Chemical Precipitation
Definition:
Formation of a separable solid substance from a solution, either by converting the
substance into a solid form or by changing the composition of the solvent to diminish
the solubility of the unwanted substance and later on filter it out. [1]
Figure 1 Precipitation
chemical
solution with reaction insoluble
soluble ions complexes + “clean water”
Process Description:
The fundamental idea of precipitation is to precipitate the chemical from dissolved
substance in the wastewater by adding a reagent, to form an insoluble substance with
the to-be-separated matter. Generally, precipitation occurs in a 1 to 1 mole ratio;
similarly to one molecule of dissolved matter with one molecule of reagent will form
an insoluble precipitate. However, a certain amount of over-dosage is needed for a
complete removal. [2] For instance, to achieve a complete derivation of certain amount
of barium from soluble barium chloride, a higher amount of caustic soda (NaOH) as
reagent is needed, compare to amount of barium occur in the particular barium
chloride solution. [3]
2 NaOH +BaC l 2 →2 NaCl+Ba ¿
(insoluble salt)
Besides, heavy metals can be also precipitated as hydroxide by increasing the pH.
Once the heavy metal has been precipitated, it can be separated from the main stream
using filtration, flotation or sedimentation. Normally, a polymer will be added to
improve the silt separation. [2]
Advantages and Disadvantages:
[2]
Advantage Disadvantage
Simple and effective Large amount of reagent is needed
Able to remove substance that other High cost
techniques can hardly remove High amount of silt is produced
Removal of a very specific
component, without removing
desired components
High degree of selectivity
Industrial Application:
Precipitation can be applied to almost any liquid waste stream containing a
precipitable hazardous constituent. [4] In this assignment will be discussed about the
precipitation technique used in the Public Water Supply Treatment. [5]
A typical precipitation reaction is used to remove 2 major classes of ions:
1. Calcium(Ca2+) ions and magnesium(Mg2+) ions
2. Iron(Fe2+) ions and manganese(Mn2+) ions
The process of removing Ca2+ and Mg2+ from the water is known as water softening.
Lime (Ca(OH)2) and soda ash (Na2CO3), are typically used to soften public water
supplies.
The reaction is shown in the chemical equations below:
Mg2+ (aq) + Ca2+ (aq) + 2 OH- (aq) Mg(OH)2 (s) + Ca2+
from Lime precipitate
Ca2+ (aq) + Ca2+ (aq) + 2 CO32- (aq) 2CaCO3 (s)
from water from Line from soda ash precipitate
The reaction of removing Fe2+ and Mn2+ from water is oxidation. By using molecular
oxygen (O2) or another oxidant such as potassium permanganate (KMnO 4), Fe(II) is
readily oxidized to Fe(III) in solution. If the solution is alkaline (high pH, basic), the
Fe(III) forms Fe(OH)3. As the concentration of Fe(OH)3 increases, the oxygens start
to coordinate between multiple iron ions, and a lattice begins to form. At some point
in this lattice formation, the Fe(OH)3 starts to look like Fe2O3 (rust) and precipitates.
Hence, by adding an oxidant to the water and raising the water's pH at the water-
treatment plant, an insoluble precipitate is formed.
3 Fe2+ (aq) + MnO4- (aq) + 2 H20 (l) 3 Fe3+ (aq) + MnO2 (s) + 4OH-
The insoluble rust can then be removed by sedimentation or filtration.
References:
1. The Editors of Encyclopaedia Britannica, Chemical Precipitation.
Available at https://www.britannica.com/science/chemical-precipitation
2. VITO, Boeretang 200, B-2400 Mol, Belgium.
Available at https://emis.vito.be/en/techniekfiche/chemical-precipitation
3. Prof. M.S.Subramanian, Indian Institute of Technology Madras, Environment
Chemistry and Analysis.
Retrieved at:
http://ebooks.bharathuniv.ac.in/gdlc1/gdlc4/FirstYear/Environmental
%20Chemistry/Notes/Chemical%20Methods%20of%20Treatment.pdf
4. Rachel Casiday, Greg Noelken and Regina Frey, Dept. of Chemistry, Washington
University, “Treating the Public Water Supply: What Is In Your Water, and How Is It
Made Safe To Drink?”
Retrieved at:
http://www.chemistry.wustl.edu/~edudev/LabTutorials/Water/PublicWaterSup
ply/PublicWaterSupply.html
