Calcium hydroxide
Calcium hydroxide, commonly known as slaked lime or hydrated lime, is an inorganic compound with the chemical formula Ca(OH)2, appearing as a white, odorless, crystalline powder that is slightly soluble in water to form a mildly alkaline solution with a pH of approximately 12.4.[1][2] It has a molecular weight of 74.09 g/mol, a density of 2.24–2.25 g/cm³, and decomposes upon heating above 580°C without a distinct melting point.[1][2] As a strong base, it reacts readily with acids and carbon dioxide to form calcium salts, and its solubility in water is low at about 1.73 g/L at 20°C but decreases with increasing temperature.[1][2] Industrially, calcium hydroxide is primarily produced through the hydration (or slaking) of calcium oxide, which is obtained by calcining high-purity limestone (calcium carbonate) at temperatures exceeding 900°C in a kiln, followed by the controlled addition of water to the quicklime in a process that generates significant heat.[2][3] This exothermic reaction, CaO + H2O → Ca(OH)2, yields a product typically containing at least 95% Ca(OH)2 for commercial grades, with the United States producing around 2.69 million metric tons annually as of 2018, supported by abundant domestic limestone resources.[3] Alternative methods, such as precipitation from aqueous solutions or extraction from industrial wastes like steel slag, are emerging but represent a minor fraction of global output due to higher costs.[2][4] Calcium hydroxide finds extensive applications across multiple sectors due to its alkaline properties and reactivity. In construction, it serves as a key ingredient in mortar, plaster, and cement production, where it facilitates binding and acts as a pH stabilizer.[1][5] In water and wastewater treatment, it is used for pH adjustment, precipitative softening, and neutralization of acidic effluents, including acid mine drainage remediation.[1][3] Food-grade variants, recognized as generally safe (GRAS) by the FDA under 21 CFR 184.1205, function as a firming agent in processes like nixtamalization for corn products and sugar refining, with usage rates up to 250 kg per ton of beet sugar.[2] In medicine, particularly endodontics, it is employed as an intracanal medicament for pulp capping and root canal disinfection due to its antimicrobial and tissue-regenerative effects.[1] Additionally, it is utilized in agriculture as a fungicide and soil conditioner under USDA organic standards (§205.601), in paper manufacturing for pulp processing, and in environmental applications like heavy metal stabilization.[2][6] Despite its utility, calcium hydroxide is corrosive and irritating to skin, eyes, and respiratory tissues upon exposure, with an oral LD50 in rats of 7,340 mg/kg, necessitating protective equipment during handling; it is not classified as carcinogenic but can cause burns in concentrated forms.[1][2] Its production contributes to CO2 emissions from limestone calcination, prompting research into low-carbon alternatives like electrochemical synthesis.[3][7]Properties
Physical properties
Calcium hydroxide appears as a white, odorless powder or colorless crystals in its pure form.[8][9] When dissolved in water, it forms a colorless solution known as limewater.[10] The solid has a density of 2.24 g/cm³ at 20 °C.[1] It decomposes at approximately 580°C before reaching a true melting point, releasing water vapor and forming calcium oxide.[11]Chemical properties
Calcium hydroxide acts as a strong base primarily due to its dissociation in water, releasing hydroxide ions that increase the solution's alkalinity.[1] The ionic dissociation follows the equilibrium \ce{[Ca](/page/CA)([OH](/page/Oh))2 ⇌ [Ca](/page/CA)^{2+} + 2[OH](/page/Oh)^-}, where the hydroxide ions (OH⁻) dominate the chemical behavior, making it effective for neutralization and pH adjustment applications.[12] In a saturated aqueous solution at 25°C, this results in a pH of approximately 12.4, confirming its strong basic character.[1] The compound is stable under normal conditions but undergoes thermal decomposition at approximately 580 °C under standard atmospheric conditions (where the partial pressure of water vapor is 1 atm), reverting to calcium oxide (CaO) and water (H₂O). This process highlights its reversible hydration-dehydration cycle, which is key to its industrial utility, though the exact onset can vary with environmental factors like pressure and particle size.[13] Calcium hydroxide is non-flammable and non-explosive, posing no ignition risk in standard handling scenarios, though it can react exothermically with water or acids to generate heat.[14] Its safety profile in this regard supports widespread use in construction and chemical processing without concerns for combustion hazards.[15]Preparation and production
Laboratory methods
