Does calcium have a reaction to water?
Calcium carbonate
Calcium carbonate is a chemical compound with the formula CaCO. It is a common substance found in rocks as the minerals calcite and aragonite (most notably as limestone, which is a type of sedimentary rock consisting mainly of calcite) and is the main component of pearls and the shells of marine organisms, snails, and eggs. Calcium carbonate is the active ingredient in agricultural lime and is created when calcium ions in hard water react with carbonate ions to create limescale. It is medicinally used as a calcium supplement or as an antacid, but excessive consumption can be hazardous.
Calcium bicarbonate
In the reverse process, dissolved carbon dioxide (CO) in rainwater (HO) reacts with limestone calcium carbonate (CaCO) to form soluble calcium bicarbonate (Ca(HCO)). This soluble compound is then washed away with the rainwater. This form of weathering is called carbonation.
Amorphous calcium phosphate
Amorphous calcium phosphate (ACP or ATCP) is a glassy precipitate of variable composition that is formed in double decomposition reactions involving a soluble phosphate and calcium salts (e.g. (NH)HPO + Ca(NO)) performed under carefully controlled pH conditions. The precipitate will either be "amorphous tricalcium phosphate", ATCP, or calcium deficient hydroxypatite, CDHA, Ca(HPO)(PO)(OH), (note CDHA is sometimes termed apatitic calcium triphosphate). The composition of amorphous calcium phosphate is CaH(PO)·HO where n is between 3 and 4.5. Precipitation from the moderately supersaturated and basic solution containing magnesium, led the precipitation of amorphous magnesium calcium phosphate (AMCP) in which magnesium incorporated into the ACP structure.
Calcium
Like other elements, a variety of processes alter the relative abundance of calcium isotopes. The best studied of these processes is the mass-dependent fractionation of calcium isotopes that accompanies the precipitation of calcium minerals such as calcite, aragonite and apatite from solution. Lighter isotopes are preferentially incorporated into these minerals, leaving the surrounding solution enriched in heavier isotopes at a magnitude of roughly 0.025% per atomic mass unit (amu) at room temperature. Mass-dependent differences in calcium isotope composition are conventionally expressed by the ratio of two isotopes (usually Ca/Ca) in a sample compared to the same ratio in a standard reference material. Ca/Ca varies by about 1% among common earth materials.
Reef aquarium
Stony corals, which are defined by their calcerous calcium carbonate skeletons (CaCO), are the focus of many advanced reef keepers. These corals require additional attention to water chemistry, especially maintenance of stable and optimal calcium, carbonate, and pH levels. These parameters may be tracked and adjusted with test kits and frequent manual dosing of calcium and pH buffer additives requiring no additional equipment. Alternatively, automated methods employing small dedicated computers with electronic water quality monitoring capabilities are often used to control water chemistry parameters via several components including calcium reactors and kalkwasser reactors. Calcium reactors are canisters filled with crushed coral skeletons. Carbon dioxide is injected into the canister acidifying the water and dissolving the coral skeletons. The acidified and CaCO rich solution is then pumped into the sump. The excess CO then diffuses out of the water and into the air leaving behind the CaCO. Kalkwasser is an aqueous solution of calcium hydroxide, Ca(OH). The kalk reactor stirs the solution, preventing precipitation, and dispenses the solution into the sump where the Ca(OH) combines with dissolved CO to produce CaCO. These components must be controlled by a computer to prevent dangerous changes in pH due to the acidic calcium reactor effluent or alkaline kalkwasser effluent.
Calcium reactor
In marine and reef aquariums, a calcium reactor creates a balance of alkalinity. An acidic solution is produced by injecting carbon dioxide into a chamber with salt water and calcium rich media. The carbon dioxide lowers the pH by producing a solution high in carbonic acid, and dissolves calcium. The effluent is returned to the reef aquarium where the calcium is consumed by organisms, primarily corals when building skeletons. A calcium reactor is an efficient method to supply calcium to a reef aquarium. Reactors may be used in elaborate freshwater and brackish aquariums where freshwater clams and other invertebrates need a constant supply of calcium.
Calcium hydride
The reaction of CaH with water can be represented as follows:
Calcium
The chemistry of calcium is that of a typical heavy alkaline earth metal. For example, calcium spontaneously reacts with water more quickly than magnesium and less quickly than strontium to produce calcium hydroxide and hydrogen gas. It also reacts with the oxygen and nitrogen in the air to form a mixture of calcium oxide and calcium nitride. When finely divided, it spontaneously burns in air to produce the nitride. In bulk, calcium is less reactive: it quickly forms a hydration coating in moist air, but below 30% relative humidity it may be stored indefinitely at room temperature.
Potassium
Potassium is an extremely active metal that reacts violently with oxygen in water and air. With oxygen it forms potassium peroxide, and with water potassium forms potassium hydroxide. The reaction of potassium with water is dangerous because of its violent exothermic character and the production of hydrogen gas. Hydrogen reacts again with atmospheric oxygen, producing water, which reacts with the remaining potassium. This reaction requires only traces of water; because of this, potassium and the liquid sodium-potassium (NaK) alloy are potent desiccants that can be used to dry solvents prior to distillation.
Thermometric titration
An excellent application is the sequential determination of calcium and magnesium. Although calcium reacts exothermically with EDTA (heat of chelation ~-23.4 kJ/mol), magnesium reacts endothermically with a heat of chelation of ~+20.1 kJ/mol. This is illustrated in the titration plot of EDTA with calcium and magnesium in sea water (Figure 18). Following the solution temperature curve, the breakpoint for the calcium content (red-tagged endpoint) is followed by a region of modest temperature rise due to competition between the heats of dilution of the titrant with the solution, and the endothermic reaction of Mg and EDTA. The breakpoint for the consumption of Mg (blue-tagged endpoint) by EDTA is revealed by upswing in temperature caused purely by the heat of dilution.
