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 produces an acidic or basic (alkaline) solution –Acid releases H + –Base combines with H +

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; = 0.1; = 0.001, etc) pH = -log 10 [H + ] Average pH of seawater is ~ 8

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.

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

CO 2 in seawater CO 2 + H 2 O  H 2 CO 3  HCO H +  CO 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 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

When you add acid to seawater CO 2 + H 2 O  H 2 CO 3  HCO H +  CO H + CO 2 + H 2 O  H 2 CO 3  HCO 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)

When you add base to seawater CO 2 + H 2 O  H 2 CO 3  HCO H +  CO H + CO 2 + H 2 O  H 2 CO 3  HCO 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

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

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 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 H +

Dominant at high pH Dominant at low pH

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

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 H +  CO H +

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

Revisit the CCD Calcium (Ca 2+ ) is much more abundant than CO 3 2- in seawater so CaCO 3 saturation is described by CO 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)

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

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

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

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