H + Fe(OH) 2 + log k = (1) Fe(OOH) 3 + 3H + Fe H 2 O log k = 41.0 (2) e - + Fe 3+ Fe 2+ log k = 13.0 (3) Sillén considered equation 1 as the most important as it gives concentration of Fe(OH) 2 + ions of M at ph of seawater 8.1 However dissolved iron concentration in seawater is higher than M Addition of equations 2&3 gives the concentration of ferrous iron Fe 2+ = M Fe(OOH) 3 + 3H + + e - Fe H 2 O log k = 17 It appears that the concentrations of Fe 2+ and Fe 2+ at equilibrium are lower than the measured iron concentration in seawater Most of the iron present in solution as colloidal iron and iron organic complexes particularly of humic acids