Soda lime
A quite didactic parallelism can be drawn between the catalytic role of sodium hydroxide in the process of carbonation of soda lime and the role of sodium hydroxide played in the alkali-silica reaction, a slow degration process causing the swelling and the cracking of concrete containing aggregates rich in reactive amorphous silica. In a very similar way, NaOH greatly facilitates the dissolution of the amorphous silica. The produced sodium silicate then reacts with the calcium hydroxide (portlandite) present in the hardened cement paste to form calcium silicate hydrate (abbreviated as C-S-H in the cement chemist notation). This silicification reaction of Ca(OH) on its turn continuously releases again sodium hydroxide in solution, maintaining a high pH, and the cycle continues up to the total disappearance of portlandite or reactive silica in the exposed concrete. Without the catalysis of this reaction by sodium or potassium soluble hydroxides the alkali-silica reaction would not proceed or would be limited to a very slow pozzolanic reaction. The alkali silica reaction can be written like the soda lime reaction, by simply substituting CO by SiO in the reactions mentioned here above as follows:
Calcite rafts
Calcium carbonate is known to precipitate as calcite crystals in water supersaturated with calcium and carbonate ions. Under quiescent conditions, calcite crystals can form on a water surface when calcium carbonate supersaturation conditions do not exist in the bulk water. Water evaporates from the surface and carbon dioxide degasses from the surface layer to create a thin layer of water with high pH and concentrations of calcium and carbonate ions far above the saturation concentration for calcium carbonate. Calcite crystals precipitate in this highly localized environment and attach to one another to form what appear to be rafts of a white material.
Water-reactive substances
Magnesium shows insignificant reaction with water, but burns vigorously with steam or water vapor to produce white magnesium oxide and hydrogen gas:
Mg + 2HO ⟶ Mg(OH) + H
Calcium hydroxide
When heated to 512 °C, the partial pressure of water in equilibrium with calcium hydroxide reaches 101kPa (normal atmospheric pressure), which decomposes calcium hydroxide into calcium oxide and water.
Tricalcium aluminate
In keeping with its high basicity, tricalcium aluminate reacts most strongly with water of all the calcium aluminates, and it is also the most reactive of the Portland clinker phases. Its hydration to phases of the form CaAlO(OH)·nHO leads to the phenomenon of "flash set" (instantaneous set), and a large amount of heat is generated. To avoid this, Portland-type cements include a small addition of calcium sulfate (typically 4-8%). Sulfate ions in solution lead to the formation of an insoluble layer of ettringite (3CaO·AlO·3CaSO·32HO) over the surface of the aluminate crystals, passivating them. The aluminate then reacts slowly to form the AFm phase 3CaO·AlO·CaSO·12HO. These hydrates contribute little to strength development.
Nephrogenic diabetes insipidus
High blood calcium causes natriuresis (increased sodium loss in the urine) and water diuresis, in part by its effect through the calcium-sensing receptor (CaSR).
Fouling
As an example, the equilibrium between the readily soluble calcium bicarbonate - always prevailing in natural water - and the poorly soluble calcium carbonate, the following chemical equation may be written:
Calcium bicarbonate
The above reaction is very important to the formation of stalactites, stalagmites, columns, and other speleothems within caves, and for that matter, in the formation of the caves themselves. As water containing carbon dioxide (including extra CO acquired from soil organisms) passes through limestone or other calcium carbonate-containing minerals, it dissolves part of the calcium carbonate, hence becomes richer in bicarbonate. As the groundwater enters the cave, the excess carbon dioxide is released from the solution of the bicarbonate, causing the much less soluble calcium carbonate to be deposited.
Coupling reaction
While many coupling reactions involve reagents that are extremely susceptible to presence of water or oxygen, it is unreasonable to assume that all coupling reactions need to be performed with strict exclusion of water. It is possible to perform palladium-based coupling reactions in aqueous solutions using the water-soluble sulfonated phosphines made by the reaction of triphenyl phosphine with sulfuric acid. Another example of coupling in aqueous media, with the main reacting agent being trimolybdenum-alkylidyne clusters, is that of Bogoslavsky et al. In general, the oxygen in the air is more able to disrupt coupling reactions, because many of these reactions occur via unsaturated metal complexes that do not have 18 valence electrons. For example, in nickel and palladium cross couplings, a zerovalent complex with two vacant sites (or labile ligands) reacts with the carbon halogen bond to form a metal halogen and a metal carbon bond. Such a zerovalent complex with labile ligands or empty coordination sites is normally very reactive toward oxygen.
Calcium
The actual reaction is more complicated and involves the bicarbonate ion (HCO) that forms when CO reacts with water at seawater pH:
At seawater pH, most of the CO is immediately converted back into . The reaction results in a net transport of one molecule of CO from the ocean/atmosphere into the lithosphere. The result is that each Ca ion released by chemical weathering ultimately removes one CO molecule from the surficial system (atmosphere, ocean, soils and living organisms), storing it in carbonate rocks where it is likely to stay for hundreds of millions of years. The weathering of calcium from rocks thus scrubs CO from the ocean and atmosphere, exerting a strong long-term effect on climate.
Calcium carbonate
Alternatively, calcium carbonate is prepared from calcium oxide. Water is added to give calcium hydroxide then carbon dioxide is passed through this solution to precipitate the desired calcium carbonate, referred to in the industry as precipitated calcium carbonate (PCC):
Calcium
Besides the simple oxide CaO, the peroxide CaO can be made by direct oxidation of calcium metal under a high pressure of oxygen, and there is some evidence for a yellow superoxide Ca(O). Calcium hydroxide, Ca(OH), is a strong base, though it is not as strong as the hydroxides of strontium, barium or the alkali metals. All four dihalides of calcium are known. Calcium carbonate (CaCO) and calcium sulfate (CaSO) are particularly abundant minerals. Like strontium and barium, as well as the alkali metals and the divalent lanthanides europium and ytterbium, calcium metal dissolves directly in liquid ammonia to give a dark blue solution.
Signal transduction
The release of calcium ions from the endoplasmic reticulum into the cytosol results in its binding to signaling proteins that are then activated; it is then sequestered in the smooth endoplasmic reticulum and the mitochondria. Two combined receptor/ion channel proteins control the transport of calcium: the InsP-receptor that transports calcium upon interaction with inositol triphosphate on its cytosolic side; and the ryanodine receptor named after the alkaloid ryanodine, similar to the InsP receptor but having a feedback mechanism that releases more calcium upon binding with it. The nature of calcium in the cytosol means that it is active for only a very short time, meaning its free state concentration is very low and is mostly bound to organelle molecules like calreticulin when inactive.
Calcium carbide
The reaction of calcium carbide with water, producing acetylene and calcium hydroxide, was discovered by Friedrich Wöhler in 1862.
