Fact-checked by Grok 2 weeks ago

Calcium

Calcium is a chemical element with the symbol Ca and atomic number 20, classified as a soft, silvery-white alkaline earth metal in group 2 of the periodic table. It has an atomic mass of 40.078 u, an electron configuration of [Ar] 4s², and occurs naturally in compounds such as limestone (calcium carbonate) and gypsum (calcium sulfate dihydrate), making it the fifth-most-abundant element by mass in Earth's crust at over 3%. Essential for life, calcium plays critical roles in biomineralization for bones, teeth, and shells, as well as in cellular signaling, muscle contraction, and nerve transmission. Physically, calcium is a solid at with a of 1.54 g/cm³, a of 842°C (1115 K), and a of 1484°C (1757 K). The metal tarnishes in air due to oxidation but is protected by a thin layer, and it reacts vigorously with to produce gas and . Calcium was first isolated in its metallic form in 1808 by English chemist through the of a mixture of (CaO) and mercuric , though its compounds like had been used since ancient times for and ; the name derives from the Latin calx, meaning . In nature, calcium is never found uncombined but constitutes key minerals including (CaF₂) and (calcium phosphate), with significant abundances in (about 4220 ppm by weight) and the (around 1.4% by weight, primarily in bones as ). Biologically, over 99% of bodily calcium is stored in the for and as a reservoir, while the remaining 1% in and tissues regulates processes like blood clotting, secretion, and activation; ionized calcium in serum normally ranges from 8.8 to 10.4 mg/dL. Deficiency can lead to conditions like , underscoring its indispensable role in health. Industrially, calcium metal serves as a in extracting metals like and , and as an alloying agent to strengthen aluminum, , and ; its compounds dominate applications, including for production, , and manufacturing. Calcium isotopes, such as ⁴²Ca and ⁴⁶Ca, are employed in nutritional to study and . Ongoing explores calcium in advanced batteries for its abundance and potential for high-energy density storage.

Characteristics

Classification

Calcium is a with 20, positioned in group 2 and period 4 of the periodic table, classifying it as an . Its ground-state is [Ar] 4s², with the two electrons occupying the 4s orbital, which contributes to its characteristic chemical behavior within the group. Alkaline earth metals like calcium are highly reactive, owing to their tendency to lose both electrons and form divalent cations (M²⁺), resulting in a +2 in most compounds./08:_Chemistry_of_the_Main_Group_Elements/8.05:_Group_2_The_Alkaline_Earth_Metals/8.5.02:_Alkaline_Earth_Metals'_Chemical_Properties) This reactivity is enhanced by their low ; calcium specifically has an electronegativity of 1.00 on the Pauling scale, making it prone to with nonmetals.80142-5) Compared to other group 2 elements, calcium's atomic radius of 197 pm is intermediate, larger than magnesium (160 pm) but smaller than strontium (215 pm), illustrating the trend of increasing size down the group due to additional electron shells. Its ionization energies reflect this position: the first is 589.8 kJ/mol and the second is 1145 kJ/mol, values higher than magnesium's (738 kJ/mol first, 1451 kJ/mol second) but lower than strontium's (549 kJ/mol first, 1064 kJ/mol second), consistent with decreasing ionization energy as atomic size increases.

Physical properties

Calcium is a soft, silvery-white metal that rapidly tarnishes in air to form a grayish-white layer. This malleable and ductile nature allows it to be easily shaped or drawn into wires, characteristic of its position among the alkaline earth metals. The density of elemental calcium is 1.55 g/cm³ at 20 °C. It has a of 842 °C (1115 K) and a of 1484 °C (1757 K). Calcium exhibits a face-centered cubic (FCC) crystal structure with a of 558.8 pm at 20 °C. Elemental calcium demonstrates good thermal conductivity of 201 W/(m·K) and electrical resistivity of 34 nΩ·m at 20 °C, reflecting its and behavior. These properties contribute to its utility in applications requiring heat dissipation or electrical conduction, though its reactivity limits direct use.

