Calcium Hydroxide
Properties
| State | Solid (white powder) |
| Color | White |
| Solubility | Slightly soluble in water (1.5 g/L at 25°C; solubility decreases with temperature) |
| Melting Point | 580°C (decomposes) |
| Boiling Point | Decomposes before boiling |
About Calcium Hydroxide
Calcium hydroxide is the second step in the lime cycle — what you get when you add water to quicklime (CaO) in the slaking reaction CaO + H2O → Ca(OH)2, ΔH ≈ −63 kJ/mol, exothermic enough to boil the slaking water if you don't control the addition. Two pieces of chemistry make Ca(OH)2 useful. First, it has retrograde solubility: 1.73 g/L at 20 °C dropping to about 0.66 g/L at 100 °C, which is unusual among bases and explains why limewater forms a fresh CaCO3 precipitate when warmed in the presence of CO2. Second, the saturated solution (limewater) sits at pH 12.4, hard enough to do real base chemistry but soft enough that you can hold a hand in it without immediate harm — a useful balance for water treatment where you need to push pH up without overshooting. The classic CO2 detection test runs on this: bubble exhaled breath through limewater and CaCO3 precipitates as a white turbidity, Ca(OH)2 + CO2 → CaCO3↓ + H2O. The same reaction is what hardens lime mortar over time as atmospheric CO2 slowly carbonates it back to limestone, a process the Romans understood empirically and that still keeps Pantheon-era concrete standing. In Mesoamerica, slaked lime has been used for at least 3,500 years to nixtamalize maize: an alkaline cook at pH ~12 saponifies the kernel pericarp and frees bound niacin, the chemistry that prevents pellagra in maize-based diets.
Where you'll encounter it
Walk into a municipal water treatment plant and you'll find lime silos feeding Ca(OH)2 slurry into the head of the process to raise pH and precipitate hardness as calcium carbonate. In a tortilla factory the same compound — sold as cal — sits in steel vats simmering whole corn kernels overnight in the nixtamal step. Endodontists pack calcium hydroxide paste into root canals between visits because its pH 12+ environment kills residual bacteria. On a construction site, masons mix it into mortar and stucco, where carbonation slowly strengthens the joint over decades. And in any general chemistry teaching lab, the limewater bottle is the demonstration of choice for the CO2 test — students breathe through a straw and watch the milky CaCO3 develop.
Common Uses
- pH adjustment and hardness softening at the head of municipal water plants
- Nixtamalization of maize to free bound niacin and remove the pericarp
- Limewater as the qualitative test for CO2 in teaching and field labs
- Antimicrobial dressing in endodontic root-canal therapy at pH > 12
- Hydraulic component of lime mortar that re-carbonates to CaCO3 over decades
- Soil conditioner to neutralize acidity at higher reactivity than ground limestone
- Flocculant in sugar refining to remove organic acids from cane juice
- Scrubber reagent for SO2 removal from coal-fired flue gas
Safety Information
GHS H315/H318/H335 — causes skin irritation, serious eye damage, and respiratory irritation. The pH-12 alkalinity is the active hazard, and prolonged skin contact gives chemical burns that are slower to develop than acid burns and easy to underestimate. OSHA PEL is 5 mg/m³ (respirable) total dust. Splash goggles, nitrile gloves, and a dust mask are the right baseline. Avoid mixing with strong acids — the neutralization is exothermic and can spit corrosive aerosol.
This safety summary is for educational reference only and may not be complete. It is not a substitute for Safety Data Sheets (SDS), medical advice, or professional chemical safety guidance. Always consult appropriate SDS and qualified professionals before handling chemicals.