Molar Mass of Calcium Hydroxide (Ca(OH)₂)
Learn how chemists calculate the molar mass of Calcium Hydroxide (Ca(OH)₂), with a clear formula breakdown, worked steps, and study notes · IUPAC name: Calcium dihydroxide.
Quick answer
The molar mass of Calcium Hydroxide (Ca(OH)₂) is
74.092g/mol
One mole of Calcium Hydroxide therefore has a mass of 74.092 grams—the value you use for stoichiometry and laboratory preparation.
Reviewed for educational accuracy · Accuracy policy
- CAS Registry Number
- 1305-62-0
- PubChem CID
- 14777
- SMILES
- [OH-].[OH-].[Ca+2]
Step-by-step calculation
Let's find the molar mass of Calcium Hydroxide (Ca(OH)₂) together—step by step, as if you are seeing the formula for the first time.
Step 1 — Look at the chemical formula
The formula is Ca(OH)₂. Each letter stands for an element. The little number after a letter (the subscript) tells you how many atoms of that element are in one molecule or formula unit.
- 1 Calcium atom (Ca)
- 2 Oxygen atoms (O)
- 2 Hydrogen atoms (H)
Step 2 — Look up each atomic mass
Atomic mass comes from the periodic table. It is the average mass of one mole of atoms of that element, in grams per mole (g/mol). Think of it as the "price tag" for one mole of that element.
- Calcium (Ca) = 40.078 g/mol
- Oxygen (O) = 15.999 g/mol
- Hydrogen (H) = 1.008 g/mol
Step 3 — Multiply atoms × atomic mass
Why multiply? If one oxygen atom "costs" about 16 g/mol, then two oxygen atoms cost twice as much. Each element's contribution is: number of atoms × atomic mass.
- 1 × 40.078 = 40.078 g/mol (Calcium)
- 2 × 15.999 = 31.998 g/mol (Oxygen)
- 2 × 1.008 = 2.016 g/mol (Hydrogen)
Step 4 — Add the contributions
Why add? The molar mass of the whole compound is simply the total mass of every atom in the formula. Add each element's contribution:
40.078 + 31.998 + 2.016 = 74.092 g/mol
Step 5 — Final answer
Molar mass of Calcium Hydroxide = 74.092 g/mol
That means one mole of Calcium Hydroxide (Ca(OH)₂) has a mass of about 74.09 grams.
Quick summary
Read the formula → count atoms → look up atomic masses → multiply → add → report g/mol. For Ca(OH)₂, the total is 74.092 g/mol.
Common beginner mistakes
- Treating Ca(OH)₂ as a monoacidic base — it provides two hydroxide ions per formula unit.
- Forgetting that excess CO₂ clears limewater turbidity by forming soluble bicarbonate, not by 'undoing' the reaction.
- Confusing calcium hydroxide (Ca(OH)₂, 74.10 g/mol) with calcium oxide (CaO, 56.08 g/mol) or calcium carbonate (CaCO₃, 100.09 g/mol) in stoichiometry problems.
Memory trick
Remember the lime cycle: CaCO₃ → CaO → Ca(OH)₂ → CaCO₃, and be ready to write each step's balanced equation.
Mini practice
Without looking above, list the atoms in Ca(OH)₂ and write one multiplication line for the heaviest element. Then check your work against Step 3.
Real-world example
If a recipe asks for 0.100 mol of Calcium Hydroxide, mass needed = 0.100 × 74.092 = 7.409 g. That is how chemists turn a mole amount into a weighable sample.
Atomic contribution table
Each row shows how much mass one element contributes to the total for Ca(OH)₂.
| Element | Atoms | Atomic mass | Contribution | Mass % |
|---|---|---|---|---|
| Ca | 1 | 40.078 | 40.078 g/mol | 54.1% |
| O | 2 | 15.999 | 31.998 g/mol | 43.2% |
| H | 2 | 1.008 | 2.016 g/mol | 2.7% |
| Total molar mass | 74.092 g/mol | 100% | ||
Mass contribution chart
Count every atom in this formula, multiply by atomic mass, then add. That total is the molar mass used in lab weighing.
Download study sheets
Save a printable summary, revision sheet, practice worksheet, or laboratory reference for Calcium Hydroxide (Ca(OH)₂).
Practice this calculation
Without looking above, write the atom count for Ca(OH)₂, then compute the molar mass. Check your answer against 74.092 g/mol.
Next challenge: how many grams are in 0.250 mol of Calcium Hydroxide? Multiply 0.250 × 74.092 to get 18.523 g.
