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Solubility of CO 2 and Carbonate Equilibrium CO 2 (g) ↔ CO 2 (aq) ↔ H 2 CO 3 ↔ HCO 3 - ↔ CO 3 2- CO 2 (g) ↔ CO 2 (aq) + H 2 O ↔ H 2 CO 3 H 2 CO 3 ↔ HCO.

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Presentation on theme: "Solubility of CO 2 and Carbonate Equilibrium CO 2 (g) ↔ CO 2 (aq) ↔ H 2 CO 3 ↔ HCO 3 - ↔ CO 3 2- CO 2 (g) ↔ CO 2 (aq) + H 2 O ↔ H 2 CO 3 H 2 CO 3 ↔ HCO."— Presentation transcript:

1 Solubility of CO 2 and Carbonate Equilibrium CO 2 (g) ↔ CO 2 (aq) ↔ H 2 CO 3 ↔ HCO 3 - ↔ CO 3 2- CO 2 (g) ↔ CO 2 (aq) + H 2 O ↔ H 2 CO 3 H 2 CO 3 ↔ HCO H + HCO 3 - ↔ CO H + H 2 O ↔ H + + OH - “Carbonic acid”

2 Solubility of CO 2 and Carbonate Equilibrium 1. K H = [H 2 CO 3 ]/P CO2 = 3 x M atm -1 = M atm K 1 = [HCO 3 - ][H + ]/[H 2 CO 3 ]= 9 x M = M 3. K 2 = [CO 3 2- ][H + ]/[HCO 3 - ] = 4.5 x M = M 4. K w = [H + ][OH - ] = M 2 Note: Values of these equil. constants are sensitive to temperature and ionic strength of the solution; these values are appropriate to seawater. CO 2 (g) ↔ CO 2 (aq) ↔ H 2 CO 3 ↔ HCO 3 - ↔ CO 3 2- KHKH K1K1 K2K2

3 CO 2 Partitioning Atmosphere - Ocean pH of Natural Waters So we want to understand what controls pH… KKC Box Fig. 8-2

4 pH of Natural Waters General concept… At least six unknowns: H +, OH - P CO2 H 2 CO 3, HCO 3 -, CO 3 2-  Need at least six equations: Equilibrium expressions Typically, constraint on either P CO2 (“open system”) or total moles carbon (“closed system”) Charge balance;  n[i n+ ] =  m[j m- ]

5 pH of Natural Waters Pure water in contact with atmosphere Six unknowns Equilibrium expressions (4 equations) P CO2 = 3.5 x atm (1 more equation) Charge balance (6th equation): [H + ] = [HCO 3 - ] + 2[CO 3 2- ] + [OH - ] Strategy: Rewrite charge balance equation in terms of [H + ] and known quantities…

6 pH of Natural Waters Example: Pure water in contact with atmosphere [H + ] = K 1 K H P CO2 /[H + ] + 2K 1 K 2 K H P CO2 /[H + ] 2 + K w /[H + ] Can solve rigorously for [H + ]. Alternatively, make a simplifying assumption: [CO 3 2- ] << [HCO 3 - ] In this case: [H + ] = K 1 K H P CO2 /[H + ] + K w /[H + ] or [H + ] 2 = K 1 K H P CO2 + K w This is easily solved: For pure K 1 = 4.45 x M; K H = 3.39 x M/atm  [H+] = 2.4 x 10-6 M pH = 5.62  “Acid rain” is a term applied to pH < 5

7 pH of Natural Waters Assess simplifying assumption… Is it fair to assume [CO 3 2- ] << [HCO 3 - ]? K 2 = [CO 3 2- ][H + ]/[HCO 3 - ] = (pure water, 25°C) So: [CO 3 2- ]/[HCO 3 - ] = /[H + ] Clearly, [CO 3 2- ]/[HCO 3 - ] > i.e., as long as pH << This is true in most natural waters

