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 moreacidic the solution is. The scale is logarithmic so a change of 1 pH unit is a 10-fold changein H+ concentration (100 = 1; 10-1 = 0.1; 10-2 = 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 CO2 in the ocean so shouldn’t it be acidic if the CO2 combines with H2O to form carbonic acid?

• CO2 is actually present in several different forms in water.

CO2

• Soluble in water• Ocean is a big reservoir for CO2 (only bigger reservoir is

sediments and sed. rocks)• The total amount of CO2 in seawater is about 60 times

the amount of CO2 in the atmosphere• Because it forms a variety of chemical species in water,

more CO2 can dissolve than is predicted by gas solubility alone!

• Two major processes affect CO2 in the ocean– Addition and removal by organisms– Carbonate mineral precipitation & dissolution

CO2 in seawaterCO2 + H2O H2CO3 HCO3

- + H+ CO32- + 2 H+

1: CO2 + H2O H2CO3 (carbonic acid) Carbonic acid rapidly dissociates to form ions

2: H2CO3 HCO3- (bicarbonate) + H+

3: HCO3- + H+ CO3

2- (carbonate) + 2 H+

Some of the bicarbonate combine with H+ ions to form carbonate

**At a given pH, CO2, H2CO3, HCO3-, CO3

2- and H+ are in equilibrium

When you add acid to seawater

CO2 + H2O H2CO3 HCO3- + H+ CO3

2- + 2 H+

CO2 + H2O H2CO3 HCO3- + H+

2H+ + CO32-

Acid will react with carbonateMay get dissolution of carbonate skeletons

and production of CO2 gas (and evasion to the atmosphere)

When you add base to seawater

CO2 + H2O H2CO3 HCO3- + H+ CO3

2- + 2 H+

CO2 + H2O H2CO3 HCO3- + H+

2H+ + CO32-

More CO2 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 [HCO3

-] and carbonate [CO3

2-])

• Equilibrium reactions and the oceans carbonate system

• Two major processes affect CO2 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 HCO3-

(bicarbonate)• A decrease in dissolved CO2 (e.g., from

photosynthesis) will cause more bicarbonate to change to CO2 (or more to go in from atmosphere)

HCO3- + H+ H2CO3 CO2 + H2O

• A decrease in bicarbonate will cause more carbonate to go to bicarbonate (may get carbonate dissolution

2H+ + CO32- (carbonate) HCO3

- + H+

Dominant at high pH

Dominant at low pH

Ocean buffering

• Use of carbonate (CO32-) from seawater will cause more

bicarbonate disassociate to replace it:

HCO3- + H+ 2H+ + CO3

2- (carbonate)

• A decrease in bicarbonate will cause more carbon dioxide to go to bicarbonate

CO2 + H2O H2CO3 HCO3- + 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 CO2 can dissolve• Where there’s high rates of photosynthesis

CO2 + H2O CH2O + O2

removing CO2 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 CO2

CO2 + H2O H2CO3 HCO3- + H+ CO3

2- + 2 H+

Higher surface pH due to photosynthetic CO2 removal and warmer temperatures

Lower subsurface pH due to respiratoryCO2 inputs and decay or organisms

Lower deep pH due to cold temperatures,high pressure and no CO2 removal byplants also have CaCO3 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?)

CO2 removal from CaCO3 formation

Revisit the CCD

• Calcium (Ca2+) is much more abundant than CO32- in

seawater so CaCO3 saturation is described by CO32-.

• Increase pressure, increase CaCO3 solubility• Decrease temperature, increase CaCO3 solubility• So, with depth, CaCO3 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

O2 and CO2 are about opposite due to

complimentary sources and sinks

O2 produced at surface by ps & CO2 is

consumedAt depth (no light),

respiration exceeds ps

Compensation depthWhere O2 production =

CO2 production

Now what about carbon

• This buffering assumes that there is a relatively constant carbon concentration in seawater (steady state)

• We’re increasing the CO2 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 CO2 in the ocean• Most dissolved inorganic C is present in the ocean as

bicarbonate ion (very little as CO2 or carbonic acid (e.g., 1%)

• CO2 and O2 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 CO2