Chemical properties

Calcium is a highly reactive that tarnishes rapidly in air, forming a gray-white coating of (CaO) and calcium (Ca₃N₂) due to its affinity for and . Finely divided calcium is pyrophoric and ignites spontaneously in air, burning with an intense white flame. It reacts vigorously with at , liberating gas and forming (Ca(OH)₂), as represented by the equation: \ce{Ca + 2H2O -> Ca(OH)2 + H2} This exothermic reaction demonstrates calcium's strong reducing properties. The predominant oxidation state of calcium is +2, resulting from the facile loss of its two 4s valence electrons, enabled by relatively low successive ionization energies (first: 589.8 kJ/mol; second: 1145.3 kJ/mol). This electron configuration promotes ionic bonding in calcium compounds, where the Ca²⁺ cation electrostatically interacts with anions, favoring the formation of ionic lattices over covalent structures. Calcium's position in the reactivity series places it above magnesium but below sodium, underscoring its ability to displace hydrogen from dilute acids, yielding hydrogen gas and soluble calcium salts. Calcium readily forms compounds such as oxides (e.g., CaO), halides (e.g., CaCl₂), and sulfates (e.g., ), with solubility trends governed by and effects. Most calcium salts exhibit good solubility in , but notable exceptions include (CaCO₃), which is insoluble (K_{sp} = 2.8 \times 10^{-9}), and calcium sulfate dihydrate (, ), which has limited (K_{sp} = 2.4 \times 10^{-5} at 25°C). The standard reduction potential for the Ca²⁺/Ca half-cell is -2.87 V versus the standard hydrogen electrode, indicating calcium's potent reducing capability and thermodynamic tendency to oxidize.

Isotopes

Calcium has 25 known isotopes, with mass numbers ranging from 34 to 58. Six of these isotopes are and occur naturally on , while the remainder are radioactive with half-lives spanning from microseconds to over 100,000 years. The isotopes dominate the natural composition of calcium, with variations in abundance reflecting both and minor cosmogenic production. The isotopes of calcium, along with their approximate natural abundances in terrestrial materials, are listed below:
IsotopeNatural Abundance (%)
⁴⁰Ca96.941
⁴²Ca0.647
⁴³Ca0.135
⁴⁴Ca2.086
⁴⁶Ca0.004
⁴⁸Ca0.187
These abundances are based on measurements from standard geological samples and show minor variations due to mass-dependent processes in natural systems. Among the radioactive isotopes, ⁴¹Ca has a half-life of 99,400 years and is produced cosmogenically through neutron capture on ⁴⁰Ca; it serves as a long-lived tracer in studies of calcium metabolism and has applications in paleodietary research by enabling dating of ancient bone samples beyond the range of radiocarbon methods. Another notable example is ⁴⁵Ca, with a half-life of 162.7 days, which is widely used as a radioactive tracer in biological and medical investigations of calcium uptake and transport in living organisms. The other radioactive isotopes decay primarily via beta emission or electron capture, with most having short half-lives that limit their natural occurrence. The stable isotopes of calcium are primarily synthesized through , where explosive stellar events produce neutron-rich nuclei via rapid neutron capture (r-process) and silicon burning, supplemented by contributions from asymptotic giant branch stars through the slow neutron capture (s-process). Calcium isotopes find specialized applications in and . For instance, ⁴⁴Ca is employed in theoretical and experimental studies of , a rare process that could reveal insights into properties if observed. Unlike heavier elements, calcium has no fissile isotopes capable of sustaining a .

History

Discovery and isolation

Lime, or (CaO), has been utilized as a compound since prehistoric times, with evidence of its use in dating back to approximately 4000 BCE in , where it served as a binding and in the building of early pyramids, such as those at around 2600 BCE. This early recognition of lime's properties as a reactive material for highlights its role in foundational , though the elemental nature of calcium remained unrecognized for millennia. Efforts to isolate calcium as a pure began in the early amid advances in . In 1803, Swedish chemists and Magnus Martin Pontin conducted of dissolved in mercury, producing a calcium-mercury amalgam but failing to obtain the pure metal due to the amalgam's stability. Inspired by their work, British chemist pursued similar experiments at the Royal Institution in . Davy successfully isolated metallic calcium in 1808 through electrolysis of a mixture of moist (CaO) and mercuric (HgO), using a mercury to form a calcium amalgam, which he then heated in a to distill away the mercury and yield impure calcium metal. This breakthrough, announced in a series of lectures to the Royal Society, marked the first production of calcium in its al form and demonstrated the power of for decomposing refractory s, a technique Davy had previously applied to and sodium. The resulting calcium appeared as a soft, silvery-white solid that rapidly tarnished in air, confirming its reactivity and placement among the alkaline earth metals.

