1. 1 Acid-base reactions and carbonate system

2. 2 Topics for this chapter Acid base reactions and their importance Acid base reactions and their importance Carbonate system and their importance in biogeochemical reactions Carbonate system and their importance in biogeochemical reactions pH range and controlling factors in the natural waters pH range and controlling factors in the natural waters

3. 3 Objectives Better understand acid-base equilibrium in natural waters Better understand acid-base equilibrium in natural waters Better understand carbonate species change and alkalinity measurement Better understand carbonate species change and alkalinity measurement Understand the method to study species distributions in natural waters Understand the method to study species distributions in natural waters

4. 4 DEFINITION: ACIDS AND BASES Bronsted definition Bronsted definition Acid: any substance that can donate a proton Acid: any substance that can donate a proton Base: any substance that can accept a proton Base: any substance that can accept a proton HCl(aq)  H + + Cl - NaOH(aq) + H +  Na + + H 2 O(l) Lewis definition: Lewis definition: Acid: any substance that can accept electrons Acid: any substance that can accept electrons Base: any substance that can donate electrons Base: any substance that can donate electrons H 3 BO 3 (aq) + OH -  B(OH) 4 - H 2 O(l)  H + + OH - H 3 BO 3 (aq) + H 2 O(l)  B(OH) 4 - + H +

5. 5 Acid-Base Acids and bases exist as conjugate acid-base pairs. Every time a Brønsted acid acts as an H + -ion donor, it forms a conjugate base. HA + H 2 O  H 3 O + + A - Every time a base gains an H+ ion, the product is a Brønsted acid, HA. A - + H 2 O  HA + OH -

6. 6 Conjugate acid-base pairs AcidsBases H 3 O + H 2 O H 2 OOH - H 2 CO 3 HCO 3 - HCl Cl -

7. 7 THE HYDRONIUM ION The proton does not actually exist in aqueous solution as a bare H + ion. The proton does not actually exist in aqueous solution as a bare H + ion. The proton exists as the hydronium ion (H 3 O + ). The proton exists as the hydronium ion (H 3 O + ). Consider the acid-base reaction: Consider the acid-base reaction: HCO 3 - + H 2 O  H 3 O + + CO 3 2- Here water acts as a base, producing the hydronium ion as its conjugate acid. For simplicity, we often just write this reaction as: HCO 3 -  H + + CO 3 2-

8. 8 Expression of acid strength The strength of an acid is expressed by the value of the equilibrium constant for its dissociation reaction. The strength of an acid is expressed by the value of the equilibrium constant for its dissociation reaction. Consider: H 2 CO 3  H + + HCO 3 - Consider: H 2 CO 3  H + + HCO 3 - The dissociation constant for this reaction at 25°C is: The dissociation constant for this reaction at 25°C is: This can also be expressed as pK a = 6.35. This can also be expressed as pK a = 6.35. Bicarbonate is considered to be a relatively weak acid. Bicarbonate is considered to be a relatively weak acid.

9. 9 Strong acids Now consider: HNO 3  H + + NO 3 - Now consider: HNO 3  H + + NO 3 - Nitric acid is considered to be a very strong acid; in fact, its pK a is not well defined because the concentration of undissociated acid HNO 3 0 is too small to be measured accurately. Nitric acid is considered to be a very strong acid; in fact, its pK a is not well defined because the concentration of undissociated acid HNO 3 0 is too small to be measured accurately. The conjugate bases of weak acids are strong, and the conjugate bases of strong acids are weak. The conjugate bases of weak acids are strong, and the conjugate bases of strong acids are weak. Thus, NO 3 - is a very weak base, but CO 3 2- is a comparatively strong base. Thus, NO 3 - is a very weak base, but CO 3 2- is a comparatively strong base.

10. 10 K a and pKa HA (aq) + H 2 O (l)  H 3 O + (aq) + A - (aq) The stronger the acids, the larger the K a The stronger the acids, the smaller the pK a pK a = -log K a

11. 11 The larger the pK a, the weaker the acid

12. 12 DISSOCIATION CONSTANTS OF WEAK ACIDS AT 25°C

13. 13 Range of pH values in the natural environment Most natural waters have pH between 4-9. Most natural waters have pH between 4-9. The acids are usually weak, mainly carbonic acid and organic acids (e.g. fulvic and humic acids, formic and acetic acid) The acids are usually weak, mainly carbonic acid and organic acids (e.g. fulvic and humic acids, formic and acetic acid) pH values > 8.5 are rare, occurring only in evaporitic lakes, lakes clogged with photosynthetic plants, and springs discharging from serpentine or ultramafic rocks.