Physical and chemical properties
Physical properties
| Appearance | White powder or crystalline solid; forms a cloudy suspension or clear saturated solution (limewater) in water |
| Color | White |
| Odor | Odorless |
| State (STP) | Solid |
| Density | 2.211 g/cm³ |
| Melting point | ~580 °C (decomposes to CaO + H₂O before/at melting) |
| Boiling point | Decomposes before boiling |
| Solubility | 1.5 g/L water at 25 °C (decreases with increasing temperature, unusual for an ionic solid) |
| Crystal structure | Hexagonal (brucite-type layered structure) |
Chemical properties
| Classification | Alkaline earth metal hydroxide / moderately strong base |
| Family | Group 2 hydroxides (alkaline earth hydroxides) |
| Basicity | Strong base character in solution despite low solubility (fully dissociated Ca²⁺ and OH⁻) |
| Polarity | Ionic |
| Oxidation states | Ca: +2, O: −2, H: +1 |
Applications
Industrial uses
- Lime mortar, plaster, and whitewash in construction and heritage restoration
- Water and wastewater treatment (softening, pH adjustment, coagulation)
- Flue-gas desulfurization scrubbers (reacts with SO₂ to form calcium sulfite/sulfate)
- Sugar refining (clarification of raw sugar juice) and leather tanning (dehairing hides)
Laboratory uses
- Limewater as a qualitative test reagent for carbon dioxide
- Standard base for demonstrating precipitation and neutralization reactions
- Source of hydroxide ion in introductory acid-base titration demonstrations
Used to neutralize acidic industrial effluents and acid mine drainage; central reagent in flue-gas desulfurization to reduce SO₂ emissions from coal combustion.
Nixtamalization of corn improves nutritional availability of niacin; dilute Ca(OH)₂ solutions are used in some food processing and as a dental pulp-capping material in endodontics.
Preparation and production
Industrially and in the laboratory, slaked lime is made by adding water to calcium oxide (quicklime): CaO + H₂O → Ca(OH)₂, a strongly exothermic reaction that can boil off excess water if added too quickly. It can also be precipitated by mixing solutions of a soluble calcium salt with a strong base.
Quicklime is produced by calcining limestone in kilns above 900 °C (CaCO₃ → CaO + CO₂), then hydrated with a controlled amount of water to produce hydrated lime powder or lime putty on an industrial scale exceeding tens of millions of tonnes annually worldwide.
Important reactions of Calcium Hydroxide
CaO(s) + H₂O(l) → Ca(OH)₂(s)
- Reaction type
- Slaking (hydration)
- Conditions
- Room temperature, exothermic
- Explanation
- Quicklime reacts vigorously with water, releasing substantial heat and forming slaked lime; large-scale slaking can cause boiling and spattering if not controlled.
- Products
- Calcium hydroxide
- Why it matters
- Production of hydrated lime for construction and industry
Related ideas: Exothermic reactions · Hydration · Industrial chemistry
Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)
- Reaction type
- Precipitation / carbonation
- Conditions
- Room temperature, CO₂ gas bubbled through limewater
- Explanation
- Dissolved hydroxide reacts with carbon dioxide to precipitate insoluble calcium carbonate, turning clear limewater cloudy — the standard qualitative CO₂ test.
- Products
- Calcium carbonate and water
- Why it matters
- CO₂ detection, lime mortar hardening, water treatment
Related ideas: Qualitative analysis · Precipitation · Carbonation curing
Ca(OH)₂(aq) + 2 HCl(aq) → CaCl₂(aq) + 2 H₂O(l)
- Reaction type
- Acid–base neutralization
- Conditions
- Aqueous, room temperature
- Explanation
- Both hydroxide ions are neutralized by strong acid, forming calcium chloride and water in a 1:2 base-to-acid mole ratio.
- Products
- Calcium chloride and water
- Why it matters
- Acid neutralization in industrial effluent treatment
Related ideas: Neutralization · Stoichiometry · Diacidic bases
2 Ca(OH)₂(s) + 2 Cl₂(g) → Ca(OCl)₂(s) + CaCl₂(s) + 2 H₂O(l)
- Reaction type
- Disproportionation with halogen
- Conditions
- Cool, dilute Cl₂ gas exposure
- Explanation
- Chlorine reacts with slaked lime to produce bleaching powder, a mixture of calcium hypochlorite and calcium chloride used historically and industrially as a bleaching/disinfecting agent.
- Products
- Calcium hypochlorite (bleaching powder), calcium chloride, water
- Why it matters
- Bleaching powder manufacture, water disinfection
Related ideas: Disproportionation · Halogen chemistry · Disinfection chemistry
History and discovery
Lime plaster and mortar were used as early as 7000 BCE in the Middle East and extensively by the Romans, who perfected lime-pozzolan concrete. Joseph Black's mid-18th-century studies of limewater and "fixed air" (CO₂) were foundational to the discovery of gases as distinct chemical substances, helping launch modern pneumatic chemistry.