Etymology and early recognition

The term "calcium" originates from the Latin word , meaning lime or limestone, reflecting the element's long association with calcium oxide (CaO), commonly known as quicklime. This nomenclature was proposed by British chemist in 1808, following his electrolytic isolation of the metal from a mixture of lime and mercuric oxide, to denote the new element distinct from earlier uses of "calx" for calcined substances like ashes. Prior to this, in alchemical and early chemical traditions, calcium compounds were referred to as "calcareous earth" or "earth of lime," encompassing materials such as limestone and chalk that were manipulated in processes like calcination to produce lime for various applications. Human recognition of calcium-based materials dates back millennia, with evidence of their use in construction predating the element's identification. (CaSO₄·2H₂O), a hydrated calcium sulfate, was employed as plaster in settlements, notably at in modern-day around 7500 BCE, where it was mixed with for wall coatings and decorative purposes. By the Roman era, quicklime played a pivotal role in durable formulations; for instance, the Pantheon's massive dome, completed in 126 CE under Emperor , utilized pozzolanic incorporating quicklime, , and aggregates, enabling self-healing properties through hot-mixing techniques that formed lime clasts for crack repair. In the pre-modern scientific framework, calcium was understood as an elemental "earth" rather than a metal. , in his 1789 treatise Traité Élémentaire de Chimie, classified "terre calcaire" (calcareous earth) among the simple substances, listing it alongside other earths like magnesia and baryta based on their resistance to decomposition and role in forming salts. , who assisted Davy in early experiments, further refined the terminology in 1810 by adopting and promoting "calcium" in , aligning it with emerging conventions for metallic elements derived from earths. This progression from alchemical "earth of " to a recognized underscored calcium's foundational role in both practical arts and systematic chemistry.

Occurrence and production

Natural occurrence

Calcium is produced primarily through stellar nucleosynthesis in massive stars and is relatively abundant in the universe compared to heavier elements. Its cosmic distribution reflects contributions from Big Bang nucleosynthesis for lighter elements and subsequent stellar processes for heavier ones like calcium. On Earth, calcium ranks as the fifth most abundant element in the crust, comprising approximately 3.6% by mass. In the oceans, it is present at a concentration of about 0.4 g/L primarily as Ca²⁺ ions, dissolved from continental weathering and contributing to marine chemistry. Calcium occurs naturally in a variety of minerals, with (CaCO₃) being the most prevalent, accounting for roughly 4% of the and forming extensive deposits. Other significant minerals include (CaMg(CO₃)₂), which is common in sedimentary rocks; (CaSO₄·2H₂O), found in deposits; and (CaF₂), a key source of . In the biosphere, calcium is highly concentrated in the exoskeletons, shells, and bones of organisms, where it forms structural components like in marine shells and in vertebrate bones. Marine organisms, such as mollusks, corals, and , extract calcium from to build these structures, facilitating its distribution across ecosystems. Within the solar system, calcium is a major component of chondritic meteorites, representing primitive solar material with typical CaO contents around 1–2%, and lunar rocks, particularly anorthosites in the highlands, where CaO levels can reach up to 10–15%.

Industrial production

The primary industrial method for producing elemental calcium is the aluminothermic reduction of (CaO), a process conducted under at approximately 1200°C. In this reaction, aluminum serves as the , following the equation: $3\text{CaO} + 2\text{Al} \rightarrow 3\text{Ca} + \text{Al}_2\text{O}_3 This yields calcium metal that can be further processed for higher purity. An alternative , electrolysis of molten (CaCl₂), has been employed since the 1890s and remains in use for producing calcium of about 99% purity. The process involves electrolyzing the molten with a sacrificial , where calcium deposits at the and gas is evolved at the . Global annual production of calcium metal is approximately 10,000 tons as of 2023, with the majority originating from and . Following initial production, calcium is purified by to achieve 99.9% purity, removing volatile impurities and residual oxides. The raw material for these processes is typically derived from the of (CaCO₃), which decomposes to form CaO and releases CO₂ as a : CaCO₃ → CaO + CO₂.