14. 14 Natural Water pH

15. 15 Control of pH on natural reactions The solubility rate of dissolution of most minerals is strongly pH- dependent. Weathering of carbonate, silicate, and alumino-silicate minerals consumes protons and releases metal cations. The solubility rate of dissolution of most minerals is strongly pH- dependent. Weathering of carbonate, silicate, and alumino-silicate minerals consumes protons and releases metal cations. Aqueous acid-base equilibria, including hydrolysis and polymerization. Aqueous acid-base equilibria, including hydrolysis and polymerization. Adsorption, because protons compete with cations and hydroxyl ions compete with anions for adsorption sites. Also, the surface charge of most minerals is pH dependent. Adsorption, because protons compete with cations and hydroxyl ions compete with anions for adsorption sites. Also, the surface charge of most minerals is pH dependent. The formation of metal ligand complexes, because protons compete with metal ions to bond with weak-acid ions, and OH- competes with other ligands that would form complexes. The formation of metal ligand complexes, because protons compete with metal ions to bond with weak-acid ions, and OH- competes with other ligands that would form complexes. Oxidation-reduction reactions, whether abiological or biologically mediated. Oxidation usually produced protons, whereas reduction consumes them. Oxidation-reduction reactions, whether abiological or biologically mediated. Oxidation usually produced protons, whereas reduction consumes them.

16. 16 pH effect on adsorption of metals by ferrihydrite

17. 17 Carbonic acid system: the importance   Carbonic acid is the most abundant acid in natural water systems   Carbonic acid is the acid most responsible for rock weathering.   The pH of most natural waters is controlled by reactions involving the carbonate system   Bicarbonate ion is generally the dominant anion in fresh surface- and ground- waters.   Bicarbonate and carbonate ions are the chief contributors to total alkalinity in natural waters.   An example of acid-base systems in general; the species relationship developed for carbonate equilibrium can be used with little modification for equilibira involving such species as phosphate, sulfide, and silicic acid.

18. 18 Unanswered questions from previous chapters Alkalinity = [HCO 3 – ] + 2[CO 3 2- ] + [OH – ] – [H + ] ≈ [HCO 3 – ] +2[CO 3 2– ] Alkalinity = [HCO 3 – ] + 2[CO 3 2- ] + [OH – ] – [H + ] ≈ [HCO 3 – ] +2[CO 3 2– ] ≈ [HCO 3 – ] [H + ] = [HCO 3 - ] + 2[CO 3 2- ] + [OH - ] [H + ] = [HCO 3 - ] + 2[CO 3 2- ] + [OH - ] When pH<7, [H + ] ≈ [HCO 3 - ] [H + ] ≈ [HCO 3 - ]  Why?

19. 19 THE CO 2 -H 2 O SYSTEM Carbonic acid is a weak acid of great importance in natural waters. The first step in its formation is the dissolution of CO 2 (g) in water according to: CO 2 (g)  CO 2 (aq) At equilibrium we have: Once in solution, CO 2 (aq) reacts with water to form carbonic acid: CO 2 (aq) + H 2 O(l)  H 2 CO 3 0

20. 20 THE CO 2 -H 2 O SYSTEM In practice, CO 2 (aq) and H 2 CO 3 0 are combined and this combination is denoted as H 2 CO 3 *. It’s formation is dictated by the reaction: CO 2 (g) + H 2 O(l)  H 2 CO 3 * For which the equilibrium constant at 25°C is: Most of the dissolved CO 2 is actually present as CO 2 (aq); only a small amount is actually present as true carbonic acid H 2 CO 3 0.

21. 21 THE CO 2 -H 2 O SYSTEM Carbonic acid (H 2 CO 3 *) is a weak acid that dissociates according to: H 2 CO 3 *  HCO 3 - + H + For which the dissociation constant at 25°C and 1 bar is: Bicarbonate then dissociates according to: HCO 3 -  CO 3 2- + H +

22. 22 THE RELATIONSHIP BETWEEN H 2 CO 3 * AND HCO 3 - We can rearrange the expression for K 1 to obtain: This equation shows that, when pH = pK 1 (when pH = 6.35), the activities of carbonic acid and bicarbonate are equal. We can also rearrange the expression for K 2 to obtain: This equation shows that, when pH = pK 2 (when pH = 10.33), the activities of bicarbonate and carbonate ion are equal.

