Hard Water Titration

Measure water hardness using EDTA chelation
Intermediate🕐30 min▲Moderate hazardanalyticalcomplex-ions

Summary

You’ll measure exactly how hard a water sample is by titrating1 it with EDTA, a molecule that chelates2 calcium and magnesium and snatches them from a colour-changing indicator. You’ll be able to explain what makes water “hard,” why EDTA out-competes the indicator, and why a six-point grip makes such a stable complex.

History

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.

Hazards & preparation

Warning

PPE: safety glasses and gloves; work with ventilation.

  • The ammonia buffer gives off pungent ammonia fumes — mix and use it in a well-ventilated area and avoid inhaling it.
  • EDTA and the indicator are low-hazard — avoid ingestion.

Disposal: the small, dilute volumes are low-hazard — flush down the drain with water. See the Safety page.

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

  1. Place 50mL of hard water (or tap water) into the flask.
  2. Add 2mL of ammonia buffer — this raises the pH to ~10, where the indicator works correctly.
  3. Add 3-4 drops of Eriochrome Black T. The solution turns wine-red (purple-red), showing free Ca²⁺ and Mg²⁺.
  4. Add EDTA dropwise from the burette, swirling between drops.
  5. As EDTA is added, it chelates the metal ions one by one.
  6. At the endpoint, the last metal ion is chelated and the indicator changes sharply from wine-red to clear blue.
  7. Record the volume of EDTA used.

Calculation: with 0.01 M EDTA and a 50 mL sample, each mL of EDTA ≈ 20 mg/L hardness as CaCO₃. (1 mL of 0.01 M EDTA = 0.01 mmol, chelating 0.01 mmol Ca²⁺ ≡ 1.0 mg CaCO₃; in 50 mL that’s 1.0 mg / 0.050 L = 20 mg/L.)

What you should see

With the buffer and indicator in, the sample glows a clear wine-red. As you drip in EDTA it stays red… stays red… then, right at the endpoint, one last drop flips the whole flask sharply to a clean blue. The sharper that colour change, the better your endpoint.

Symptom Likely cause Fix
No wine-red to start No free metal ions, or wrong pH Use genuinely hard water; add the ammonia buffer to reach pH ~10
Colour change is gradual, not sharp Titrating too fast near the end Slow to one drop at a time as blue tinges appear
Endpoint colour is muddy purple Old indicator, or Mg²⁺ absent Use fresh indicator; a trace of magnesium sharpens the change

The reactions

\[\ce{Ca^{2+}(aq) + EDTA^{4-}(aq) -> [Ca·EDTA]^{2-}(aq)}\]

The Eriochrome Black T indicator (In) forms a weaker complex with metal ions than EDTA does:

\[\ce{[Ca·In]^{} (red) + EDTA -> [Ca·EDTA] + In (blue)}\]

EDTA grabs the calcium and magnesium; once the last one is taken from the indicator, the freed indicator turns from red to blue, marking the endpoint.

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.

Questions to Explore

  1. What makes water “hard”? Why do calcium and magnesium cause scale and scum, but not sodium?

    Hint / answer

    Hardness is dissolved Ca²⁺ and Mg²⁺. Being doubly charged, they bridge soap molecules into insoluble scum and precipitate as carbonate scale. Sodium is singly charged and its salts stay dissolved, so it doesn’t scale or scum.

  2. Why does EDTA “win” against the indicator? And why doesn’t it steal the metal instantly when first added?

    Hint / answer

    EDTA wraps a metal ion at six points, forming a far more stable complex than the indicator’s looser grip — so it eventually takes the metal. It works through the free metal first; only once free metal is gone does it pull the last ions off the indicator, which is exactly when the colour flips.

  3. What is “hexadentate” binding and why so stable? Why does gripping at six sites beat single-site ligands like water?

    Hint / answer

    One molecule latching at six points is much harder to dislodge than six separate molecules each holding once — losing one grip doesn’t release the metal (the “chelate effect”). It’s also entropically favoured, so the complex is enormously more stable.

  4. Where does scale come from? Boiling hard water deposits calcium carbonate. Why does heat do this?

    Hint / answer

    Dissolved calcium bicarbonate breaks down on heating, driving off CO₂ and leaving insoluble calcium carbonate that plates out as scale. Heat shifts the equilibrium toward the less-soluble carbonate.

  5. Temporary vs. permanent hardness? Why can boiling remove one but not the other?

    Hint / answer

    Bicarbonate hardness decomposes on boiling to insoluble carbonate (removed — “temporary”). Sulfate hardness has no heat-driven breakdown and stays dissolved no matter how long you boil (“permanent”), needing softening chemistry instead.

Going further

  • Survey your water. Titrate tap water, bottled water, and rainwater and rank their hardness in mg/L CaCO₃.
  • Before and after softening. Titrate a hard sample, then soften it (see Water Softening) and titrate again to measure how much hardness you removed.

Footnotes

  1. Titration — Measuring an unknown concentration by adding a reagent of known concentration until the reaction just completes.↩︎

  2. Chelation — The binding of a metal ion by a molecule that grips it at several points at once, like a claw.↩︎