Environmental Applications. Part I.
Common Forms of the Elements in Water
You will study how the variation of pH affects the availability of different elements
as nutrients or pollutants in a natural body of water, Purity Springs Lake. Although the chemistry of
a natural body of water such as a lake is complex, let's assume that Purity
Springs Lake contains no oxidizing, reducing or complexing agents. The only reactions
occurring in the lake come from the hydrolysis of hydrated cations.
|All hydrolysis processes are equilibria.
will exist in more than one form at any given pH.
A distribution diagram shows the relative concentrations of the different forms
of an element in solution as a function of pH. This type of diagram is generated from
experimentally measured concentrations of the species present plotted
A predominance diagram shows the predominant species that is present in solution
at a given pH. This type of diagram can be generated with a minimum amount of calculation
using basic concepts about the hydrolysis characteristics of metal cations and oxoanions.
You will use predominance diagrams in analyzing Purity Springs Lake.
Most natural waters have pH ranges between 6 and 9.
Predominance Diagrams For Some Selected Elements
A number of generalizations come from studying predominance diagrams for different elements.
||Type of Species
- nonacidic cation
- feebly acidic cation
- weakly acidic cation
in aqueous solution
|insoluble hydroxide or oxide
- moderately acidic cation
- strongly acidic cation
|in bottom sediments
|The pH of hydroxide precipitation is roughly equal to its pKa
- element in high oxidation state
may be partially hydrolyzed
The radioisotopes 99Tc, 90Sr and244Pu (fallout from atomic weapon testing) are deposited
in Purity Springs Lake. The lake has a pH in the range 5.5 to 7. In what predominant form does each
- Predict the oxidation state of each element
- Evaluate the acidity of each cation to determine if it can remain as a hydrated
ion in neutral water
- Evaluate the possibility of forming an oxo anion (for strongly acidic cations)
Click here to see a tabulation of the forms in which the elements predominate in moderately aerated water of
pH 5.5 to 7.
Consequences of Hydrolysis Reactions of the Elements
- Elements having moderately and strongly acidic cations are rarely available
for biological activity
- unavailable elements are usually not essential for life or are toxic to life
- exceedingly toxic species such as Pu4+ may be harmful even though
most of it is in the sludge at the bottom of the body of water and only a minute
amont is present in solution due to equilibrium
- Unavailable elements may become soluble when a body of water becomes highly
polluted with acidic or basic pollutants
- acid rain can cause the pH of a poorly buffered lake to drop sufficiently to
convert insoluble Al(OH)3 into Al3+ (aq) which is toxic to fish
- raising the pH of an acidic body of water by neutralization may precipitate
weakly and moderately acidic ions as hydroxides
- Al3+ (aq)in acidic lake water is converted into gelatinous
Al(OH)3 in the more basic environment of fish gills, coating the gills
- fish "sneeze" themselves to death trying to rid themselves of the precipitate
- In acid mine drainage, Fe2(SO4)3 which results
from air oxidation of pyrite,dissolves in the water draining from the mine forming
- acidic mine drainage becomes diluted as it enters uncontaminated streams
- at the higher pH, the hydroxy cation of iron is converted into insoluble Fe(OH)3,
an unsightly yellow precipitate called yellow boy
- Elements having weakly or moderately acidic cations which serve as nutrients for plant and animal life
may become unavailable in neutral or slightly basic solutions
- Zn2+, Cu2+, Co2+, Fe2+, Mn2+ are weakly acidic essential micronutrients
- these ions precipitate from dilute solutions at pH's in the range 5.3-8.5
- liming acidic soil or lakes with CaO may remove these essential micronutrients
- the moderately acidic Fe3+cation precipitates from dilute solution
at pH's above 2.0
- iron-deficiency anemia cannot be treated by ingesting Fe3+ salts
because the pH of the intestine where iron absorption occurs is much higher than 2.0