Hard Water Titration
Difficulty: Medium | Time: 30 minutes | Visual Impact: Medium
Historical Context
Water hardness was one of the first chemical properties of water to be measured systematically, because it had direct economic consequences - hard water consumed soap and deposited scale in industrial boilers. Early measurements used soap itself: the amount of soap needed to produce lasting lather in a water sample defined its hardness.
Analytical titration methods came in the early 20th century, but they required unpleasant procedures involving lengthy precipitation steps. The revolution came in 1946 when Gerold Schwarzenbach at the University of Zurich realized that EDTA could chelate calcium and magnesium tightly and predictably, and that indicator dyes could show when all metal ions were complexed. Suddenly, a water hardness measurement that had taken hours could be done in minutes with a burette.
The EDTA method remains standard practice worldwide. Expressed as mg/L CaCO₃ equivalent, water hardness is still routinely measured in municipal water treatment, brewing, aquarium keeping, and industrial cooling systems.
Materials
- EDTA disodium - 0.4g dissolved in 100mL water (0.01M solution)
- Calcium chloride - a small amount to make test hard water, or use real tap water
- Eriochrome Black T indicator - a few crystals dissolved in ethanol, or purchased as solution
- Ammonia buffer solution - 7mL ammonia + 57g ammonium chloride in 500mL water (pH ~10)
- Burette or marked dropper/syringe for EDTA
- 100mL conical flask or glass
Procedure
- Place 50mL of hard water (or 50mL tap water) into the flask.
- Add 2mL of ammonia buffer solution. This raises pH to ~10, where the indicator works correctly.
- Add 3-4 drops of Eriochrome Black T indicator. The solution turns wine-red (purple-red), indicating free Ca²⁺ and Mg²⁺.
- Add EDTA solution dropwise from a burette, swirling between drops.
- As EDTA is added, it chelates the metal ions one by one.
- At the endpoint, the last metal ion is chelated and the indicator changes sharply from wine-red to clear blue.
- Record the volume of EDTA used.
Calculation: With 0.01M EDTA and 50mL sample, each mL of EDTA used = 20 mg/L hardness as CaCO₃ (approximate, for a rough figure). Derivation: 1 mL of 0.01M EDTA contains 0.01 mmol, which chelates 0.01 mmol Ca²⁺ equivalent to 1.0 mg CaCO₃. In a 50 mL sample, that gives 1.0 mg / 0.050 L = 20 mg/L.
Reaction
\[\ce{Ca^{2+}(aq) + EDTA^{4-}(aq) -> [Ca·EDTA]^{2-}(aq)}\]
The Eriochrome Black T indicator (In) forms a weaker complex with metal ions than EDTA does:
\[\ce{[Ca·In]^{} (red) + EDTA -> [Ca·EDTA] + In (blue)}\]
The Science
EDTA is a hexadentate ligand - it has six donor atoms (two nitrogen and four oxygen) that can all coordinate to a single metal ion simultaneously, wrapping around it like a claw (the name “chelate” comes from the Greek for claw). The resulting complex is extremely stable, with stability constants many orders of magnitude higher than simple complexes.
Eriochrome Black T binds metal ions with a weaker complex than EDTA. When EDTA is added to a solution containing metal-indicator complexes, the EDTA “wins” the competition, stealing the metal ions away and releasing the free indicator, which is a different color. The endpoint of the titration is the moment when all metal ions have been captured by EDTA.