Chemical Reactions: Acids and Bases

1
Chemical Reactions Acids and Bases

• Significance and measurement of pH (5.1)
• Acids and bases overview (5.2)
• Carbon dioxide and carbonic acid species in
  natural waters (5.3)
• Influence of CO2 dissolution on pH (notes)
• Distribution of carbonate species as function of
  pH (notes Bjerrum plot transparency)
• CO2-pH relations in natural waters (Table 5.3,
  Fig. 5.3)

2
Humic and fulvic acids

• Humic compounds are refractory, high molecular
  weight organic compounds - mixtures of long
  polymer chains of variable lengths gfw. They
  act as acids, giving up H from carboxylic and
  phenolic groups, which gives them high surface
  charge and CEC (i.e., they are effective
  chelates).
• Fulvic acids low molecular weight (500-1500
  g/mol), soluble in acid and alakaline solutions
• Humic acid Intermediate molecular weight,
  soluble in alkaline solutions, insoluble below pH
  2
• Humin highest molecular weights, insoluble.
• pKa of humic and fulvic acids gt 3.6 (weak)
• Fulvic acid produced by organic decay in soil A
  horizon. Forms soluble complexes with Fe3, Al3
  and other metals, which are transported during
  downward percolation. Aerobic decay of fulvic
  acid chelates causes metals to precipitate in B
  horizon.

3
Metal hydroxides
Acid-base properties of minerals and rocks
Most minerals are salts of weak acids (silicic,
carbonic, phosphoric) and strong bases (e.g.,
NaOH, HOH, Ca(OH)2). During dissolution, anion
hydrolyzes to form weak acid OH-, increasing pH
of water. Thus, minerals help to neutralize
acidity, and ? contribute to alkalinity. Solubilit
ies increase with decreasing pH. Examples
calcite, apatite

• Metal hydroxides are amphoteric, meaning they can
  either accept or donate protons, depending on pH.
  Ex
• Al(OH)2 H Al(OH)2 H2O
• Al(OH)2 H2O Al(OH)3? H
• Hydrolysis of metal cations (Fe3, Al3, Mn4,
  etc.) contributes to acidity. Intermediate IP
  causes cation to repel H from H2O molecules in
  inner sphere, forming aqueous hydroxycation. The
  concentrations of these species affects pH.

4
Acidity Alkalinity

• Acidity capacity of water to give or donate
  protons
• Immediate acidity (and alkalinity) considers
  species present only in solution (not solids)
• Acidity undesirable because it increases
  solubilities of many minerals ? TDS, ? hardness,
  and ? concentrations of toxic heavy metals.
  Corrosive to fish other aquatic lifeforms Acid
  waters must be treated by dilution and
  neutralization, at great expense.
• In dilute, potable waters, most of acidity is
  cause by carbonate species, and CA 2H2CO3
  HCO3- H - OH-
• Alkalinity capacity of a water to accept protons
• Carbonate alkalinity HCO3- 2CO32-
• Total alkalinity CB HCO3- 2CO32- OH- - H
• Bicarbonate is the only significant base in
  natural water with pH lt 8.3, and derives from
  weathering of silicates and dissolution of
  carbonates.

5
Acidity Alkalinity Titrations

• Total equivalents of acidity CA or alkalinity CB
  measured by titrating with standardized solution
  of acid or base Cstd (eq/l) Vstd(l)
  Csample(eq/l) Vsample(l)
• - for freshwaters, may also be computed from a
  total analysis
• Inflections in titration curve are equivalence
  points

6
Summary of controls on the pH of natural waters

• pH of most natural waters range between 4 and 9
  (transparency, Fig. 5.4).
• Water-dominated systems humid climate soils
  stream sediments. Carbonic and organic acids
  frequently replenished by rainfall fresh
  recharge, extensive reaction with minerals during
  chemical weathering. Intermediate pH results from
  balance between dissociation of weak acids and
  consumption of H (and CO2) by weathering
  reactions.
• Rock-dominated systems (arid soils, deep
  groundwaters) - low fluid/rock ratio. Minerals
  (salts of weak acids strong bases) dissolve
  increase pH.