Weathering and Reverse weathering

Step I:Weathering of igneous rocks

1. Igneous rocks are mainly composed of Al, Si and O2 with minor and varying quantities

of Na, K, Ca and Mg composing pheldspar minerals

2. these minerals are weathered according to the following equations:

CaAl2Si2O8

Anorthite Ca2+

2KAlSi3O8 + 9H2O + 6CO2 2K+ + 8SiO2(aq) + 3Al2Si2O5(OH)4

Potassium feldspar 2Na+ (Kaolinite)

2NaAlSi3O8 6HCO3-

Sodium feldspar

Igneous rock Rain water Seawater Detritus

Step II: Equilibration in ocean

3Al2Si2O5(OH)4 + 2K+ + 2HCO3- 2K(AlSiO4)(OH2)O2(Si2O4) + 5H2O + 2CO2 (deep water)

kaolinite + seawater illite

Ca 2+ + 2HCO3- organisms CaCO3 + H2O + CO2 shallow water

2HCl + 2HCO3- Volcanisms 3Cl- + 2H2O + 2CO2

Step III: Metamorphosis of Clay

KAlSi3O8

Potassium pheldspar

Heat NaAlSi3O8

2K(AlSiO4)(OH2)O2(Si2O4) + Na+ + Cl- + 8SiO2 Sodium Pheldspar + HCl + 2SiO2 + AlSi2O5(OH)

Pressure KAl2(AlSi3O10)(OH)2

Potassium Mica

SiO2

CLAY + interstitial water Granite + Volcanic gases + Quartz + Pyrophyllite:

Step IV: left behind in Ocean

Na+ + Cl-

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: H2CO3; HCO3-; H3BO3; H2BO3

-

2. Hydrated metallic ions that react with water (doubly or more charged cations)

M(H2O)xn+

+ H2O M(H2O)x-1(n-1)+ + H3O

+

3. Cations of very weak acidity (alkali and alkaline earth elements)

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

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 Al2Si2O5(OH)4(solid) + 4SiO2 + 2K+ + 2Ca2+ +9H2O 2KCaAl3Si5O16(H2O)

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+, Ca2+ and Mg2+

5. Silicate minerals control the concentration of dissolved silica in seawater

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:

pE = - log e- = E0/RTF

-1ln10

E0

= 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 = 59.15 mV at 25O C

= 54.19 mV at 0O C

Calculation of pE

The following equilibrium determines the pE of seawater or an ideal system:

0.5 O2 + 2H+

+ 2e-

H2O log k = 41.55

log k = log (H2O) – 0.5 log pO2 – 2log (H+) – 2log (e-)

41.55 = 0.0 – 0.5 log (0.21) + 2 pH + 2 pE

41.55 0.0 = 0.34 + (2 x 8.1) + 2 pE

pE = 0.5 (25.5) = 12.5

Any increase of pH is accompanied with a decrease of pE

pO2 has weak influence on the pE

pE is the master variable in any oxidation-reduction system

Influence of pE on the ferric/ferrous system:

Concentration of dissolved iron (Fe2+, Fe3+) 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 = -2.35 (1)

Fe(OOH)3 + 3H+ Fe3+ + 2H2O log k = 41.0 (2)

e-

+ Fe3+ Fe2+ log k = 13.0 (3)

Sillén considered equation 1 as the most important as it gives concentration of Fe(OH) 2

+

ions of 10-10.45 M at ph of seawater 8.1

However dissolved iron concentration in seawater is higher than 10-7.2 M

Addition of equations 2&3 gives the concentration of ferrous iron

Fe(OOH)3 + 3H+ + e- Fe2+ + 2H2O log k = 17

Fe2+ = 10-20 M

It appears that the concentrations of Fe2+ and Fe2+ 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