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PH and Chemical Equilibrium. Acid-base balance Water can separate to form ions H + and OH - In fresh water, these ions are equally balanced An imbalance.

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Presentation on theme: "PH and Chemical Equilibrium. Acid-base balance Water can separate to form ions H + and OH - In fresh water, these ions are equally balanced An imbalance."— Presentation transcript:

1 pH and Chemical Equilibrium

2 Acid-base balance Water can separate to form ions H + and OH - In fresh water, these ions are equally balanced An imbalance produces an acidic or basic (alkaline) solution –Acid releases H + –Base combines with H +

3 pH scale measures the concentration of H + ions in solution. The more H + ions there are, the more acidic the solution is. The scale is logarithmic so a change of 1 pH unit is a 10-fold change in H + concentration (10 0 = 1; 10 -1 = 0.1; 10 -2 = 0.001, etc) pH = -log 10 [H + ] Average pH of seawater is ~ 8

4 Why is the ocean slightly basic or alkaline? There is a large amount of CO 2 in the ocean so shouldn’t it be acidic if the CO 2 combines with H 2 O to form carbonic acid? CO 2 is actually present in several different forms in water.

5 CO 2 Soluble in water Ocean is a big reservoir for CO 2 (only bigger reservoir is sediments and sed. rocks) The total amount of CO 2 in seawater is about 60 times the amount of CO 2 in the atmosphere Because it forms a variety of chemical species in water, more CO 2 can dissolve than is predicted by gas solubility alone! Two major processes affect CO 2 in the ocean –Addition and removal by organisms –Carbonate mineral precipitation & dissolution

6 CO 2 in seawater CO 2 + H 2 O  H 2 CO 3  HCO 3 - + H +  CO 3 2- + 2 H + 1: CO 2 + H 2 O  H 2 CO 3 (carbonic acid) Carbonic acid rapidly dissociates to form ions 2: H 2 CO 3  HCO 3 - (bicarbonate) + H + 3: HCO 3 - + H +  CO 3 2- (carbonate) + 2 H + Some of the bicarbonate combine with H + ions to form carbonate **At a given pH, CO 2, H 2 CO 3, HCO 3 -, CO 3 2- and H + are in equilibrium

7 When you add acid to seawater CO 2 + H 2 O  H 2 CO 3  HCO 3 - + H +  CO 3 2- + 2 H + CO 2 + H 2 O  H 2 CO 3  HCO 3 - + H +  2H + + CO 3 2- Acid will react with carbonate May get dissolution of carbonate skeletons and production of CO 2 gas (and evasion to the atmosphere)

8 When you add base to seawater CO 2 + H 2 O  H 2 CO 3  HCO 3 - + H +  CO 3 2- + 2 H + CO 2 + H 2 O  H 2 CO 3  HCO 3 - + H +  2H + + CO 3 2- More CO 2 will dissolve in seawater and you may get carbonate production These equilibrium reactions help the ocean buffer itself from changes in pH

9 Why is the ocean slightly basic or alkaline? Alkalinity = the amount of acid needed to neutralize a base (in this case, the amount of H + needed to neutralize bicarbonate [HCO 3 - ] and carbonate [CO 3 2- ]) Equilibrium reactions and the oceans carbonate system Two major processes affect CO 2 in the ocean –Addition and removal by organisms –Carbonate mineral precipitation & dissolution

10 Ocean buffering At the normal pH of seawater, about 80% of the carbon compounds are in the form of HCO 3 - (bicarbonate) A decrease in dissolved CO 2 (e.g., from photosynthesis) will cause more bicarbonate to change to CO 2 (or more to go in from atmosphere) HCO 3 - + H +  H 2 CO 3  CO 2 + H 2 O A decrease in bicarbonate will cause more carbonate to go to bicarbonate (may get carbonate dissolution 2H + + CO 3 2- (carbonate)  HCO 3 - + H +

11 Dominant at high pH Dominant at low pH

12 Ocean buffering Use of carbonate (CO 3 2- ) from seawater will cause more bicarbonate disassociate to replace it: HCO 3 - + H +  2H + + CO 3 2- (carbonate) A decrease in bicarbonate will cause more carbon dioxide to go to bicarbonate CO 2 + H 2 O  H 2 CO 3  HCO 3 - + H + Carbonate precipitation causes a net loss of carbon dioxide from the ocean

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14 Variability in ocean pH Really, there’s not too much (~7.5 – 8.5) Surface pH in warm productive waters is ~ 8.5 In warmer surface waters less CO 2 can dissolve Where there’s high rates of photosynthesis CO 2 + H 2 O  CH 2 O + O 2 removing CO 2 from the water column, pH can increase slightly (more basic or alkaline) and reactions move to the left to try and free H + will be try to replace CO 2 CO 2 + H 2 O  H 2 CO 3  HCO 3 - + H +  CO 3 2- + 2 H +

15 Higher surface pH due to photosynthetic CO 2 removal and warmer temperatures Lower subsurface pH due to respiratory CO 2 inputs and decay or organisms Lower deep pH due to cold temperatures, high pressure and no CO 2 removal by plants also have CaCO 3 dissolution. Deep, cold water (e.g., 4500 m) has a pH of about 7.5 Can go as low as 7.0 at the bottom (remember the CCD?) CO 2 removal from CaCO 3 formation

16 Revisit the CCD Calcium (Ca 2+ ) is much more abundant than CO 3 2- in seawater so CaCO 3 saturation is described by CO 3 2-. Increase pressure, increase CaCO 3 solubility Decrease temperature, increase CaCO 3 solubility So, with depth, CaCO 3 becomes undersaturated pH is lower with depth Depth of CCD is controlled by carbonate equilibrium and pH –High productivity areas have deeper CCD (because more carbon)

17 Distribution with depth O 2 and CO 2 are about opposite due to complimentary sources and sinks O 2 produced at surface by ps & CO 2 is consumed At depth (no light), respiration exceeds ps Compensation depth Where O 2 production = CO 2 production

18 Now what about carbon This buffering assumes that there is a relatively constant carbon concentration in seawater (steady state) We’re increasing the CO 2 content of the atmosphere by about 0.2% per year. About half that has gone into the ocean This has caused ocean pH to decrease and more is projected

19 Simplified C cycle We’re changing fluxes and sizes of major reservoirs

20 Take home points What is the average pH of the oceans The 4 major forms of CO 2 in the ocean Most dissolved inorganic C is present in the ocean as bicarbonate ion (very little as CO 2 or carbonic acid (e.g., 1%) CO 2 and O 2 concentrations tied with biological processes Carbonate system is important for buffering ocean’s pH Reactions between the different forms of C in water produce or consume H + Also reactions of different forms of C in water buffer big changes in atmospheric CO 2


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