In laboratory settings, calcium hydroxide is commonly synthesized through the hydration, or slaking, of calcium oxide (quicklime) with water, a process that produces a fine white powder suitable for small-scale educational or research applications. The reaction proceeds as follows: \text{CaO} + \text{H}_2\text{O} \rightarrow \text{Ca(OH)}_2 This exothermic reaction generates significant heat, often reaching temperatures above 100°C, necessitating the slow, controlled addition of water to calcium oxide—typically in a molar ratio of 1:1—to prevent boiling or splashing and ensure complete conversion without excessive agglomeration.[16][17] An alternative synthesis method employs a double displacement reaction between aqueous solutions of a calcium salt (such as calcium chloride) and sodium hydroxide, yielding calcium hydroxide as a precipitate. The balanced equation is: \text{CaCl}_2 + 2\text{NaOH} \rightarrow \text{Ca(OH)}_2 \downarrow + 2\text{NaCl} The solutions are mixed at room temperature, resulting in the immediate formation of a gelatinous precipitate due to the low solubility of calcium hydroxide. This method is advantageous for demonstrating precipitation reactions in teaching labs, as it avoids the handling of reactive quicklime.[11] Following synthesis by either method, purification involves filtration through filter paper or a Buchner funnel to separate the solid calcium hydroxide from unreacted reagents or soluble byproducts, such as sodium chloride in the double displacement approach. The collected precipitate is then washed with distilled water to remove residual impurities and dried under vacuum or in a low-temperature oven (around 60°C) to yield pure, anhydrous crystals without decomposition.[18] Laboratory preparation requires strict safety protocols due to the exothermic heat release during slaking, which can cause burns, and the generation of caustic dust that irritates skin, eyes, and respiratory tract. Operators must wear protective equipment including gloves, safety goggles, lab coats, and respirators; perform the reaction in a well-ventilated fume hood; and have cooling measures like ice baths available for temperature control.[19]Industrial production
Calcium hydroxide is primarily produced on an industrial scale through the slaking of quicklime (CaO), which is obtained by calcining high-calcium limestone in rotary or vertical kilns at temperatures exceeding 900°C. This quicklime is then hydrated in large-scale slakers—typically vertical or horizontal reactors—by controlled addition of water, forming a slurry or dry powder of calcium hydroxide with yields typically ranging from 95% to 98%, accounting for minor losses due to impurities and unreacted material.[17] The slaking reaction is highly exothermic, releasing approximately 490 BTU per pound of quicklime for high-reactivity variants, and industrial processes incorporate heat recovery systems to capture this energy, often via steam generation or indirect heat exchangers, enhancing overall energy efficiency by up to 20-30% in integrated facilities.[17][20] Global production of calcium hydroxide reached approximately 34 million tons in 2024, with major contributions from China, the world's leading lime producer accounting for over 50% of output, and the United States, which produced approximately 2.5 million metric tons of hydrated lime.[21][22][23] Recent trends emphasize sustainable practices, including sourcing from recycled lime recovered from industrial wastes like steel slag, reducing reliance on virgin limestone and minimizing environmental footprints. Advancements in production integrate carbon capture technologies during the upstream calcination step, such as indirectly heated carbonate looping (IHCaL) processes, which can capture over 90% of CO₂ emissions from limestone decomposition, enabling near-zero net emissions in modern lime plants.[24]Crystal structure
Molecular arrangement
Calcium hydroxide, also known as portlandite, adopts a layered structure in its solid state, crystallizing in the hexagonal (trigonal) crystal system with space group P-3m1 (No. 164).[25] In this arrangement, each Ca²⁺ ion is octahedrally coordinated by six OH⁻ groups, forming CaO₆ octahedra that share edges to create infinite two-dimensional sheets parallel to the (001) plane.[26] These sheets are stacked along the c-axis, with adjacent layers connected via hydrogen bonds between the hydroxyl groups.[27] The Ca-O bond length within the octahedra is approximately 2.36 Å, reflecting the ionic nature of the coordination.[28] Hydrogen bonds between layers involve O-H···O interactions, with O···O distances around 3.0 Å, contributing to the overall stability of the structure.[27] The unit cell parameters are a = b ≈ 3.59 Å and c ≈ 4.90 Å, with α = β = 90° and γ = 120°, resulting in a compact layered motif that accommodates the larger ionic radius of Ca²⁺ compared to analogous compounds.[25] This polymeric structure is analogous to that of brucite, Mg(OH)₂, which also features octahedral MO₆ layers (M = metal) linked by hydrogen bonds, but calcium hydroxide exhibits larger lattice parameters (a ≈ 3.14 Å, c ≈ 4.77 Å for brucite) due to the greater ionic radius of Ca²⁺ (1.00 Å) versus Mg²⁺ (0.72 Å), leading to increased interlayer spacing.[29]Polymorphism