Geochemical cycling

Calcium plays a central role in the carbonate-silicate cycle, a fundamental geochemical process that regulates Earth's long-term and climate stability. During chemical of calcium-rich silicate rocks, such as plagioclase feldspars, derived from atmospheric CO₂ reacts with minerals to release Ca²⁺ ions into solution, which are then transported to the oceans via rivers. In marine settings, these ions combine with (HCO₃⁻) to form (CaCO₃), primarily through biogenic by organisms like coccolithophores and . This effectively sequesters carbon, with global riverine inputs delivering approximately 0.65 Gt of Ca per year to the oceans from both riverine and minor hydrothermal sources. The reverse process occurs through volcanic and metamorphic , releasing CO₂ and enabling renewed . In the oceanic realm, the calcium cycle involves biological uptake, , and physical mixing. Calcifying incorporate Ca²⁺ into CaCO₃ structures, which sink as and accumulate as sediments, forming vast deposits over geological time. A significant portion—over 80%—of produced CaCO₃ dissolves in the or upper sediments due to increasing pressure and decreasing with depth, but the net burial balances inputs at roughly 0.6 Gt Ca per year. circulates dissolved calcium back to surface waters, sustaining in nutrient-limited regions. The average of calcium in is approximately 1 million years, reflecting its conservative behavior and well-mixed distribution throughout the global ocean. Volcanic processes contribute to calcium inputs primarily through the formation and subsequent of at mid-ocean ridges and arcs, where CaO constitutes about 10-12% of composition. The from these sources, including hydrothermal alteration, is estimated at around 0.1 Gt Ca per year, providing a steady supply that supports the cycle's continuity. On land, the influences local calcium dynamics: absorb Ca²⁺ from soils for structural roles in cell walls, while herbivores and decomposers recycle it through and , with minimal net loss over short timescales but significant observable in isotopic signatures. Human activities have introduced perturbations to this ancient cycle, notably through intensified mining of and the production of , which consumes over 4 Gt of CaCO₃ annually and releases about 1 Gt of CO₂ per year via (CaCO₃ → CaO + CO₂). This process accelerates calcium mobilization and atmospheric CO₂ enrichment, bypassing natural rates and enhancing the reverse flux in the carbonate-silicate cycle, with potential long-term implications for ocean alkalinity and .

Calcium compounds

Inorganic compounds

Calcium oxide (CaO), commonly known as quicklime, is a white, caustic, alkaline solid produced primarily by the of or other sources. The reaction involves heating to temperatures exceeding 900°C, yielding and gas:
\ce{CaCO3 ->[Δ] CaO + CO2}.
Quicklime exhibits high reactivity with water, undergoing a strongly exothermic to form (slaked lime):
\ce{CaO + H2O -> Ca(OH)2}.
It has a of 2,613°C and is widely used as a key ingredient in production due to its binding properties upon . Calcium carbonate (CaCO₃) occurs naturally in several polymorphic forms, including the stable (rhombohedral structure) and the metastable (orthorhombic structure), both of which are principal components of , , and seashells.
It is sparingly soluble in , with a solubility product constant (Ksp) of approximately 3.7 × 10⁻⁹ at 25°C, which governs its precipitation in aqueous environments and contributes to its role in geological formations like stalactites.
Calcium carbonate decomposes thermally above 840°C to form calcium oxide and , mirroring the preparation of quicklime. Among the calcium halides, (CaCl₂) is highly hygroscopic and deliquescent, readily absorbing moisture from the air to form hydrates such as the dihydrate (CaCl₂·2H₂O).
It is typically prepared by reacting with :
\ce{CaCO3 + 2HCl -> CaCl2 + H2O + CO2}.
has a of 772°C and is soluble in to the extent of about 74 g/100 mL at 20°C, making it useful in de-icing applications due to its exothermic dissolution.
In contrast, (CaF₂), known as or fluorspar, is nearly insoluble in (Ksp ≈ 3.9 × 10⁻¹¹ at 25°C) and adopts a cubic .
It occurs naturally as a and has a high of 1,418°C, contributing to its use in and for its transparency to radiation. Calcium sulfate exists predominantly as the dihydrate (CaSO₄·2H₂O), or , a soft, white with a Mohs of 2, commonly found in sedimentary deposits.
Upon heating to around 150°C, loses three-quarters of its of to form hemihydrate (CaSO₄·0.5H₂O), known as of , which sets rapidly upon rehydration to regenerate the dihydrate structure.
The dihydrate has low in (about 0.21 g/100 mL at 20°C) and is chemically stable under neutral conditions. Calcium nitrate (Ca(NO₃)₂) is a highly water-soluble ( ≈ 121 g/100 mL at 20°C), often encountered as the tetrahydrate (Ca(NO₃)₂·4H₂O), which is colorless and deliquescent.
It can be synthesized by dissolving in or via the reaction of with .
The anhydrous form melts at 561°C and decomposes at higher temperatures, while its solutions are neutral to slightly acidic due to partial .