Mitochondrial calcium uniporter
The MCU has a very low affinity for calcium, so the cytosolic calcium concentration needs to be approximately 5-10 uM for significant transport of calcium into the mitochondria. Mitochondria are closely associated with the endoplasmic reticulum (ER), at contact sites, which contains stores of cellular calcium ions for calcium signaling. The presence of 1,4,5-triphosphate (IP) triggers the release of calcium from these intracellular stores, which creates microdomains of high calcium concentration between the ER and the mitochondria, creating the conditions for the MCU to take up calcium.
Common-ion effect
A practical example used very widely in areas drawing drinking water from chalk or limestone aquifers is the addition of sodium carbonate to the raw water to reduce the hardness of the water. In the water treatment process, highly soluble sodium carbonate salt is added to precipitate out sparingly soluble calcium carbonate. The very pure and finely divided precipitate of calcium carbonate that is generated is a valuable by-product used in the manufacture of toothpaste.
Edible water bottle
Sodium alginate (NaAlg) coagulates when exposed to calcium chloride (CaCl) and forms calcium alginate (CaAlg ) and sodium chloride (NaCl), according to the following reaction:
Clarifying agent
Many flocculants are multivalent cations such as aluminium, iron, calcium or magnesium. These positively charged molecules interact with negatively charged particles and molecules to reduce the barriers to aggregation. In addition, many of these chemicals, under appropriate pH and other conditions such as temperature and salinity, react with water to form insoluble hydroxides which, upon precipitating, link together to form long chains or meshes, physically trapping small particles into the larger floc.
Calcium cycle
Calcium is an essential component of soil. When deposited in the form of lime, it can't be used by plants. To combat this, carbon dioxide produced by plants reacts with water in the environment to produce carbonic acid. Carbonic acid is then able to dissolve limestone, enabling the release of calcium ions. This reaction is more readily available with smaller particles of limestone than it is with large pieces of rock due to the increased surface area. When lime is leached into soil, calcium levels inevitably increase, both stabilising pH and enabling calcium to mix with water to form a Ca 2+ ions, thus making it soluble and accessible to plants to be absorbed and utilised by the root system. The calcium ions travel up the xylem of the plant alongside water to reach the leaves. The plant can utilise this calcium in the form of calcium pectate to stabilise cell walls and provide rigidity. Calcium is also used by plant enzymes to signal growth and coordinate life-promoting processes . Additionally, the release of calcium ions enables microorganisms to access phosphorus and other micro nutrients with greater ease, improving the soil ecosystem drastically thus indirectly promoting plant growth and nutrition.
Calcium bicarbonate
All waters in contact with the atmosphere absorb carbon dioxide, and as these waters come into contact with rocks and sediments they acquire metal ions, most commonly calcium and magnesium, so most natural waters that come from streams, lakes, and especially wells, can be regarded as dilute solutions of these bicarbonates. These hard waters tend to form carbonate scale in pipes and boilers and they react with soaps to form an undesirable scum.
Calcium carbonate
make it possible to solve simultaneously for the remaining five unknown concentrations (note that the above form of the neutrality equation is valid only if calcium carbonate has been put in contact with pure water or with a neutral pH solution; in the case where the initial water solvent pH is not neutral, the equation is modified).
Pozzolan
Pozzolans are a broad class of siliceous or siliceous and aluminous materials which, in themselves, possess little or no cementitious value but which will, in finely divided form and in the presence of water, react chemically with calcium hydroxide at ordinary temperature to form compounds possessing cementitious properties. The quantification of the capacity of a pozzolan to react with calcium hydroxide and water is given by measuring its pozzolanic activity. "Pozzolana" are naturally occurring pozzolans of volcanic origin.
Residual sodium carbonate index
The following chemical reactions take place in lime soda softening process which precipitates the calcium and magnesium salts as calcium carbonate and magnesium hydroxide which have very low solubility in water.The excess soda ash after precipitating the calcium and magnesium salts is in carbonates & bicarbonates of sodium which imparts high pH or alkalinity to soil water.
Solubility
The chart shows solubility curves for some typical solid inorganic salts (temperature is in degrees Celsius i.e. kelvins minus 273). Many salts behave like barium nitrate and disodium hydrogen arsenate, and show a large increase in solubility with temperature. Some solutes (e.g., sodium chloride in water) exhibit solubility that is fairly independent of temperature. A few, such as calcium sulfate (gypsum) and cerium(III) sulfate, become less soluble in water as temperature increases. This is also the case for calcium hydroxide (portlandite), whose solubility at 70 °C is about half of its value at 25 °C. The dissolution of calcium hydroxide in water is also an exothermic process and obeys Le Chatelier's principle. A lowering of temperature will thus favor the dissolution heat elimination and increases the equilibrium constant of dissolution of Ca(OH), and so increase its solubility at low temperature. This temperature dependence is sometimes referred to as "retrograde" or "inverse" solubility. Occasionally, a more complex pattern is observed, as with sodium sulfate, where the less soluble decahydrate crystal loses water of crystallization at 32 °C to form a more soluble anhydrous phase.
The solubility of organic compounds nearly always increases with temperature. The technique of recrystallization, used for purification of solids, depends on a solute's different solubilities in hot and cold solvent. A few exceptions exist, such as certain cyclodextrins.
Calcium oxide
Calcium oxide (CaO), commonly known as quicklime or burnt lime, is a widely used chemical compound. It is a white, caustic, alkaline, crystalline solid at room temperature. The broadly used term "lime" connotes calcium-containing inorganic materials, in which carbonates, oxides and hydroxides of calcium, silicon, magnesium, aluminium, and iron predominate. By contrast, "quicklime" specifically applies to the single chemical compound calcium oxide. Calcium oxide that survives processing without reacting in building products such as cement is called free lime.
Hydrofluoric acid
Once absorbed into blood through the skin, it reacts with blood calcium and may cause cardiac arrest. Burns with areas larger than 160 cm (25 square inches) have the potential to cause serious systemic toxicity from interference with blood and tissue calcium levels. In the body, hydrofluoric acid reacts with the ubiquitous biologically important ions Ca and Mg. Formation of insoluble calcium fluoride is proposed as the etiology for both precipitous fall in serum calcium and the severe pain associated with tissue toxicity. In some cases, exposures can lead to hypocalcemia. Thus, hydrofluoric acid exposure is often treated with calcium gluconate, a source of Ca that sequesters the fluoride ions. HF chemical burns can be treated with a water wash and 2.5% calcium gluconate gel or special rinsing solutions. However, because it is absorbed, medical treatment is necessary; rinsing off is not enough. Intra-arterial infusions of calcium chloride have also shown great effectiveness in treating burns.
