Inorganic Carbon

pH

• When water ionizes, a proton (hydrogen ion) is transferred from one water molecule to another, resulting in a hydrated hydrogen ion and a hydroxyl ion

• H2O + H2O ↔ H3O+ + OH-

Lowry-Bronsted concept of acids and bases

• Acid - a chemical species that has the capacity to donate or transfer a proton to an acceptor

• Base - a chemical species that has the capacity to accept a proton from a donor

Equilibrium constants

• For the hydration of water, the equilibrium reaction is represented by:

(H3O+)(OH-)• Keq =

(H2O)2

Equilibrium constants

• Kw = (H3O+) (OH-)

• At 25oC, Kw = 1 x 10-14 M2

pH

• Hydrogen ion activity is expressed as pH

• pH = - log (H+)

pH

• In pure distilled water:

• pH = - 1/2 log Kw

• pH = - 1/2 log (1 x 10-14)

• pH = - log (1 x 10-7)

• pH = 7

pH

• At neutrality, the pH of a solution is 7

• pH of an acidic solution is less than 7

• pH of a basic solution is more than 7

Inorganic Carbon

• Most of the carbon in a lake is in the form of dissolved inorganic carbon

• Less is in the form of organic detritus

• A small amount is in the form of living organic matter (organisms)

Sources of Inorganic Carbon

• Carbon dioxide in the atmosphere (0.027% - 0.044%; average concentration in the air is 0.033%)

• Rainwater contains roughly 0.6 mg/l• Carbonate rock• Respiration by living organisms

– Microbes in the soil– Organisms in the lake

• Burning of fossil fuels

Hydration of carbon dioxide

• Carbon dioxide in the atmosphere dissolves in water

• CO2 (air) ↔ CO2 (dissolved) + H2O

• CO2 + H2O ↔ H2CO3

Hydration of carbon dioxide

• This reaction is slow (t1/2 = 15 seconds)

• The equilibrium concentration of H2CO3is 1/400th of unhydrated carbon dioxide

• Formation of carbonic acid by hydration of carbon dioxide predominates at a pH of less than 8

Forms of Inorganic Carbon

• Dissolved carbon dioxide or free carbon dioxide - CO2

• Carbonic acid - H2CO3• Bicarbonate - HCO3

-

• Carbonate - CO32-

Ionization of Inorganic Carbon

• Carbonic acid rapidly dissociates to yield a proton and bicarbonate

• H2CO3 ↔ H+ + HCO3-

Ionization of Inorganic Carbon

(HCO3-)(H+)

• K1 = (H2CO3)

• pK1 = 6.43 (at 15oC)

Ionization of Inorganic Carbon

• Bicarbonate dissociates to yield a proton and carbonate

• HCO3- ↔ H+ + CO3

2-

Ionization of Inorganic Carbon

(CO32-)(H+)

• K2 = _______________(HCO3

-)

• pK2 = 10.43 (at 15oC)

Dissociation of Bicarbonate and Carbonate

• CO32- + H2O ↔ HCO3

- + OH-

• HCO3- + H2O ↔ H2CO3 + OH-

• H2CO3 ↔ H2O + CO2

Proportions of Inorganic Carbon

• Free carbon dioxide and carbonic acid dominates at pH less than 4.3

• Bicarbonate dominates at a pH of 8.3

• Carbonate dominates at a pH of greater than 12.6

Proportions of Inorganic Carbon

Buffering of pH

• Addition of hydrogen ions (acid solutions) is neutralized by reaction of hydrogen ions directly with bicarbonate and carbonate

Buffering of pH

• H2CO3 ↔ H+ + HCO3-

• HCO3- ↔ H+ + CO3

2-

Buffering of pH

• Addition of hydrogen ions (acid solutions) is neutralized by reaction with the hydroxyl ions formed by the hydration of bicarbonate and carbonate

Buffering of pH

• HCO3- + H2O ↔ H2CO3 + OH-

• CO32- + H2O ↔ HCO3

- + OH-

• H+ + OH- ↔ H2O

Supply of Inorganic Carbon

• As water percolates through the soil, microbial respiration increases the concentrations of carbon dioxide dissolved in water

• Dissolved carbon dioxide dissociates to form carbonic acid

Supply of Inorganic Carbon

• Carbonic acid solubilizes limestone of calcium-rich rock and produces calcium bicarbonate (Ca(HCO3)2), which is relatively soluble

• H2CO3 + CaCO3 ↔ Ca(HCO3)2

Supply of Inorganic Carbon

• The calcium bicarbonate increases the amount of ionized Ca2+ and HCO3

- in the lake water

• pH increases from the hydroxyl ions released by the dissociation of bicarbonate

Solubility of Inorganic Carbon

• Solubility of carbon dioxide increases markedly in water that contains carbonate

• A definite amount of free carbon dioxide will remain in solution after equilibrium is reached between calcium, bicarbonate, carbonate, andundissociated carbonate

Solubility of Inorganic Carbon

• A certain amount of free carbon dioxide is required to keep Ca(HCO3)2 in solution and this free carbon dioxide will not dissolve any more CaCO3

• Ca (HCO3)2 ↔ CaCO3 + H2O + CO2

Solubility of Inorganic Carbon

• The amount of free carbon dioxide that is required to keep the calcium bicarbonate in solution is called the carbon dioxide of equilibrium

Solubility of Inorganic Carbon

• If more free carbon dioxide is added in excess of the carbon dioxide of equilibrium, it will dissolve more CaCO3

• The amount of free carbon dioxide in excess of the carbon dioxide of equilibrium is termed aggressive carbon dioxide

Solubility of Inorganic Carbon

• If the level of free carbon dioxide is reduced below the carbon dioxide of equilibrium (such as by removal of CO2 through photosynthesis of algae andmacrophytes), CaCO3 will precipitate until equilibrium is reestablished by the formation of CO2

Influence of Metabolism on pH

• 6 CO2 + 6 H2O ↔ C6H12O6 + 6 O2

• Primary production lowers the concentration of CO2 , causing an increase in pH

• Respiration increases the concentration of CO2 , causing a decrease in pH

Acid Neutralizing Capacity

• The inorganic carbon system provides the major buffering of natural water against changes in pH

• Alkalinity is the milliequivalents of acid necessary to neutralize the hydroxyl, bicarbonate, and carbonate ions in a liter of water

Acid Neutralizing Capacity

• Alkalinity is most correctly expressed in terms ofmiliequivalents per liter

• Alkalinity is often expressed less accurately as concentration of all species as CaCO3 in mg/l

Hardness

• Hardness is a measure of the calcium and magnesium salts in water

• Carbonate hardness is the amount of calcium and magnesium associated with bicarbonates and carbonates

• Hardness is related to the capacity to precipitate soap

Additional Sources of Acid Neutralizing Capacity

• Inorganic carbon accounts for most of the acid neutralizing capacity in lake water

• Other compounds contribute minor portions of the acid neutralizing capacity– Borate– Silicate– Phosphate– Organic matter

pH in Natural Waters

• Photosynthetic activity increases the pH

• Respiration by animals, plants, and microbes releases CO2 and decreases the pH

• Water draining volcanic and pyritic rock formations is often low in pH

pH in Natural Waters

• Release of organic acids by plants can slightly reduce the pH

•• Cation exchange by Sphagnum, a

moss, can greatly increase the hydrogen ion activity and lower the pH– Sphagnum bogs are frequently acidic