8 pH of Natural Waters Alkalinity Consider charge balance: [H + ] = [HCO 3 - ] + 2[CO 3 2- ] + [OH - ] By definition, at pH = 7, [H + ] = [OH - ] So, if [HCO 3 - ] or [CO 3 2- ] are at all comparable to [OH], [OH - ] must be less than [H + ]. This would imply pH 7. So how does this work… How do we get pH > 7??? Must add cations…

9 pH of Natural Waters Alkalinity Imagine we dissolve some CaCO 3 in the system Now: 2[Ca 2+ ] + [H + ] = [HCO 3 - ] + 2[CO 3 2- ] + [OH - ] In this case, [H + ] is free to have lower values (pH > 7) as long as [Ca 2+ ] is present

10 Consider CO 2 Dissolution (again) P CO2 = [H 2 CO 3 ]/K H = [H + ][HCO 3 - ]/(K 1 K H ) = (K 2 [HCO 3 - ]/[CO 3 2- ])([HCO 3 - ]/(K 1 K H )) = K 3 [HCO 3 - ] 2 /[CO 3 2- ]; (K 3 =K 2 /(K 1 K H )) So what?

11 CO 2 Dissolution, HCO 3 - and CO 3 2- P CO2 = K 3 [HCO 3 - ] 2 /[CO 3 2- ] This is the equilibrium expression for the reaction: CO 2 + CO H 2 O ↔ 2HCO 3 - i.e., dissolution of CO 2 consumes CO 3 2-, produces HCO 3 - ; Note that dissolution of CO 2 itself does not affect Alkalinity (gain 2 moles HCO 3 - for every CO 3 2- ) Note also that capacity for CO 2 uptake determined by [CO 3 2- ]

12 CO 2 Dissolution, HCO 3 - and CO 3 2- Another perspective CO 2 + H 2 O ↔ H 2 CO 3 H 2 CO 3  H + + HCO 3 - Net direction if pH >~ 6 H + + CO 3 2-  HCO 3 - Net direction if pH <~ 10 CO 2 + CO H 2 O  2HCO 3 -

13 CO 2 Dissolution, HCO 3 - and CO 3 2- Alk ~ [HCO 3 - ] + 2[CO 3 2- ] ΣCO 2 = [H 2 CO 3 ] + [HCO 3 - ] + [CO 3 2- ] ~ [HCO 3 - ] + [CO 3 2- ] P CO2 = K 3 [HCO 3 - ] 2 /[CO 3 2- ] Algebra… [CO 3 2- ] = Alk - ΣCO 2 [HCO 3 - ] = 2 ΣCO 2 - Alk P CO2 = K 3 (2 ΣCO 2 - Alk) 2 /(Alk - ΣCO 2 ) i.e., P CO2 is controlled by Alk and ΣCO 2

14 CO 3 2- in Ocean Didn’t we say [HCO 3 - ] >> [CO 3 2- ]? Yes, but that doesn’t mean CO 3 2- is irrelevant: Today: [CO 3 2- ] ~ 1/50 [HCO 3 - ] As long as there is excess CO 3 2-, ocean can take up CO 2 without change in alkalinity (or pH). Figure 11-3 Distribution of dissolved carbon species in seawater as a function of pH. Average oceanic pH is about 8.2. The distribution is calculated for temperature of 15 o C and a salinity of 35‰. The equilibrium constants are from Mehrbach et al. (1973).

15 * Reactions are quite congruent, structure not attacked; simple reactions. * Acid Precipitation: 2CaCO 3 + 2H + + SO 4 =  2Ca SO 4 = + 2HCO 3 - * There are natural acids ie. In rain CO 2 in the atmosphere (pH = 5.6) CO 2 + H 2 O  H 2 CO 3  H + + HCO 3- PCO 2 ~ 350 ppm  pH = 5.6 * Rate of carbonate weathering is about 100 times the rate of silicate weathering. ie. Florida- formation of karst due to the dissolving of calcite- wouldn't see that with silica.

