Determination of water hardness

The hardness of a water sample can be determined by direct titration of a known volume with a standardised solution of a soap such as potassium oleate. Initially, this soap will precipitate, and the calcium and magnesium salts and the water will not produce any foam on shaking. Gradual addition of the potassium oleate solution is continued until the agitation of the solution produces a reasonably permanent foam. F rom Scheme 8.4, 1.00 ml of 0.02 M potassium oleate solution 2R– K+ (aq) + Ca2+ (aq)= R2Ca(s) + 2K+ (aq) R= Oleate radical is equivalent to 0.01 mmol Ca2+. Thus, if a 100.0 ml sample of water is titrated with 0.02 M potassium oleate solution to just give a permanent foam after addition of V ml of soap, equivalent to V 0.01 mmol Ca2+, the hardness willbe V 10 ppm CaCO3.Despite the simplicity of this method, it is not always possible to detect precisely the point at which a layer of foam persists. Hardness is therefore more usually determined by titration with a standardised solution of the disodium salt of ethylenediamine tetra-acetic acid (EDTA). EDTA forms very stable, soluble complexes with calcium and magnesium ions (Figure 8.1). EDTA is a tetraprotic acid and the various acid–base forms are usually abbreviated by H4Y, H3Y–, H2Y2–,HY3–,and Y4–, as the degree of dissociation progresses. Scheme 8.5 illustrates complex formation with calcium and magnesium ions. The equilibrium constants are large and ensure quantitative formation of the colourless complexes on addition of EDTA to a solution containing calcium and magnesium ions. The titration involves gradual addition of standardised EDTA solution(Na2H2Y) to the water sample, buffered at pH 10 with an ammonia–ammonium

chloride solution. The indicator is a dye that changes colour on complexing with calcium or magnesium, such as Eriochrome Black T (CI Mordant Black 11). At the beginning of the titration, the large concentrations of calcium and magnesium ions ensure that all the dye indicator is in the form of the red complexes with these metals (Ca–Dye–, Scheme 8.6). Very close to the equivalence point, the free Ca2+ and Mg2+ ion concentrations become so low that the added EDTA begins to remove the metals from their red complexes with the dye. This liberates the blue dye ion. The colour of the solution therefore changes from red (dye–metal complex) to blue (free dye anion) at the end-point.

In the titration of calcium ions alone at pH 10, the red calcium–dye complex(Ca–Dye–) is not sufficiently stable and the calcium transfers to EDTA liberating the blue dye anion (H–Dye2–) slightly before the true equivalence point. For titrations of magnesium ion, the magnesium–dye complex is more stable than that of calcium and the blue colour does not occur until all the Mg2+ ions in solution have reacted with the EDTA. Thus, the red-to-blue colour change occurs at the true equivalence point. Towards the end of the titration of hard water, as long as calcium ion is present, it will liberate magnesium ion from the magnesium–EDTA complex because the calcium–EDTA complex is more stable. The liberated magnesium thus ensures that the more stable magnesium–dye complex is always present, until all the calcium has been complexed with EDTA. The last addition of this reagent then removes the magnesium from the red magnesium–dye complex and the solution becomes blue. To ensure that a water sample contains a sufficient quantity of magnesium, it is common to add a little to the EDTA solution before standardisation with a known calcium ion solution.