Organocalcium compounds

Organocalcium compounds, featuring direct calcium-carbon bonds, are exceedingly rare owing to the element's pronounced reactivity, rendering most species highly unstable toward air and moisture; consequently, they must be synthesized and manipulated under strict inert atmospheric conditions. These compounds exhibit bonding that is predominantly ionic, reflecting calcium's electropositive nature, yet displays subtle covalent character in the Ca-C interaction, which contributes to their enhanced nucleophilicity compared to analogous magnesium species. In the gas phase, computational and spectroscopic studies reveal that many such compounds favor dimeric structures, stabilized by bridging ligands or metal-metal interactions. Prominent examples include calcium cyclopentadienide (\ce{Cp2Ca}), a metallocene analog of featuring two \eta^5-coordinated rings, first synthesized in the late 1960s by the reaction of calcium metal with in , yielding an insoluble polymeric solid. Another key representative is calcium acetylide (\ce{CaC2}), formed via the high-temperature reaction of elemental calcium with carbon (typically above 2000°C), which serves as the primary industrial organocalcium and is hydrolyzed to generate gas for . The practical applications of organocalcium compounds remain confined largely to academic research, where their strong nucleophilicity enables roles as initiators in anionic polymerization processes, such as the ring-opening polymerization of lactones or the synthesis of polyolefins, often outperforming traditional Grignard reagents due to reduced side reactions under optimized conditions. Advancements in stabilization have yielded notable complexes like bis[tris(trimethylsilyl)methyl]calcium (\ce{[Ca{C(SiMe3)3}2]}), prepared in the 1990s through metathesis of calcium diiodide with the corresponding lithium alkyl in hydrocarbon solvents; the enormous steric bulk of the \ce{-C(SiMe3)3} ligands shields the reactive Ca-C bonds, enabling isolation as a two-coordinate, bent monomeric species stable enough for structural characterization and reactivity studies.

Industrial applications

Metallurgical uses

Calcium plays a critical role in production through desulfurization processes, where it is added to molten in amounts of 0.1–0.3% to react with , forming calcium (CaS) inclusions. These spherical CaS inclusions replace elongated inclusions, which can cause and reduce during ; the modification improves the 's and by preventing along inclusion boundaries. This treatment is typically performed in the ladle during secondary metallurgy, using calcium wire or cored wire injection to ensure and minimize losses due to calcium's high reactivity with oxygen. In lead alloys for maintenance-free batteries, calcium is incorporated at levels of 0.03–0.15% in combination with lead-antimony (Pb-Sb) to enhance performance. The addition reduces rates by forming stable passivation layers on the surfaces, minimizing and water loss, which extends battery life and eliminates the need for maintenance. These Pb-Ca-Sb alloys are into for automotive and lead-acid , where the calcium stabilizes the alloy against grid growth during charge-discharge cycles. Calcium also contributes to aluminum alloys through the CaAl₂, which acts as a grain refiner to promote finer equiaxed grains during solidification. This refinement enhances the alloy's mechanical strength, fatigue resistance, and castability by increasing sites and restricting growth, particularly in wrought aluminum alloys used for and automotive components. The addition is controlled to avoid excess calcium, which could lead to brittle phases. In rare earth metal production, calcium serves as a to eliminate fluoride impurities during . High-purity calcium is added to rare earth s (e.g., NdF₃ or DyF₃) in a sealed under , where it preferentially reduces the fluorides to volatile CaF₂, allowing of the pure rare earth metal at temperatures around 1000–1400°C. This calciothermic process is essential for achieving metallic purity above 99.5%, as residual fluorides degrade magnetic properties in applications like permanent magnets.

Chemical and construction uses

Calcium compounds play a pivotal role in construction and chemical industries, particularly through the use of (CaO) and its derivatives. In cement production, is calcined to produce CaO, which constitutes 60–67% of clinker and reacts with silicates, aluminates, and ferrites to form the binding phases essential for hydraulic . Global Portland cement output reached 4.1 gigatons in , underscoring its scale in infrastructure development. Lime, primarily as (Ca(OH)₂), is widely employed in for softening and adjusting . The process involves raising the water's to precipitate hardness-causing ions as (CaCO₃), effectively removing temporary hardness while stabilizing for downstream and control. This application reduces scaling in pipes and improves in municipal systems. In fertilizers, (CaCN₂) serves as a slow-release nitrogen source derived from processes, providing both nitrogen and calcium to enhance and suppress weeds or pathogens. Its use has declined relative to synthetic alternatives like due to handling requirements and slower nutrient release, though it remains valued in niche applications for . Calcium oxide is integral to the paper and sugar industries for purification and bleaching. In papermaking, quicklime (CaO) is used to regenerate caustic soda in the and produce for bleaching pulp, aiding removal and brightness enhancement. In sugar refining, clarifies raw juice by neutralizing acids, precipitating impurities, and forming calcium saccharate for filtration, typically at rates of 0.25 tons per ton of sugar. Calcium fluoride (CaF₂) finds application in toothpaste formulations as a mild abrasive, polishing tooth surfaces while releasing fluoride ions to support remineralization and caries prevention. This dual functionality makes it suitable for sensitive enamel formulations.