Calcium
Although calcium is in the same group as magnesium and organomagnesium compounds are very commonly used throughout chemistry, organocalcium compounds are not similarly widespread because they are more difficult to make and more reactive, although they have recently been investigated as possible catalysts. Organocalcium compounds tend to be more similar to organoytterbium compounds due to the similar ionic radii of Yb (102 pm) and Ca (100 pm). Most of these compounds can only be prepared at low temperatures; bulky ligands tend to favor stability. For example, calcium dicyclopentadienyl, Ca(CH), must be made by directly reacting calcium metal with mercurocene or cyclopentadiene itself; replacing the CH ligand with the bulkier C(CH) ligand on the other hand increases the compound's solubility, volatility, and kinetic stability.
Milk fever
Intravenous calcium, though indicated in many cases, is potentially fatal through "heart blockade", or transient high calcium levels stopping the heart, so should be administered with care.
Cows are to be fed jaggery along with the lime water mixture.
In unclear cases of downer cows, intravenous calcium injection can lead to diagnosis. The typical reaction will be a generalized tremor of the skeletal muscles, and sometimes cardiac arrhythmia. Defecation, urination and eructation are frequent during the treatment, due to pharmacological effect of calcium on the smooth muscles.
Salt (chemistry)
Many ionic compounds exhibit significant solubility in water or other polar solvents. Unlike molecular compounds, salts dissociate in solution into anionic and cationic components. 
The lattice energy, the cohesive forces between these ions within a solid, determines the solubility. The solubility is dependent on how well each ion interacts with the solvent, so certain patterns become apparent. For example, salts of sodium, potassium and ammonium are usually soluble in water. Notable exceptions include ammonium hexachloroplatinate and potassium cobaltinitrite. Most nitrates and many sulfates are water-soluble. Exceptions include barium sulfate, calcium sulfate (sparingly soluble), and lead(II) sulfate, where the 2+/2− pairing leads to high lattice energies. For similar reasons, most alkali metal carbonates are not soluble in water. Some soluble carbonate salts are: sodium carbonate, potassium carbonate and ammonium carbonate.
Calcium hydroxide
At ambient temperature. Calcium hydroxide (portlandite) dissolves in pure water to produce an alkaline solution with a pH of about 12.4. Calcium hydroxide solutions can cause chemical burns. At high pH value (see common ion effect), its solubility drastically decreases. This behavior is relevant to cement pastes. Aqueous solutions of calcium hydroxide are called limewater and are medium strength bases that reacts with acids and can attack some metals such as aluminium (amphoteric hydroxide dissolving at high pH) while protecting other metals from corrosion such as iron and steel by passivation of their surface. Limewater turns milky in the presence of carbon dioxide due to formation of calcium carbonate, a process called carbonatation:for example lime water
Calcium reactor
Inside the reaction chamber, a calcium rich media (aragonite), mainly CaCO, is forced into contact with water injected with carbon dioxide (CO) in order to create carbonic acid (HCO). This increases the solubility of the calcium carbonate. The reaction frees the calcium and carbonate, supplying the aquarium with water rich in Ca and CO, important for maintaining alkalinity and calcium levels.
Salt water chlorination
The chlorinator cell consists of parallel titanium plates coated with ruthenium and sometimes iridium. Older models make use of perforated (or mesh) plates rather than solid plates. Electrolysis naturally attracts calcium and other minerals to the plates. Thus, depending on water chemistry and magnitude of use, the cell will require periodic cleaning in a mild acid solution (1 part HCl to 15 parts pool water) which will remove the buildup of calcium compound crystals, such as calcium carbonate or calcium nitrate. Excessive buildup can reduce the effectiveness of the cell. Running the chlorinator for long periods with not enough salt in the pool can strip the coating off the cell which then requires an expensive replacement, as can using too strong an acid wash.
Calcium channel blocker
Side effects of these drugs may include but are not limited to:Research indicates ethanol is involved in the inhibition of L-type calcium channels. One study showed the nature of ethanol binding to L-type calcium channels is according to first-order kinetics with a Hill coefficient around 1. This indicates ethanol binds independently to the channel, expressing noncooperative binding. Early studies showed a link between calcium and the release of vasopressin by the secondary messenger system. Vasopressin levels are reduced after the ingestion of alcohol. The lower levels of vasopressin from the consumption of alcohol have been linked to ethanol acting as an antagonist to voltage-gated calcium channels (VGCCs). Studies conducted by Treistman et al. in the aplysia confirm inhibition of VGCC by ethanol. Voltage clamp recordings have been done on the aplysia neuron. VGCCs were isolated and calcium current was recorded using patch clamp technique having ethanol as a treatment. Recordings were replicated at varying concentrations (0, 10, 25, 50, and 100 mM) at a voltage clamp of +30 mV. Results showed calcium current decreased as concentration of ethanol increased. Similar results have shown to be true in single-channel recordings from isolated nerve terminal of rats that ethanol does in fact block VGCCs.
Lime softening
As lime in the form of limewater is added to raw water, the pH is raised and the equilibrium of carbonate species in the water is shifted. Dissolved carbon dioxide (CO) is changed into bicarbonate (HCO) and then carbonate (CO). This action causes calcium carbonate to precipitate due to exceeding the solubility product. Additionally, magnesium can be precipitated as magnesium hydroxide in a double displacement reaction.
Calcium stearoyl-2-lactylate
As described by the Food Chemicals Codex 7th edition, CSL is a cream-colored powder. CSL is currently manufactured by the esterification of stearic acid and lactic acid with partial neutralization using food-grade hydrated lime (calcium hydroxide). Commercial grade CSL is a mixture of calcium salts of stearoyl lactic acid, with minor proportions of other salts of related acids. The HLB for CSL is 5.1. It is slightly soluble in hot water. The pH of a 2% aqueous suspension is approximately 4.7.