23. 23 BJERRUM PLOTS Plot of the log of the concentrations of various species in a closed CO 2 -H 2 O system as a function of pH. Plot of the log of the concentrations of various species in a closed CO 2 -H 2 O system as a function of pH. The species in the CO 2 -H 2 O system: H 2 CO 3 *, HCO 3 -, CO 3 2-, H +, and OH -. The species in the CO 2 -H 2 O system: H 2 CO 3 *, HCO 3 -, CO 3 2-, H +, and OH -. At each pK value, conjugate acid-base pairs have equal concentrations, and At each pK value, conjugate acid-base pairs have equal concentrations, and At pH < pK 1, H 2 CO 3 * is predominant, and accounts for nearly 100% of total carbonate. At pH < pK 1, H 2 CO 3 * is predominant, and accounts for nearly 100% of total carbonate. pH < 6.35 pH < 6.35 At pK 1 < pH < pK 2, HCO 3 - is predominant, and accounts for nearly 100% of total carbonate. At pK 1 < pH < pK 2, HCO 3 - is predominant, and accounts for nearly 100% of total carbonate. 6.35 < pH < 10.33 6.35 < pH < 10.33 At pH > pK 2, CO 3 2- is predominant. At pH > pK 2, CO 3 2- is predominant. pH < 10.33 pH < 10.33

24. 24 Bjerrum plot showing the activities of inorganic carbon species as a function of pH for a value of total inorganic carbon of 10 -3 mol L -1. In most natural waters, bicarbonate is the dominant carbonate species!

25. 25 Alkalinity end points Alkalinity = [HCO 3 – ] + 2[CO 3 2- ] + [OH – ] – [H + ]

26. 26 Bjerrum plot showing the activities of reduced sulfur species as a function of pH for a value of total reduced sulfur of 10 -3 mol L -1.

27. 27 Bjerrum plot: full solution Closed system, with total carbonate concentration (K 1, K 2 known) Closed system, with total carbonate concentration (K 1, K 2 known) C T = [H 2 CO 3 ]+[HCO 3 - ] + [CO 3 2- ](1) K 1 =[H + ][HCO 3 - ]/[H 2 CO 3 ](2) K 2 =[H + ][CO 3 2- ]/[HCO 3 2- ](3)  Solution: [CO 3 2- ] = C T /  H [HCO 3 - ]=C T [H + ]/K 2  H [H 2 CO 3 ]=C T [H + ] 2 /K 1 K 2  H where:  H = (1 + [H + ]/K 2 +[H + ] 2 /K 1 K 2 )

28. 28 SPECIATION IN OPEN CO 2 - H 2 O SYSTEMS In an open system, the system is in contact with its surroundings and components such as CO 2 can migrate in and out of the system. Therefore, the total carbonate concentration (C T ) is not constant. In an open system, the system is in contact with its surroundings and components such as CO 2 can migrate in and out of the system. Therefore, the total carbonate concentration (C T ) is not constant. In an open system, the solubility of CO 2 increases dramatically with pH, once pH has increased beyond pK 1 In an open system, the solubility of CO 2 increases dramatically with pH, once pH has increased beyond pK 1 At low pH, the solubility of CO 2 is independent of pH. At low pH, the solubility of CO 2 is independent of pH. Let us consider two natural waters Let us consider two natural waters open to the atmosphere, for which p CO 2 = 10 -3.5 atm. open to the atmosphere, for which p CO 2 = 10 -3.5 atm. open to local exchange, for which p CO 2 = 10 -2.0 atm. open to local exchange, for which p CO 2 = 10 -2.0 atm.

29. 29 Plot of log concentrations of inorganic carbon species H + and OH -, for open-system conditions with a fixed p CO 2 = 10 -3.5 atm.

30. 30 Plot of log concentrations of inorganic carbon species H + and OH -, for open-system conditions with a fixed p CO 2 = 10 -2.0 atm.

31. 31 SOURCES OF CO 2 IN NATURAL WATERS When we determine p CO 2 in natural waters, particularly ground waters and soil solutions, values greater than atmospheric are commonly obtained. When we determine p CO 2 in natural waters, particularly ground waters and soil solutions, values greater than atmospheric are commonly obtained. System essentially closed to atmospheric CO 2 (little exchange) System essentially closed to atmospheric CO 2 (little exchange) Respiration by plant roots and microbes consumes organic matter and produces CO 2 : Respiration by plant roots and microbes consumes organic matter and produces CO 2 : CH 2 O + O 2  CO 2 + H 2 O Amount of CO 2 production depends on temperature, soil moisture content, and the amount of organic matter. Amount of CO 2 production depends on temperature, soil moisture content, and the amount of organic matter.