Biological role

Dietary sources and nutrition

Calcium is an essential mineral obtained primarily through dietary sources, with recommended daily intakes varying by age, sex, and physiological status to support bone health and other functions. According to the National Institutes of Health (NIH), adults aged 19 to 50 years require 1,000 mg of calcium per day, while women over 50 and men over 70 need 1,200 mg; for pregnant and lactating individuals, the recommended dietary allowance (RDA) is 1,300 mg for those aged 14 to 18 and 1,000 mg for ages 19 to 50. These guidelines align with broader international recommendations from organizations like the International Osteoporosis Foundation, emphasizing intakes around 1,000 mg for most adults to prevent deficiency. Rich dietary sources of calcium include dairy products, which provide highly bioavailable forms. For example, cow's milk contains approximately 120 mg per 100 g, while hard cheeses like cheddar offer about 700 mg per 100 g. Non-dairy options encompass leafy green vegetables such as , with around 150 mg per 100 g, and fortified foods like , which can deliver 300 mg per cup when enriched with calcium citrate malate or similar compounds. Other contributors include fish with edible bones (e.g., sardines) and certain nuts or seeds, though plant-based sources often require larger portions to meet daily needs due to varying absorption rates. The of dietary calcium, or the fraction absorbed in the intestines, typically ranges from 30% to 40% for sources like and fortified products. is enhanced by factors such as , which promotes intestinal uptake, and in , which aids solubility. Conversely, inhibitors like phytates (found in grains and ) and oxalates (present in and ) can bind calcium, reducing its absorption by up to 50% in high-fiber diets. For individuals unable to meet requirements through diet, such as (CaCO₃) or are commonly used, with effective doses of 500 to 600 mg taken in divided portions to maximize absorption. , which requires for dissolution, is inexpensive but best taken with meals, whereas citrate is more readily absorbed on an empty . The (WHO) includes on its List, recommending 1.5 to 2 g of elemental calcium daily for pregnant women in regions with low dietary intake to prevent complications like . Globally, inadequate calcium intake affects an estimated 3.5 billion people, particularly in low- and middle-income countries where consumption is limited and plant-based diets predominate. This deficiency heightens the risk of , a condition characterized by reduced and increased fracture susceptibility, underscoring the need for targeted nutritional interventions.

Physiological functions

Calcium plays a fundamental structural role in the , with over 99% of total body calcium stored in bones and teeth in the form of , \ce{Ca10(PO4)6(OH)2}, which imparts rigidity and strength to these tissues. In an average adult, this amounts to approximately 1–2 kg of calcium, primarily contributing to the mineralization and mechanical support of skeletal structures. crystals integrate with an organic matrix of and other proteins, forming a that withstands compressive forces and maintains overall body posture. As a key signaling molecule, the calcium ion (\ce{Ca^2+}) serves as a ubiquitous second messenger in cellular processes, particularly in neurons and muscle cells, where it orchestrates responses to stimuli. In neurons, \ce{Ca^2+} influx during action potentials triggers release at synapses by binding to proteins like synaptotagmin, while in muscle cells, it is released from (ER) stores via channels such as receptors, amplifying signals for contraction and other functions. These transient elevations in cytosolic \ce{Ca^2+} concentration, often from 100 nM to 1–10 μM, activate downstream pathways including calmodulin-dependent kinases, ensuring precise temporal and spatial control of cellular events. In muscle contraction, \ce{Ca^2+} binds specifically to , a regulatory subunit of the complex on thin filaments, inducing a conformational change that shifts away from myosin-binding sites on . This exposure allows myosin heads to interact with , forming cross-bridges powered by , which generates the sliding filament mechanism essential for force production in skeletal, cardiac, and smooth muscles. The process is tightly regulated, with \ce{Ca^2+} dissociation leading to relaxation as re-blocks the sites. Calcium ions are required for the activity of numerous enzymes across various metabolic pathways, either as direct cofactors or through calcium-binding regulatory proteins such as , facilitating their activation. In blood clotting, for instance, \ce{Ca^2+} is required for the activation of coagulation factors VII and IX, enabling the formation of the tenase that amplifies the proteolytic leading to clot formation. This cofactor role extends to other processes, such as and , where \ce{Ca^2+} modulates by altering or inducing allosteric changes. Finally, extracellular \ce{Ca^2+} contributes to cell stability by modulating permeability and maintaining the integrity of bilayers. It interacts with proteins and to regulate gating and prevent excessive leakage, thereby preserving the electrochemical gradients vital for cellular excitability and volume control. Intracellularly, controlled \ce{Ca^2+} levels further support function by influencing cytoskeletal attachments and vesicular trafficking.