Carbon dioxide
Acids liberate from most metal carbonates. Consequently, it may be obtained directly from natural carbon dioxide springs, where it is produced by the action of acidified water on limestone or dolomite. The reaction between hydrochloric acid and calcium carbonate (limestone or chalk) is shown below:
The carbonic acid () then decomposes to water and :
Such reactions are accompanied by foaming or bubbling, or both, as the gas is released. They have widespread uses in industry because they can be used to neutralize waste acid streams.
Hard water
In swimming pools, hard water is manifested by a turbid, or cloudy (milky), appearance to the water. Calcium and magnesium hydroxides are both soluble in water. The solubility of the hydroxides of the alkaline-earth metals to which calcium and magnesium belong (group 2 of the periodic table) increases moving down the column. Aqueous solutions of these metal hydroxides absorb carbon dioxide from the air, forming the insoluble carbonates, giving rise to the turbidity. This often results from the pH being excessively high (pH > 7.6). Hence, a common solution to the problem is, while maintaining the chlorine concentration at the proper level, to lower the pH by the addition of hydrochloric acid, the optimum value being in the range of 7.2 to 7.6.
Blackwater river
Black and white waters differ significantly in their ionic composition, as shown in Table 1. Black waters are more acidic, resulting in an aluminum concentration greater than that of the more neutral white waters. The major difference is the concentrations of sodium, magnesium, calcium and potassium; these are very low in black waters. This has ecological implications. Some animals need more calcium than is available in blackwaters, so for example, snails, which need much calcium to build shells, are not abundant in blackwaters. The lack of dissolved ions in black waters results in a low conductivity, similar to that of rainwater.
Calcium chloride
Calcium chloride dissolves in water, producing chloride and the aquo complex [Ca(HO)]. In this way, these solutions are sources of "free" calcium and free chloride ions. This description is illustrated by the fact that these solutions react with phosphate sources to give a solid precipitate of calcium phosphate:
Calcium chloride has a very high enthalpy change of solution, indicated by considerable temperature rise accompanying dissolution of the anhydrous salt in water. This property is the basis for its largest-scale application.
Calcium phosphate
Calcium phosphates are found in many living organisms, e.g., bone mineral and tooth enamel. In milk, it exists in a colloidal form in micelles bound to casein protein with magnesium, zinc, and citrate - collectively referred to as colloidal calcium phosphate (CCP). Various calcium phosphate minerals are used in the production of phosphoric acid and fertilizers. Overuse of certain forms of calcium phosphate can lead to nutrient-containing surface runoff and subsequent adverse effects upon receiving waters such as algal blooms and eutrophication.
Calcium sulfate
The dissolution of the different crystalline phases of calcium sulfate in water is exothermic and release heat (decrease in Enthalpy: ΔH < 0). As an immediate consequence, to proceed, the dissolution reaction needs to evacuate this heat that can be considered as a product of reaction. If the system is cooled, the dissolution equilibrium will evolve towards the right according to the Le Chatelier principle and calcium sulfate will dissolve more easily. The solubility of calcium sulfate increases thus when the temperature decreases. If the temperature of the system is raised, the reaction heat cannot dissipate and the equilibrium will regress towards the left according to Le Chatelier principle. The solubility of calcium sulfate decreases thus when temperature increases. This contra-intuitive solubility behaviour is called retrograde solubility. It is less common than for most of the salts whose dissolution reaction is endothermic (i.e., the reaction consumes heat: increase in Enthalpy: ΔH > 0) and whose solubility increases with temperature. Another calcium compound, calcium hydroxide (Ca(OH), portlandite) also exhibits a retrograde solubility for the same thermodynamic reason: because its dissolution reaction is also exothermic and releases heat. So, to dissolve higher amounts of calcium sulfate or calcium hydroxide in water, it is necessary to cool down the solution close to its freezing point instead of increasing its temperature.
Tadelakt
The soap chemically reacts with the lime plaster, forming lime (calcium) soaps. Calcium soaps are insoluble in water, and fairly hard. They are familiar, in areas with calcium-rich ("hard") water, as deposits in bathtubs, sinks, and showers; when soap is mixed with the water's dissolved calcium carbonate/lime, calcium soaps form.
Hypocalcaemia
Physiologically, blood calcium is tightly regulated within a narrow range for proper cellular processes. Calcium in the blood exists in three primary states: bound to proteins (mainly albumin), bound to anions such as phosphate and citrate, and as free (unbound) ionized calcium. Only the ionized calcium is physiologically active. Normal blood calcium level is between 8.5 to 10.5 mg/dL (2.12 to 2.62 mmol/L) and that of ionized calcium is 4.65 to 5.25 mg/dL (1.16 to 1.31 mmol/L).
Calthemite
Calcium hydroxide will readily react with any free CO to form calcium carbonate (CaCO) [Equation 2]. The solution is typically pH 9 - 10.3, however this will depend on what other chemical reactions are also occurring at the same time within the concrete.
Flocculation
Yeast flocculation is primarily determined by the calcium concentration, often in the 50-100ppm range. Calcium salts can be added to cause flocculation, or the process can be reversed by removing calcium by adding phosphate to form insolubable calcium phosphate, adding excess sulfate to form insoluble calcium sulfate, or adding EDTA to chelate the calcium ions. While it appears similar to sedimentation in colloidal dispersions, the mechanisms are different.
Calcium hypochlorite
Similarly, calcium hypochlorite reacts with hydrochloric acid to form calcium chloride, water and chlorine:
Calcium acetate
Calcium acetate is a chemical compound which is a calcium salt of acetic acid. It has the formula Ca(CHO). Its standard name is calcium acetate, while calcium ethanoate is the systematic name. An older name is acetate of lime. The anhydrous form is very hygroscopic; therefore the monohydrate (Ca(CHCOO)•HO) is the common form.
Caesium
Caesium metal is highly reactive and very pyrophoric. It ignites spontaneously in air, and reacts explosively with water even at low temperatures, more so than the other alkali metals (first group of the periodic table). It reacts with solid water at temperatures as low as . Because of this high reactivity, caesium metal is classified as a hazardous material. It is stored and shipped in dry, saturated hydrocarbons such as mineral oil. It can be handled only under inert gas, such as argon. However, a caesium-water explosion is often less powerful than a sodium-water explosion with a similar amount of sodium. This is because caesium explodes instantly upon contact with water, leaving little time for hydrogen to accumulate. Caesium can be stored in vacuum-sealed borosilicate glass ampoules. In quantities of more than about 100 grams (3.5 oz), caesium is shipped in hermetically sealed, stainless steel containers.