16 pH of Natural Waters Alkalinity Not all cation sources will work this way… Imagine we dissolve some NaCl into the system Now: [Na + ] + [H + ] = [HCO 3 - ] + 2[CO 3 2- ] + [Cl - ] + [OH - ] But: [Na + ] = [Cl - ], right? So, no effect on charge balance equation To cope: Distinguish between “conservative” and “nonconservative” ions…

17 pH of Natural Waters Alkalinity Conservative ions: Ions whose concentrations are not affected by pH (or pressure or temperature; not important variables here) Examples: Ca 2+, Na +, NO 3 -, K +, Cl -, etc. Nonconservative ions: Ions whose concentrations are affected by pH Examples: CO 3 2-, HCO 3 -, NH 4 +, B(OH) 4 -, H +, OH - “Alkalinity” ≡  n[i n+ ] -  m[j m- ] where i and j are only conservative ions; alkalinity is what’s left over after these are accounted for. Units: equiv./liter

18 pH of Natural Waters Alkalinity If we consider only HCO 3 -, CO 3 2-, OH - and H + and conservative ions, then we may write:  n[i n+ ] + [H + ] = [HCO 3 - ] + 2[CO 3 2- ] + [OH - ] +  m[j m -]  n[i n+ ] -  m[j m- ] = [HCO 3 - ] + 2[CO 3 2- ] + [OH - ] - [H + ] Alkalinity = [HCO 3 - ] + 2[CO 3 2- ] + [OH - ] - [H + ] Typically, Alkalinity ~ [HCO 3 - ] + 2[CO 3 2- ] ≡ Alk carb For seawater, Alkalinity ~ 2.3 x equiv/liter

19 pH of Natural Waters Alkalinity Alkalinity of seawater allows it to dissolve more CO 2 Higher alkalinity leads to lower [H + ] and higher pH Any reaction that introduces [H + ] lowers alkalinity e.g., NH O 2 → NO H 2 O + 2H + Any reaction that raises [CO 3 2- ] or [HCO 3 - ] raises alkalinity – e.g., CaCO 3 + H 2 O + CO 2 → Ca HCO 3 -

20 Distribution of major species of dissolved inorganic carbon at 20 o C Figure 9.1 Fetter, Applied Hydrogeology 4 th Edition

21 Organic acids (e.g.. Oxalic acid) * Organic material breaks down and releases acids ( pH ~ 5 ). These natural acids play an important role in weathering in absence of human activity; behaves very much like carbonic acid 4H 2 C 2 O 4 (oxalic acid) + 2O 2  8CO 2 + 4H 2 O H 2 CO 3 - -

22 Carbonate Weathering Simplified reactions (removes H + ions) CaCO 3 + H 2 CO 3  Ca++ + 2HCO 3 -  pH ~ of water with bicarbonate; about neutral (Calcite) (Carbonic acid) (bicarbonate) large amounts of calcite is the reason, in general, why places such as Rochester and the Mid-Eastern states don't have problems with acid rain. Problem of acid rain: Determined basically by the presence of absence of calcite- Tells how sensitive an area will be to acid. Calcite reacting to acid

23 Distribution of major species of dissolved inorganic carbon at 20 o C Figure 9.1 Fetter, Applied Hydrogeology 4 th Edition

24 Concentration of calcium in equilibrium with calcite as a function of P C0 2 in the system of CaCO 3 -CO 2 -H 2 0 at 25 o C and 1 atm total pressure. Figure 3-4. Drever, The Geochemistry of Natural Waters

25 Concentration of calcium in equilibrium with calcite as a function P C0 2 and Na + concentration in the system of CaCO 3 - Na 2 CO 3 -CO 2 -H 2 0 at 25 o C and 1 atm total pressure. Concentration of calcium in equilibrium with calcite as a function P C0 2 and Cl - concentration in the system of CaCO 3 - CaCl 2 -CO 2 -H 2 0 at 25 o C and 1 atm total pressure. Figure 3-5,6. Drever, The Geochemistry of Natural Waters

26 Changes in composition of carbonated water as it equilibrates with calcite when the system is either open or closed to exchange of CO2 gas. Initial P C0 2 values of and atm. Figure 3-8. Drever, The Geochemistry of Natural Waters

27 Florida karst formation

28 Solution and collapse features of karst and karren topography


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