Homeostasis and regulation

Calcium homeostasis in humans maintains calcium concentrations within a narrow range of 2.2–2.6 mmol/L (8.8–10.4 mg/dL) to support essential physiological processes. In this total calcium, approximately 50% circulates as ionized (free) calcium, which is biologically active; 40% is bound to proteins, primarily ; and 10% forms complexes with anions such as , citrate, and . This distribution ensures that only the ionized fraction directly influences cellular functions and signaling, while bound and complexed forms serve as reservoirs. The primary hormonal regulators of calcium are (PTH), calcitonin, and the active form of , 1,25-dihydroxyvitamin D₃ (). PTH, secreted by the parathyroid glands in response to low serum calcium, elevates calcium levels by stimulating osteoclast-mediated bone resorption to release calcium from crystals, enhancing renal reabsorption of calcium in the , and promoting the renal synthesis of . Calcitonin, produced by C-cells, opposes PTH by inhibiting activity and promoting renal calcium excretion, thereby lowering serum calcium, though its role in humans is less prominent under normal conditions. , synthesized in the kidneys from 25-hydroxyvitamin D under PTH stimulation, primarily increases intestinal calcium absorption efficiency from about 10–15% in deficiency states to 30–40% in sufficiency by upregulating the expression of calcium transport proteins in enterocytes. Cellular transport mechanisms are crucial for maintaining calcium balance across tissues. In the intestine, apical entry of calcium into enterocytes occurs primarily through TRPV6 channels, which are highly selective for calcium and upregulated by to facilitate transcellular absorption, particularly in the and . Once inside the cell, calcium is buffered by and extruded basolaterally via plasma membrane Ca²⁺-ATPase (PMCA1b) pumps or the Na⁺/Ca²⁺ exchanger (NCX1), ensuring vectorial transport into the bloodstream. Similar PMCA and NCX mechanisms operate in other tissues, such as the and cells, to fine-tune calcium flux and prevent intracellular overload. Bone serves as the main reservoir for calcium, with homeostasis achieved through continuous remodeling involving osteoclasts and osteoblasts. Osteoclasts dissolve in bone matrix via acidification and proteolytic enzymes, releasing calcium into the circulation to counteract ; this process is stimulated by PTH and (receptor activator of nuclear factor kappa-B ligand) produced by osteoblasts. Osteoblasts, in turn, deposit new to rebuild bone, maintaining structural integrity; they regulate osteoclastogenesis by balancing secretion with (OPG), a receptor that inhibits RANKL-RANK interactions on osteoclast precursors. This RANKL/OPG ratio ensures that and formation remain coupled, preventing net calcium loss under steady-state conditions. Feedback mechanisms tightly control these processes to stabilize serum calcium. Detection of low ionized calcium by calcium-sensing receptors on parathyroid chief cells triggers rapid PTH release, which acts within minutes to hours to restore levels through the aforementioned pathways. Conversely, rising calcium inhibits PTH secretion and stimulates calcitonin release, providing a counter-regulatory loop, while further amplifies intestinal uptake as a longer-term adjustment. These integrated systems—hormonal, transport, and remodeling—collectively maintain calcium despite varying dietary intake or physiological demands.

Pathological conditions

Hypocalcemia is defined as a serum calcium concentration below 2.2 mmol/L (8.8 mg/dL), resulting from impaired calcium mobilization from , reduced intestinal , or increased renal . This condition can lead to neuromuscular irritability, manifesting as , muscle cramps, , and in severe cases, seizures or . In children, chronic hypocalcemia often contributes to , characterized by softened bones and skeletal deformities due to impaired mineralization, while in adults, it may cause , leading to and increased fracture risk. Hypercalcemia occurs when serum calcium exceeds 2.6 mmol/L (10.4 mg/dL), most commonly due to or malignancy-associated factors such as bone metastases or humoral hypercalcemia of malignancy. Symptoms arise from , renal effects, and involvement, including , , , kidney stones, and ; severe cases can progress to cardiac arrhythmias or . Osteoporosis is a systemic skeletal disorder defined by low bone mass and microarchitectural deterioration, diagnosed when bone mineral density (BMD) is more than 2.5 standard deviations below the young adult mean (T-score ≤ -2.5), often linked to chronic calcium dysregulation and deficiency in postmenopausal women. This condition affects over 200 million people worldwide, with postmenopausal women at highest risk due to accelerated outpacing formation, increasing susceptibility to fragility fractures at the , , and . Primary hyperparathyroidism, with a prevalence of approximately 1 in 1000 adults, involves autonomous overproduction of (PTH) by one or more parathyroid glands, typically due to , leading to excessive , hypercalcemia, and increased fracture risk. Elevated PTH promotes activity, releasing calcium from bone into the bloodstream, which can exacerbate renal stone formation and cardiovascular complications if untreated. Calcium channelopathies, genetic disorders affecting voltage-gated calcium channels, represent an emerging area of research in the 2020s, with conditions like Timothy syndrome—caused by mutations in the CACNA1C gene—resulting in gain-of-function alterations that disrupt calcium influx, contributing to neurodevelopmental issues such as autism spectrum disorder, seizures, and cardiac arrhythmias. Recent studies highlight the multisystem impact of these channel defects, emphasizing targeted therapies to modulate channel activity.