Calcium cycle
The concentration of marine calcium depends on the depth of the ocean, as its solubility is altered with pressure. In deep ocean, roughly 1% more calcium is present when compared to surface concentrations. This is due to the solubility of calcium at higher pressure, thus associating higher concentration with increasing depth. The higher pressure stops calcium carbonate precipitation, whilst simultaneously enabling the dissolution of calcium carbonate sediment such as shells. As a result, the sedimentation is only found in shallower oceans, with the ocean floor at greater depths and pressure not being conducive to limestone formation.
Downs cell
The calcium does not enter into the reaction because its reduction potential of -2.87 volts is lower than that of sodium, which is -2.71 volts. Hence the sodium ions are reduced to metallic form in preference to those of calcium. If the electrolyte contained only calcium ion and no sodium, it would be calcium metal produced as the cathode product (which indeed is how metallic calcium is produced).
Calcium cycle
The calcium cycle is closely related to other cycles, particularly the carbon cycle. The reaction of atmospheric carbon with water forms carbonic acid, commonly referred to as ‘acid rain’ that falls to the ground, resulting in the chemical weathering of rock. The breakdown of rock, namely carbonate and silicate rocks releases calcium into soil. As a result it can then enter marine environments from ground runoff and affect many ecosystems.
Calcium carbonate
In ceramics/glazing applications, calcium carbonate is known as "whiting", and is a common ingredient for many glazes in its white powdered form. When a glaze containing this material is fired in a kiln, the whiting acts as a flux material in the glaze. Ground calcium carbonate is an abrasive (both as scouring powder and as an ingredient of household scouring creams), in particular in its calcite form, which has the relatively low hardness level of 3 on the Mohs scale of mineral hardness, and will therefore not scratch glass and most other ceramics, enamel, bronze, iron, and steel, and have a moderate effect on softer metals like aluminium and copper. A paste made from calcium carbonate and deionized water can be used to clean tarnish on silver.
Calcium
Due to the large size of the Ca ion, high coordination numbers are common, up to 24 in some intermetallic compounds such as CaZn. Calcium is readily complexed by oxygen chelates such as EDTA and polyphosphates, which are useful in analytic chemistry and removing calcium ions from hard water. In the absence of steric hindrance, smaller group 2 cations tend to form stronger complexes, but when large polydentate macrocycles are involved the trend is reversed.
Calcium cycle
Acidic rain does however release soluble calcium cations. This makies it available for plants and animals to use, as absorption is only possible in the ionic stage. The calcium may also runoff into marine environments, enabling the formation of shells and corals once reacted to form calcium carbonate.
Soft water
The hardness of water is determined by the concentration of multivalent cations in the water. Common cations found in hard water include Ca and Mg. These ions enter a water supply by leaching from minerals within an aquifer. Common calcium-containing minerals are calcite and gypsum. A common magnesium mineral is dolomite (which also contains calcium). Rainwater and distilled water are soft, because they contain few ions.
Hyalin
Hyalin's structure is dependent upon calcium ions. It stabilizes against denaturation from the high concentrations of NaCl in seawater. Stabilization happens when concentrations are as low as 1mM. Calcium also causes hyalin to precipitate and form aggregates with itself and other proteins. Doing this would require higher concentrations of calcium. Another divalent ion, Mg2+, causes further precipitation of hyalin. When acting alone, magnesium cannot cause precipitation, but increases the effect of calcium precipitation.
Monocalcium aluminate
Monocalcium aluminate is formed when the appropriate proportions of calcium carbonate and aluminium oxide are heated together until the mixture melts. It melts incongruently at 1390 °C. The crystal is monoclinic and pseudohexagonal, and has density 2945 kg.m. In calcium aluminate cements, it exists as a solid solution in which the amount of minor elements depends upon the bulk composition of the cement. A typical composition is CaAlFeSiO. It reacts rapidly with water, forming the metastable hydrate CaO·AlO·10HO, or a mixture of 2CaO·AlO·8HO, 3CaO·AlO·6HO and Al(OH) gel. These reactions form the first stage of strength development in calcium aluminate cements.
Alkali metal
All the alkali metals react vigorously or explosively with cold water, producing an aqueous solution of a strongly basic alkali metal hydroxide and releasing hydrogen gas. This reaction becomes more vigorous going down the group: lithium reacts steadily with effervescence, but sodium and potassium can ignite and rubidium and caesium sink in water and generate hydrogen gas so rapidly that shock waves form in the water that may shatter glass containers. When an alkali metal is dropped into water, it produces an explosion, of which there are two separate stages. The metal reacts with the water first, breaking the hydrogen bonds in the water and producing hydrogen gas; this takes place faster for the more reactive heavier alkali metals. Second, the heat generated by the first part of the reaction often ignites the hydrogen gas, causing it to burn explosively into the surrounding air. This secondary hydrogen gas explosion produces the visible flame above the bowl of water, lake or other body of water, not the initial reaction of the metal with water (which tends to happen mostly under water). The alkali metal hydroxides are the most basic known hydroxides.
Taste
Other monovalent cations, e.g. ammonium (NH), and divalent cations of the alkali earth metal group of the periodic table, e.g. calcium (Ca), ions generally elicit a bitter rather than a salty taste even though they, too, can pass directly through ion channels in the tongue, generating an action potential. But the chloride of calcium is saltier and less bitter than potassium chloride, and is commonly used in pickle brine instead of KCl.
Concrete degradation
In very old concrete where the calcium hydroxide has been leached from the leachate seepage path, the chemistry may revert to that similar to "speleothem" chemistry in limestone cave. This is where carbon dioxide enriched rain or seepage water forms a weak carbonic acid, which leaches calcium carbonate (CaCO) from within the concrete structure and carries it to the underside of the structure. When it contacts the atmosphere, carbon dioxide is degasses and calcium carbonate is precipitated to create calthemite deposits, which mimic the shapes and forms of speleothems. This degassing chemistry is not common in concrete structures as the leachate can often find new paths through the concrete to access free calcium hydroxide and this reverts the chemistry to that previously mentioned where CO is the reactant.