Calcium hydroxide, also known as portlandite in its mineral form, primarily exists in a hexagonal crystal structure with space group P-3m1, which is stable under ambient temperature and pressure conditions. This polymorph features layers of calcium atoms coordinated in octahedral geometry by hydroxide ions, forming the basis for its common occurrence in hydrated cement and natural settings.[25] At elevated pressures exceeding approximately 6 GPa, portlandite undergoes a reversible phase transition to a high-pressure polymorph with monoclinic symmetry (space group I121), characterized by a distorted sevenfold coordination around calcium atoms due to shifts in the hydroxide layers. This structural change results in a volume reduction of about 5.8% and represents an intermediate state before further transformations at higher pressures and temperatures, such as to a P21/c phase above 23 GPa.[30] An amorphous phase of calcium hydroxide forms under conditions of rapid precipitation or in nanoscale particles (typically 10-30 nm in size), arising from prenucleation clusters and dense liquid precursors before crystallizing into the stable hexagonal form. This amorphous variant is metastable, exhibiting higher solubility than crystalline portlandite, and tends to transform into the hexagonal polymorph over time, though it can be temporarily stabilized by additives.[31] These polymorphic transitions, particularly the pressure-induced shift from hexagonal to monoclinic forms, hold relevance for understanding the behavior of hydrous minerals in geological processes, such as water transport in subduction zones where pressures surpass several gigapascals.[30]Solubility and aqueous behavior
Solubility data
Calcium hydroxide exhibits low solubility in water, with its dissolution showing an inverse temperature dependence unusual among ionic compounds. The solubility is approximately 1.89 g/L at 0°C, decreasing to 1.73 g/L at 20°C and further to 0.66 g/L at 100°C.[32][8] This behavior is quantified by the solubility product constant, K_{sp} = [\ce{Ca^{2+}}][\ce{OH^{-}}]^2 = 5.02 \times 10^{-6} at 25°C, reflecting the equilibrium \ce{Ca(OH)2(s) ⇌ Ca^{2+}(aq) + 2OH^{-}(aq)}.[33] The solubility is further reduced by the common ion effect, where the presence of additional \ce{Ca^{2+}} or \ce{OH^{-}} ions shifts the equilibrium toward the solid phase in accordance with Le Châtelier's principle./19%3A_Equilibrium/19.15%3A_Common_Ion_Effect) In non-aqueous media, calcium hydroxide remains poorly soluble, being insoluble in alcohols such as ethanol. However, its solubility increases in sugar solutions, such as those containing glucose or sucrose, due to the formation of soluble complexes between the hydroxide ions and carbohydrate molecules.[9][34]Limewater
Limewater is a saturated aqueous solution of calcium hydroxide (Ca(OH)2) prepared at room temperature, resulting in a clear, colorless liquid with a pH of approximately 12.4 due to its weakly basic nature. This solution is distinct from more concentrated slurries, as its saturation is limited by the low solubility of calcium hydroxide in water, typically around 1.73 g/L at 20°C, ensuring minimal undissolved particles in the final product. To prepare limewater, an excess of calcium hydroxide powder is added to distilled or tap water in a clean container, such as a glass jar, and vigorously shaken or stirred for 1–2 minutes to facilitate dissolution.[35] The mixture is then allowed to stand undisturbed for several hours or overnight, allowing undissolved solids to settle at the bottom. The clear supernatant liquid is carefully decanted or filtered through filter paper to obtain the saturated solution, avoiding disturbance of the sediment. For storage, limewater must be kept in stoppered or sealed bottles to minimize exposure to atmospheric carbon dioxide, which would otherwise react to form insoluble calcium carbonate and cloud the solution over time.[36] A classical application of limewater is as a qualitative test for carbon dioxide gas, where bubbling CO2 through the solution produces immediate turbidity from the white precipitate of calcium carbonate, as shown in the reaction: \ce{Ca(OH)2 (aq) + CO2 (g) -> CaCO3 (s) + H2O (l)} This visible change confirms the presence of CO2, a method long used in laboratory demonstrations and gas identification experiments.[37] In excess CO2, the precipitate may redissolve to form soluble calcium bicarbonate, clearing the solution again, but the initial turbidity serves as the diagnostic indicator.[38] In contemporary settings, limewater finds use in pH adjustment for aquariums and fish ponds, where it helps buffer water against acidic fluctuations to maintain optimal conditions for aquatic life, typically targeting a pH range of 7.0–9.0.[39] It is also employed in laboratory qualitative analysis for detecting carbon dioxide in samples, such as exhaled breath or chemical reactions, providing a simple, visual confirmation without complex instrumentation.[40]Chemical reactions
Acid-base reactions