Safety

Handling elemental calcium

Elemental calcium is a highly reactive that poses significant hazards during handling due to its flammability and sensitivity to environmental factors. Finely divided calcium can spontaneously ignite in moist air, forming a dark oxide-nitride layer and potentially leading to fire if exposed to humidity levels above trace amounts. It reacts violently with water, liberating gas according to the Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g), which can result in explosions if the gas accumulates in confined spaces or ignites. To prevent such reactions, calcium must be stored and manipulated under an inert atmosphere, such as , or submerged in to exclude oxygen and moisture. Proper storage is essential for maintaining the integrity of elemental calcium, which should be kept in tightly sealed containers made of compatible materials like or certain metals, in a cool, dry, and well-ventilated area isolated from acids, oxidizers, , and any sources of ignition. Grounding and bonding of containers help prevent static discharge that could spark a . When stored under these conditions, elemental calcium maintains its usability for up to 36 months, though regular for surface tarnishing or degradation is recommended. Individuals handling elemental calcium require appropriate (PPE) to minimize risks of burns, irritation, or inhalation hazards. This includes or gloves, chemical-resistant coveralls, safety goggles or a full , and respiratory protection such as an or P3 filter respirator in dusty environments; is advised for high-exposure scenarios exceeding 30 mg/m³. Fire safety measures are critical, with Class D extinguishers containing chemicals, , or recommended for metal fires—water, , or halogenated agents must be avoided as they intensify the reaction by promoting evolution or . Industrial incidents involving calcium metal are uncommon owing to stringent handling protocols, but when they occur, they often stem from inadvertent exposure leading to rapid or formation, which causes severe burns to or eyes. Regulatory guidelines from the (OSHA) classify calcium dust under particulates not otherwise regulated, mandating an exposure limit of 5 mg/m³ as an 8-hour time-weighted average for the respirable fraction to protect against risks.

Health effects of imbalances

Ingestion of elemental calcium, a highly reactive metal, can lead to the formation of (Ca(OH)₂) upon contact with moisture, resulting in gastrointestinal irritation, , , and due to its alkaline and corrosive nature. Excessive intake or exposure to specifically contributes to , characterized by elevated blood pH, often alongside hypercalcemia and renal impairment in severe cases like . Inhalation of (CaO) or dust, common in industrial settings, irritates the and, with prolonged exposure, can cause , a fibrotic disease marked by scarring and reduced function. Overdose from exceeding the tolerable upper intake level of 2,500 mg/day for adults aged 19–50 (or 2,000 mg/day for those over 50) increases the risk of hypercalcemia, where serum calcium levels rise above 10.5 mg/dL, potentially leading to symptoms like , , and cardiac . Some 2024–2025 studies and reviews have suggested an association between and increased risks of cardiovascular events or in older adults, though evidence remains mixed and dietary calcium is generally safer. Common side effects include and , while chronic excess promotes renal calculi formation by elevating urinary calcium excretion (). In environmental contexts, drinking water hardness from calcium concentrations above 100 mg/L has been associated with reduced (CVD) risk in multiple studies and meta-analyses, potentially due to magnesium co-occurrence or anti-atherogenic effects, though remains unestablished and requires further randomized trials. Post-2010 research on CaO is limited, with most studies focusing on its or stress-alleviating properties rather than human health risks like or from or dermal . As of 2025, updates from health authorities confirm the safety of calcium-fortified foods, with rates around 30% and no widespread adverse effects when consumed within recommended dietary guidelines. Chronic exposure to excess (CaF₂), a fluoride source in some minerals or industrial dusts, is rare but can contribute to through elevated levels during tooth development, manifesting as or discoloration. This condition arises from disrupting enamel mineralization, though calcium itself may mitigate severity by binding ions.