Hard water
Temporary hardness is a type of water hardness caused by the presence of dissolved bicarbonate minerals (calcium bicarbonate and magnesium bicarbonate). When dissolved, these minerals yield calcium and magnesium cations (Ca, Mg) and carbonate and bicarbonate anions (CO, HCO). The presence of the metal cations makes the water hard. However, unlike the permanent hardness caused by sulphate and chloride compounds, this "temporary" hardness can be reduced either by boiling the water, or by the addition of lime (calcium hydroxide) through the process of lime softening. Boiling promotes the formation of carbonate from the bicarbonate and precipitates calcium carbonate out of solution, leaving water that is softer upon cooling.
Calcium aluminate cements
The reactive aluminates react with water initially to form a mixture of:    CaO · AlO · 10 HO, 2 CaO · AlO · 8 HO, 3 CaO · AlO · 6 HO, and Al(OH) gel, the amounts of each depending upon the curing temperature. 
The first two hydrates subsequently decompose to a mixture of 3 CaO · AlO · 6 HO, Al(OH) gel and water, this process being called "conversion". Because of the loss of water, conversion causes an increase in porosity, which can be accompanied by a decrease in strength. This need not be a problem in structural concrete provided that a sufficiently high cement content and a sufficiently low water/cement ratio are employed.
Limescale
The type found deposited on the heating elements of water heaters has a main component of calcium carbonate. Hard water contains calcium (and often magnesium) bicarbonate or similar ions. Calcium salts, such as calcium bicarbonate and calcium carbonate are both more soluble in hot water than cold water. Thus, heating water does not cause calcium carbonate to precipitate per se. However, there is an equilibrium between dissolved calcium bicarbonate and dissolved calcium carbonate:
Calcium stearate
It is also the main component of soap scum, a white solid that forms when soap is mixed with hard water. Unlike soaps containing sodium and potassium, calcium stearate is insoluble in water and does not lather well.
Commercially it is sold as a 50% dispersion in water or as a spray dried powder. As a food additive it is known by the generic E number E470.
Period 4 element
Calcium (Ca) is the second element in period 4, between potassium and scandium. An alkali earth metal, calcium is almost never found in nature due to its high reactivity with water. It has one of the most widely known and acknowledged biological roles in all animals and some plants, making up bones and teeth, and used in some applications in cells, such as signals for cellular processeses. It is regarded as the most abundant mineral in the body's mass.
Hydrochloric acid
Hydrochloric acid has been used for dissolving calcium carbonate, i.e. such things as de-scaling kettles and for cleaning mortar off brickwork, but it is a hazardous liquid which must be used with care. When used on brickwork the reaction with the mortar only continues until the acid has all been converted, producing calcium chloride, carbon dioxide, and water:
Potassium
Because of the sensitivity of potassium to water and air, reactions with other elements are possible only in an inert atmosphere such as argon gas using air-free techniques. Potassium does not react with most hydrocarbons such as mineral oil or kerosene. It readily dissolves in liquid ammonia, up to 480 g per 1000 g of ammonia at 0 °C. Depending on the concentration, the ammonia solutions are blue to yellow, and their electrical conductivity is similar to that of liquid metals. In a pure solution, potassium slowly reacts with ammonia to form , but this reaction is accelerated by minute amounts of transition metal salts. Because it can reduce the salts to the metal, potassium is often used as the reductant in the preparation of finely divided metals from their salts by the Rieke method. For example, the preparation of magnesium by this method employs potassium as the reductant:
Calcium
Calcium is a very ductile silvery metal (sometimes described as pale yellow) whose properties are very similar to the heavier elements in its group, strontium, barium, and radium. A calcium atom has twenty electrons, arranged in the electron configuration [Ar]4s. Like the other elements placed in group 2 of the periodic table, calcium has two valence electrons in the outermost s-orbital, which are very easily lost in chemical reactions to form a dipositive ion with the stable electron configuration of a noble gas, in this case argon. Hence, calcium is almost always divalent in its compounds, which are usually ionic. Hypothetical univalent salts of calcium would be stable with respect to their elements, but not to disproportionation to the divalent salts and calcium metal, because the enthalpy of formation of MX is much higher than those of the hypothetical MX. This occurs because of the much greater lattice energy afforded by the more highly charged Ca cation compared to the hypothetical Ca cation.
Calthemite
If the leachate finds a new path through micro cracks in old concrete, this could provide a new source of calcium hydroxide (Ca(OH)) which can change the dominant reaction back to [equation 2]. The chemistry of concrete degradation is quite complex and only the chemistry relating to calcium carbonate deposition is considered in [Equations 1 to 9]. Calcium is also part of other hydration products in concrete, such as calcium aluminium hydrates and calcium aluminium iron hydrate. The chemical [Equations 1 to 4] are responsible for creating the majority of calthemite stalactites, stalagmites, flowstone etc., found on manmade concrete structures.
Concrete degradation
When water flows through cracks present in concrete, water may dissolve various minerals present in the hardened cement paste or in the aggregates, if the solution is unsaturated with respect to them. Dissolved ions, such as calcium (Ca), are leached out and transported in solution some distance. If the physico-chemical conditions prevailing in the seeping water evolve with distance along the water path and water becomes supersaturated with respect to certain minerals, they can further precipitate, making calthemite deposits (predominately calcium carbonate) inside the cracks, or at the concrete outer surface. This process can cause the self-healing of fractures in particular conditions.
Sodium
Sodium forms alloys with many metals, such as potassium, calcium, lead, and the group 11 and 12 elements. Sodium and potassium form KNa and NaK. NaK is 40–90% potassium and it is liquid at ambient temperature. It is an excellent thermal and electrical conductor. Sodium-calcium alloys are by-products of the electrolytic production of sodium from a binary salt mixture of NaCl-CaCl and ternary mixture NaCl-CaCl-BaCl. Calcium is only partially miscible with sodium. In a liquid state, sodium is completely miscible with lead. There are several methods to make sodium-lead alloys. One is to melt them together and another is to deposit sodium electrolytically on molten lead cathodes. NaPb, NaPb, NaPb, NaPb, and NaPb are some of the known sodium-lead alloys. Sodium also forms alloys with gold (NaAu) and silver (NaAg). Group 12 metals (zinc, cadmium and mercury) are known to make alloys with sodium. NaZn and NaCd are alloys of zinc and cadmium. Sodium and mercury form NaHg, NaHg, NaHg, NaHg, and NaHg.
Hard water
Several indices are used to describe the behaviour of calcium carbonate in water, oil, or gas mixtures.
