OBJECTIVES:

• Introduce principles of coagulation for water treatment
• Construct and interpret logC-pH diagrams for solid-aqueous phase equilibrium
• Solve example problems with precipitation of solid phase

 
    Many water-borne pollutants of concern to human health are either particles (pathogenic organisms) or tend to associate with  particulate materials due to low solubility (many organic compounds, toxic metals). A common method of removing these particles is to allow them to settle out, after "scavenging" them on active surfaces (see Chapter # 8). However, very small particles have settling out times on the order of days or even weeks. Coagulation can reduce the time required to separate these particles by increasing the size and mass of the average particle, reducing their settling time. For this reason, coagulation is an essential feature of most water treatment systems.

     In order to overcome the stability of particles in treated water, a coagulant is added. This coagulant is a material that causes particles to aggregate into larger heavier masses (known as flocs) which can be more easily removed through settling or filtration. A common treatment process is to add metal salts in concentrations sufficient to cause precipitation of the metal hydroxide. These metal hydroxides are known as "sweep flocs". They collect colloidal particles both during formation of the precipitate and through subsequent collisions. In addition, dissolved organic material will adsorb to these precipitates, helping to remove color and odor from the water.

     The most common metals used in coagulation are aluminium and iron(III), as they both form insoluble metal hydroxides at low concentrations. These metals are usually added as sulfates or chloride species: for example, FeCl₃, Fe₂(SO₄)₃, AlCl₃, and alum (Al₂(SO₄)₃*14 H₂O). The exact dose required is a function of the chemistry of the treatment water, particularly the pH, alkalinity, hardness, ionic strength, and temperature.

• When iron(III) is added to water, the Fe⁺³ ion reacts with OH^- to form hydrolysis species and Fe(OH)₃ (s).
• Species: (H₂O), H⁺, OH^-, Fe⁺³, FeOH²⁺, Fe(OH)₂⁺ , Fe(OH)₄^-, Fe₂(OH)₂⁴⁺, Fe₃(OH)₄⁵⁺, Fe(OH)_3(s).
• Reactions:

            H₂O <=> H⁺ + OH^-                                Log K= -14.0
            FeOH²⁺<==>  Fe⁺³ + OH^-                    Log K= -11.8
            Fe(OH)₂⁺ <==>  Fe⁺³ + 2 OH^-              Log K= -22.3
            Fe(OH)₄^- <==> Fe⁺³ + 4 OH^-                Log K= -34.4
            Fe₂(OH)₂⁴⁺<==> 2 Fe⁺³ + 2 OH^-          Log K= -25.1
            Fe₃(OH)₄⁵⁺<==>3 Fe⁺³ + 4 OH^-           Log K= -49.7
             Fe(OH)₃ (s) <==> Fe⁺³ + 3 OH^-           Log K= -38.8
 
• Tableau and logC-pH diagram (see your book)
• For the logC-pH diagram of this system click here
• For the MATLAB code used to create this graph, click here.

Add 10^-4 M FeCl₃ to water at pH 7. Will a precipitate form? What will be the dominant species in the aqueous phase? What is the pH of the solution?

 

Other Considerations in Real Systems
 

• Alkalinity- the presence of high alkalinity can buffer the water, keeping pH closer to 7-8.
• Other ligands- Ions that form stable complexes with Fe⁺³ can increase the total iron solubility
• Other precipitates- If ions (such as PO₄^3-) that form precipitates with iron are present, these phases may form before the hydroxide phase (depending on the K_sp values).

        Below, you will find a few solubility diagrams. Test yourself to find out if you can interpret them!

    The first two are solubility diagrams for iron-oxides. What has changed? Why?