Known and used since prehistoric times for lime plaster and mortar; its role in carbon dioxide chemistry was clarified by Joseph Black in the 1750s.
Interesting facts
- Roman concrete used lime mortar mixed with volcanic ash, contributing to structures like the Pantheon that have survived nearly two millennia.
- Calcium hydroxide's low but non-negligible solubility (~1.5 g/L) means limewater is one of the few common strong bases you can make by simply stirring a base into water and filtering off the excess solid.
- The 'whitewash' historically used to coat barns and fences is essentially a Ca(OH)₂ suspension that hardens by carbonating into a durable CaCO₃ coating.
- Limewater's cloudiness with CO₂ was used by Joseph Black in the 1750s to help characterize 'fixed air,' contributing to the discovery of carbon dioxide.
Comparison with similar compounds
Ca(OH)₂ (74.10 g/mol) is only slightly soluble but strongly basic in solution, unlike NaOH (40.00 g/mol), which is fully soluble and gives much more concentrated hydroxide solutions for the same mass dissolved.
Storage, handling, and safety
Store hydrated lime powder in sealed, moisture-proof containers, as it slowly absorbs atmospheric CO₂ and moisture over time, degrading its reactivity. Limewater solutions should be kept tightly stoppered to minimize CO₂ absorption from air.
Ca(OH)₂ dust and solutions are caustic and irritating to skin, eyes, and respiratory tract. Wear gloves and eye protection; avoid generating airborne dust. Flush any skin or eye contact immediately and thoroughly with water.
Caustic irritant to skin, eyes, and respiratory tract; not classified as a strong hazard like NaOH but still requires standard base-handling precautions, especially avoiding dust inhalation and eye contact.
- Skin and eye irritation/burns from concentrated paste or solution
- Respiratory irritation from inhaling fine powder
- Exothermic reaction hazard when hydrating quicklime rapidly
Classification: GHS: Skin Irrit. 2, Eye Dam. 1, STOT SE 3
Exam notes and student tips
Exam notes
- Molar mass Ca(OH)₂ = 40.08 + 2(16.00 + 1.008) = 74.10 g/mol.
- Diacidic base: 1 mol Ca(OH)₂ furnishes 2 mol OH⁻ for neutralization stoichiometry.
- Limewater test: Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l); excess CO₂ redissolves precipitate as Ca(HCO₃)₂.
- Slaking reaction: CaO + H₂O → Ca(OH)₂ is strongly exothermic — do not confuse with dissolving, which is only mildly exothermic.
Student tips
- Remember the lime cycle: CaCO₃ → CaO → Ca(OH)₂ → CaCO₃, and be ready to write each step's balanced equation.
- For neutralization problems, treat Ca(OH)₂ like a diprotic-equivalent base: moles of H⁺ needed = 2 × moles of Ca(OH)₂.
- Link limewater cloudiness directly to the CO₂ test — a favorite qualitative analysis question.
Common mistakes
- Treating Ca(OH)₂ as a monoacidic base — it provides two hydroxide ions per formula unit.
- Forgetting that excess CO₂ clears limewater turbidity by forming soluble bicarbonate, not by 'undoing' the reaction.
- Confusing calcium hydroxide (Ca(OH)₂, 74.10 g/mol) with calcium oxide (CaO, 56.08 g/mol) or calcium carbonate (CaCO₃, 100.09 g/mol) in stoichiometry problems.
Misconceptions
- Slaked lime (Ca(OH)₂) and quicklime (CaO) are not the same substance — quicklime must react with water to become slaked lime.
- Limewater going clear again after excess CO₂ does not mean no reaction occurred — it means a different, soluble product (calcium bicarbonate) formed.
- Ca(OH)₂ is not a 'weak' base just because it is only slightly soluble — the portion that dissolves is essentially fully dissociated, making its solutions strongly basic.
Practice questions
1. Calculate the molar mass of Ca(OH)₂.
Show answer
40.08 + 2(16.00 + 1.008) = 74.10 g/mol
2. How many moles of HCl are needed to completely neutralize 7.41 g of Ca(OH)₂?
Hint: Ca(OH)₂ furnishes 2 mol OH⁻ per mole, so it needs 2 mol monoprotic acid.
Show answer
7.41 g ÷ 74.10 g/mol = 0.100 mol Ca(OH)₂; needs 2 × 0.100 = 0.200 mol HCl
3. What mass of CaCO₃ precipitates when excess CO₂ is bubbled through a solution containing 0.0500 mol Ca(OH)₂?
Show answer
1:1 mole ratio: 0.0500 mol CaCO₃ × 100.09 g/mol = 5.00 g (assuming no excess CO₂ redissolves it)
4. Why does limewater turn cloudy and then clear again if excess CO₂ is bubbled through it?
Show answer
Initial CO₂ forms insoluble CaCO₃ (cloudy); excess CO₂ converts CaCO₃ to soluble calcium bicarbonate, Ca(HCO₃)₂, clearing the solution.