Calcium hydroxide acts as a strong base in neutralization reactions with strong acids, producing the corresponding calcium salt and water. For example, the reaction with hydrochloric acid is represented by the equation: \mathrm{Ca(OH)_2 + 2HCl \rightarrow CaCl_2 + 2H_2O} This complete neutralization occurs because both the acid and base fully dissociate in aqueous solution, liberating heat and forming a neutral salt solution.[41] In carbonation reactions, calcium hydroxide reacts with carbon dioxide to form calcium carbonate and water, as shown: \mathrm{Ca(OH)_2 + CO_2 \rightarrow CaCO_3 + H_2O} This process serves as the chemical foundation for lime softening in water treatment, where added calcium hydroxide precipitates hardness-causing ions as carbonates.[42][43] Calcium hydroxide also undergoes sulfation with sulfur dioxide, yielding calcium sulfite and water: \mathrm{Ca(OH)_2 + SO_2 \rightarrow CaSO_3 + H_2O} This reaction is central to dry and semidry flue gas desulfurization processes in coal-fired power plants, where calcium hydroxide sorbs SO₂ from emissions to reduce acid rain precursors. Unlike stronger bases such as sodium hydroxide, calcium hydroxide is non-corrosive to iron and steel due to the formation of a passivating calcium carbonate layer on metal surfaces, which physically barriers corrosive agents.[44][45]Thermal and other reactions
Calcium hydroxide decomposes thermally into calcium oxide and water vapor upon heating to 580 °C, following the endothermic reaction: \mathrm{Ca(OH)_2 \rightarrow CaO + H_2O} This decomposition process is integral to lime reburning in industrial cycles, where calcium hydroxide is converted back to quicklime for reuse in applications such as construction materials production.[46] In high-temperature environments, calcium hydroxide reacts with silica to form calcium silicate hydrates, exemplified by the pozzolanic reaction: \mathrm{Ca(OH)_2 + SiO_2 \rightarrow CaSiO_3 \cdot H_2O} This reaction contributes to the strength development in cementitious materials by binding components during hydration.[47] Calcium hydroxide demonstrates photochemical stability, showing no significant degradation under exposure to light or photooxidation conditions.[1] Furthermore, calcium hydroxide exhibits redox inertness under typical conditions, as the +2 oxidation state of calcium remains stable and does not readily participate in oxidation-reduction processes.[1]Applications
Construction and materials
Calcium hydroxide, commonly known as slaked lime, serves as a fundamental binder in lime-based mortars and plasters used extensively in construction. When mixed with aggregates like sand, it forms a workable paste that hardens through carbonation, where calcium hydroxide reacts with atmospheric carbon dioxide to produce calcium carbonate, creating a durable, porous matrix:\ce{Ca(OH)2 + CO2 -> CaCO3 + H2O}
This process results in a flexible material that accommodates building movements and allows moisture vapor to escape, preventing trapped dampness.[48][49] Lime mortars are classified as non-hydraulic or hydraulic based on their setting mechanisms. Non-hydraulic limes, derived from high-calcium limestone, set solely through carbonation and require exposure to air for hardening, making them ideal for internal or sheltered applications where breathability is essential.[50] In contrast, hydraulic limes contain impurities like silica and alumina from the source limestone, enabling them to set via hydrolysis in the presence of water—even underwater—while still undergoing carbonation for long-term strength, suitable for exposed or damp environments.[51][52] Historically, calcium hydroxide has been integral to enduring structures, including Roman concrete and medieval cathedrals, prized for its breathable and self-healing attributes. In Roman opus caementicium, quicklime was hot-mixed with pozzolanic aggregates, forming lime clasts that react with water over time to fill cracks with calcium carbonate, contributing to the longevity of structures like the Pantheon.[53][54] Medieval builders employed slaked lime mortars in cathedrals such as Notre-Dame, leveraging its vapor permeability to protect stonework from salt crystallization and its ability to recrystallize and seal fissures.[55] In modern construction, slaked lime is often blended with Portland cement to enhance mortar performance, improving workability by increasing plasticity and reducing bleeding, while its slower hydration minimizes shrinkage cracking through distributed micro-cracks that self-heal via carbonation.[56][57] These hybrid mortars offer better adhesion to substrates and flexibility compared to pure cement mixes, reducing long-term damage in masonry.[58] Slaked lime putty, produced by prolonged hydration and aging of quicklime, plays a key role in traditional and restorative applications, particularly for frescoes and heritage preservation. Aged for years to refine particle size and promote cohesion, it forms a smooth, adhesive base for intonaco plaster in buon fresco technique, allowing pigments to bind chemically during carbonation.[59] In building restoration, lime putty mortars replicate historic formulations, enabling compatible repairs to ancient masonry without introducing incompatible rigidity.[60][61]