EDTA Hardness Titration of a Borewell Water Sample
A 50.0 mL sample of borewell water is titrated with 0.0100 M EDTA at pH 10, using Eriochrome Black T as indicator. The mean titre is 12.4 mL. A second 50.0 mL portion is boiled for 10 minutes, cooled, filtered, and titrated in the same way; it requires 9.1 mL. The water lathers only poorly with soap. The student reports total hardness as 124 ppm CaCO3. The teacher notes that the water contains about 0.5 mg L−1 Fe3+ and that the endpoint in the first titration was slow and somewhat fading. The laboratory has NH3/NH4Cl buffer, triethanolamine, Na2S, and KCN. The teacher asks the student to recalculate the hardness, explain the boiled-water result, and justify the choice of buffer and masking agent.
1. Recalculate total hardness as ppm CaCO3 and identify the student’s arithmetic error. Is the water soft, moderately hard, or very hard by the usual classification?
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EDTA reacts with Ca2+ and Mg2+ in a 1:1 mole ratio, and hardness is conventionally expressed as the equivalent mass of CaCO3 (molar mass 100 g mol−1).
n(EDTA) = 0.0100 mol L−1 × 0.0124 L = 1.24 × 10−4 mol. This equals the total moles of Ca2+ + Mg2+ in 50.0 mL. Mass of CaCO3 equivalent = 1.24 × 10−4 mol × 100 g mol−1 = 0.0124 g = 12.4 mg.
Concentration = 12.4 mg / 0.0500 L = 248 mg L−1, i.e. 248 ppm CaCO3.
The student’s value of 124 ppm is exactly half the correct value. The likely error is using 100 mL instead of 50.0 mL, or dividing by 0.100 L rather than 0.0500 L. The calculation is otherwise correct.
By the usual classification (soft <60 ppm, moderately hard 60–120 ppm, hard 120–180 ppm, very hard >180 ppm), 248 ppm is very hard. This is consistent with the poor lathering with soap, because Ca2+ and Mg2+ precipitate the soap as insoluble calcium and magnesium salts.
2. Interpret the boiled sample: calculate the permanent and temporary hardness. State the assumption behind attributing the difference to bicarbonates, and what would happen if the water contained MgSO4.
Read the worked answer
For the boiled sample: n(EDTA) = 0.0100 × 0.0091 = 9.10 × 10−5 mol. Mass CaCO3 equivalent = 9.10 × 10−5 × 100 = 9.10 × 10−3 g = 9.10 mg in 50.0 mL. Permanent hardness = 9.10 mg / 0.0500 L = 182 ppm.
Temporary hardness = total − permanent = 248 − 182 = 66 ppm.
The assumption is that boiling decomposes only the hydrogen carbonates: Ca(HCO3)2 → CaCO3(s) + CO2 + H2O, Mg(HCO3)2 → MgCO3(s) + CO2 + H2O, with some MgCO3 hydrolysing to Mg(OH)2. The precipitated solids are filtered off, so the titre after boiling measures the ions that remain in solution, i.e. the permanent hardness due to sulfates and chlorides.
If the water contained MgSO4, it would not be removed by boiling. Magnesium sulfate is soluble and stable under boiling, so it would remain in the filtrate and contribute to the 182 ppm. The difference of 66 ppm is therefore not necessarily all bicarbonate hardness; it is the fraction removed by boiling, which is usually bicarbonates but can also include some carbonate that precipitates on heating. A separate bicarbonate titration would be needed to confirm the assignment.
3. Explain the role of the NH3/NH4Cl buffer and Eriochrome Black T. Why does Fe3+ interfere, and which masking agent would you choose? Comment on KCN.
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The NH3/NH4Cl buffer holds the titration near pH 10. At this pH, EDTA is mainly in its deprotonated forms (HY3− and Y4−), so it binds Ca2+ and Mg2+ strongly. The indicator Eriochrome Black T also works in this range: it forms a wine-red complex with Ca2+/Mg2+, and at the endpoint EDTA displaces the indicator, leaving the free blue form. If the pH is too low, the metal–indicator complex is too stable or the endpoint is indistinct; if too high, Mg(OH)2 may precipitate and the result becomes unreliable.
Fe3+ interferes because it forms a very stable EDTA complex and can also bind strongly to the indicator, blocking it. This causes a slow or fading endpoint and can make the titre too high. The water contains only about 0.5 mg L−1 Fe3+, which is roughly 0.9 ppm CaCO3 equivalent-far too small to explain the difference between 124 and 248 ppm-but it can still spoil the endpoint. At pH 10 some Fe3+ may precipitate as Fe(OH)3, which can adsorb indicator or cause turbidity.
Triethanolamine is the preferred masking agent: it complexes Fe3+ (and Al3+) strongly in alkaline solution without masking Ca2+ and Mg2+. Na2S can precipitate heavy metals as sulfides but may also precipitate Ca2+ or Mg2+ under some conditions and is less selective here. KCN would mask Cu2+, Zn2+, and similar ions, but it is highly toxic and must not be used in a school laboratory. The correct approach is to add triethanolamine before the buffer, swirl, and then titrate; if the endpoint remains poor, the Fe3+ can be removed by precipitation and filtration.