32. 32 Reactions Affecting CO 2 and pH Key: Key: Blue results in pH increase (more alkaline) Blue results in pH increase (more alkaline) Red results in pH decrease (more acidic) Red results in pH decrease (more acidic) CO 2 (g) dissolution, CO 2 (aq) exsolution CO 2 (g) dissolution, CO 2 (aq) exsolution CO 2 (g) + H 2 O  H 2 CO 3 ° CO 2 (g) + H 2 O  H 2 CO 3 ° Photosynthesis, Respiration & aerobic decay Photosynthesis, Respiration & aerobic decay CO 2 (g) + H 2 O  1/6C 6 H 12 O 6 (aq) + O 2 CO 2 (g) + H 2 O  1/6C 6 H 12 O 6 (aq) + O 2 Methane fermentation (anaerobic decay) Methane fermentation (anaerobic decay) C 6 H 12 O 6 (aq) + O 2  CH 4 + H 2 O + CO 2 C 6 H 12 O 6 (aq) + O 2  CH 4 + H 2 O + CO 2 Nitrate uptake and reduction Nitrate uptake and reduction NO 3 - + 2H + + 2CH 2 O  NH 4 + + 2CO 2 + H 2 O NO 3 - + 2H + + 2CH 2 O  NH 4 + + 2CO 2 + H 2 O

33. 33 Reactions Controlling CO 2 and pH Carbonate mineral Dissolution or precipitation Carbonate mineral Dissolution or precipitation CaCO 3 (calcite ) + H +  Ca 2+ + H 2 O+ CO 2 CaCO 3 (calcite ) + H +  Ca 2+ + H 2 O+ CO 2 Sulfate reduction Sulfate reduction 2CH 2 O + SO 4 2- + H +  HS - + 2H 2 O + 2CO 2 2CH 2 O + SO 4 2- + H +  HS - + 2H 2 O + 2CO 2 Denitrification Denitrification 5CH 2 O + 4NO 3 - + 4H +  2N 2 + 5CO 2 + 7H 2 O 5CH 2 O + 4NO 3 - + 4H +  2N 2 + 5CO 2 + 7H 2 O Chemical weathering of Al-silicate weathering Chemical weathering of Al-silicate weathering KAlSi 3 O 8 + 2CO 2 + 11H 2 O  Al 2 Si 2 O 5 (OH) 4 + 2K + + 2 HCO 3 - KAlSi 3 O 8 + 2CO 2 + 11H 2 O  Al 2 Si 2 O 5 (OH) 4 + 2K + + 2 HCO 3 -

34. 34 CO 2 in Natural Settings Time of day Time of day Higher p CO2 values occur in surface waters at night because of respiration and aerobic decay and groundwater inflow. Higher p CO2 values occur in surface waters at night because of respiration and aerobic decay and groundwater inflow. Lower p CO2 values occur in surface waters during the day because of photosynthesis. Lower p CO2 values occur in surface waters during the day because of photosynthesis. Time of year Time of year Soil p CO2 values are highest during the growing season because of plant respiration. Soil p CO2 values are highest during the growing season because of plant respiration. Consequently, shallow groundwaters will have their highest p CO2 values during the growing season. Consequently, shallow groundwaters will have their highest p CO2 values during the growing season.

35. 35 Compute the pH of an aqueous system Step 1: List species in solution Step 2: Identify how concentrations of the species depend on each other Step 3: Establish the mass balance Step 4: Establish the charge balance

36. 36 Example: HAC in water CH 3 COOH  CH 3 COO - + H + K a = 10 -4.75 Species: Species: CH 3 COOH, CH 3 COO -, H +, OH - Mass balance: C T = [HAc] + [Ac - ] Charge balance: [H + ] = [OH - ] + [Ac - ] K a =[Ac - ][H + ]/[HAc]=10 -4.75 Kw=[H + ][OH - ] = 10 -14 With given C T (concentration of HAC you made), we have four equations, four unknowns, we can compute [H + ] hence pH

37. 37 pH of the sea pH of the sea can be computed following the above procedures pH of the sea can be computed following the above procedures Basically, pH of the sea is controlled by the carbonate species Basically, pH of the sea is controlled by the carbonate species The actual computation is more complicated and beyond this class The actual computation is more complicated and beyond this class