Frequently asked questions about Calcium Hydroxide
74.10 g/mol.
Chemistry of Calcium Hydroxide
The sections above give the number you need for calculations. Here we look more closely at how Calcium Hydroxide (Ca(OH)₂) behaves chemically—so the molar mass connects to real reactions, properties, and laboratory practice.
Calcium hydroxide (Ca(OH)₂) has a molar mass of 74.10 g/mol (Ca 40.08 + 2 × [16.00 + 1.008]), formed when calcium oxide ("quicklime") reacts with water in a highly exothermic process traditionally called "slaking." The resulting white powder or paste — slaked lime — is only moderately soluble in water (~1.5 g/L at 25 °C), and its saturated solution, limewater, is a strong base (pH ~12.4) despite the low solubility, because the dissolved fraction dissociates essentially completely into Ca²⁺ and OH⁻.
Slaked lime has been central to construction chemistry for millennia: mixed with sand and water to make lime mortar and plaster, it slowly reabsorbs atmospheric CO₂ over weeks to years, converting back to calcium carbonate (CaCO₃) in a hardening process distinct from the rapid setting of modern Portland cement. This "lime cycle" — limestone (CaCO₃) → quicklime (CaO) → slaked lime (Ca(OH)₂) → limestone again — is one of the oldest known chemical cycles exploited by civilization, and it remains commercially important today.
Beyond construction, Ca(OH)₂ is a workhorse base in water and wastewater treatment (softening and pH correction), food processing (nixtamalization of corn, sugar refining), flue-gas desulfurization, and the classic limewater test for carbon dioxide taught in introductory chemistry, where CO₂ bubbled through clear limewater produces a cloudy CaCO₃ precipitate.
Ca(OH)₂ contains one Ca²⁺ cation ionically bonded to two OH⁻ hydroxide anions, balancing the +2 charge of calcium. Each hydroxide ion contributes one oxygen and one hydrogen. The formula shows that one mole of calcium hydroxide can furnish two moles of hydroxide ion upon complete dissociation, making it a diacidic base in neutralization stoichiometry even though its solubility is limited.
Ca(OH)₂ is a moderately strong base, essentially fully ionized in the small amount that dissolves. It readily absorbs atmospheric CO₂ to form insoluble CaCO₃ (the basis of the limewater test and lime mortar carbonation), neutralizes acids to form calcium salts and water, and undergoes double displacement with soluble carbonates or sulfates to precipitate CaCO₃ or CaSO₄. It reacts with chlorine gas to produce bleaching powder (calcium hypochlorite mixture) and is a key reagent in the Solvay process ammonia recovery step.
The Lime Cycle
Limestone (CaCO₃) is heated to quicklime (CaO), which reacts with water to form slaked lime (Ca(OH)₂); slaked lime then slowly reacts with atmospheric CO₂ to regenerate CaCO₃, completing a cycle exploited in construction and industry since antiquity. This closed loop links calcium hydroxide directly to limestone and lime chemistry.
Lime Mortar and Plaster Carbonation
Traditional lime mortar hardens not by drying but by slow carbonation: Ca(OH)₂ + CO₂ → CaCO₃ + H₂O occurs gradually as CO₂ diffuses through the porous material, which is why historic lime plaster can take months to fully cure but produces a self-healing, breathable, and highly durable finish still valued in heritage restoration.
The Limewater Test for CO₂
Bubbling a gas through clear limewater is the classic school test for carbon dioxide: CO₂ reacts with the dissolved Ca(OH)₂ to form a milky white precipitate of CaCO₃. Excess CO₂ redissolves the precipitate by forming soluble calcium bicarbonate, Ca(HCO₃)₂ — a subtlety often tested in exams.
Water and Wastewater Treatment
Lime softening uses Ca(OH)₂ to raise pH and precipitate magnesium and excess calcium hardness as hydroxide/carbonate solids, while also serving in coagulation, phosphorus removal, and heavy-metal precipitation at municipal treatment plants.
Food Chemistry: Nixtamalization
Soaking and cooking maize in a calcium hydroxide solution (nixtamalization) loosens the hull, improves niacin bioavailability, and imparts the characteristic flavor and dough properties of masa used for tortillas and tamales — a process dating back thousands of years in Mesoamerica.
Recalculate any formula with the molar mass calculator, compare atoms on the periodic table, or browse more compounds in the base library.
References and further reading
- NIST Chemistry WebBook: Ca(OH)₂ thermodynamic data
- PubChem CID 14777: Calcium hydroxide identifiers and properties
- National Lime Association: Lime production and construction applications