Chemical equation
There are a few acid/base reactions that produce a precipitate in addition to the water molecule shown above. An example is the reaction of barium hydroxide with phosphoric acid, which produces not only water but also the insoluble salt barium phosphate. In this reaction, there are no spectator ions, so the net ionic equation is the same as the full ionic equation.
Calcium disilicide
Calcium disilicide (CaSi) is an inorganic compound, a silicide of calcium. It is a whitish or dark grey to black solid matter with melting point 1033 °C. It is insoluble in water, but may decompose when subjected to moisture, evolving hydrogen and producing calcium hydroxide. Decomposes in hot water. It is flammable and may ignite spontaneously in air.
Calcium monohydride
Gaseous calcium reacts with formaldehyde at temperatures around 1200 K to make CaH as well as some CaOH and CaO. This reaction glows orange-red.
Barium chloride
In aqueous solution BaCl behaves as a simple salt; in water it is a 1:2 electrolyte and the solution exhibits a neutral pH. Its solutions react with sulfate ion to produce a thick white precipitate of barium sulfate.
Oxalate effects a similar reaction:
When it is mixed with sodium hydroxide, it gives the dihydroxide, which is moderately soluble in water.
Water-reactive substances
Water-reactive substances are those that spontaneously undergo a chemical reaction with water, as they are highly reducing in nature. Notable examples include alkali metals, sodium through caesium, and alkaline earth metals, magnesium through barium.
Conservation and restoration of ceramic objects
Water can dissolve or deform ceramics that have been low fired, i.e. temperatures around 600 °C. Ceramic fired in high temperatures may also be susceptible to water if their mineral particles are soluble in water, for example Gypsum or calcite. Additionally the different compounds in water can flux and react in different ways to different ceramics. In naturally occurring water, carbon dioxide is dissolved and can create a chemical reaction with minerals in clay bodies that may form calcium bicarbonate which is very soluble. Stagnant water is less damaging because the carbon dioxide is not exhausted.
Carbonate compensation depth
The exact value of the CCD depends on the solubility of calcium carbonate which is determined by temperature, pressure and the chemical composition of the water – in particular the amount of dissolved in the water. Calcium carbonate is more soluble at lower temperatures and at higher pressures. It is also more soluble if the concentration of dissolved is higher. Adding a reactant to the above chemical equation pushes the equilibrium towards the right producing more products: Ca and HCO, and consuming more reactants and calcium carbonate according to Le Chatelier's principle.
Calcium
Natural calcium is a mixture of five stable isotopes (Ca, Ca, Ca, Ca, and Ca) and one isotope with a half-life so long that it can be considered stable for all practical purposes (Ca, with a half-life of about 4.3 × 10 years). Calcium is the first (lightest) element to have six naturally occurring isotopes.
Tricalcium aluminate
Water reacts instantly with tricalcium aluminate. Hydration likely begins already during grinding of cement clinker due to residual humidity and dehydration of gypsum additives. Initial contact with water causes protonation of single bonded oxygen atoms on aluminate rings and leads to the formation of calcium hydroxide. The next steps in the sequence of the hydration reaction involve the generated hydroxide ions as strong nucleophiles, which fully hydrolyze the ring structure in combination with water.
Gypsum
Gypsum is moderately water-soluble (~2.0–2.5 g/l at 25 °C) and, in contrast to most other salts, it exhibits retrograde solubility, becoming less soluble at higher temperatures. When gypsum is heated in air it loses water and converts first to calcium sulfate hemihydrate, (bassanite, often simply called "plaster") and, if heated further, to anhydrous calcium sulfate (anhydrite). As for anhydrite, its solubility in saline solutions and in brines is also strongly dependent on NaCl (common table salt) concentration.
Calcium metabolism
The plasma "total calcium" concentration is in the range of 2.2-2.6 mmol/L (9-10.5 mg/dL), and the normal "ionized calcium" is 1.3-1.5 mmol/L (4.5-5.6 mg/dL). The amount of total calcium in the blood varies with the level of plasma albumin, the most abundant protein in plasma, and therefore the main carrier of protein-bound calcium in the blood. The biologic effect of calcium is, however, determined by the amount of "ionized calcium", rather than the total calcium. It is therefore the plasma ionized calcium level which is tightly regulated to remain within very narrow limits by homeostatic negative feedback systems.
Aqua regia
Experimental evidence reveals that the reaction of platinum with aqua regia is considerably more complex. The initial reactions produce a mixture of chloroplatinous acid (HPtCl) and nitrosoplatinic chloride ((NO)PtCl). The nitrosoplatinic chloride is a solid product. If full dissolution of the platinum is desired, repeated extractions of the residual solids with concentrated hydrochloric acid must be performed.
Hierve el Agua
The rock formations are often described as "cascadas de sal" (salt waterfalls), "cascadas petrificadas" (petrified waterfalls) and "cascadas pétreas" (rock waterfalls) . The falls are formed by relatively small amounts of water which is oversaturated with calcium carbonate and comes to the surface through cracks or figures on the mountainside. The water has a temperature of . As the water runs down the rock face, it forms large stalactites similar to those found in caves. The flow of water from the spring varies significantly between the dry and rainy seasons. 95% of the surface rock formations are of calcium carbonate; however, lower layers of both rock and water currents are little known as they are sealed by the above layers. The calcium carbonate gives the formations a white or near white appearance, but other minerals that can be present in the water, such as silver, barium and iron can also be found and determines the color of a particular stalactite. The calcium carbonate in the water is due to rainwater which passes underground. First it absorbs carbon dioxide, and forms molecules of carbonic acid. This acid comes into contact with underground marble and partially dissolves creating calcium bicarbonate. When the water runs above ground, the excess minerals fall out.
Calcium nitride
It reacts with water or even the moisture in air to give ammonia and calcium hydroxide:
Like sodium oxide, calcium nitride absorbs hydrogen above 350 °C:
Calthemite
The way stalactites form on concrete is due to different chemistry than those that form naturally in limestone caves and is the result of the presence of calcium oxide (CaO) in cement. Concrete is made from aggregate, sand and cement. When water is added to the mix, the calcium oxide in the cement reacts with water to form calcium hydroxide (Ca(OH)), which under the right conditions can further dissociate to form calcium (Ca) and hydroxide (OH) ions [Equation 1]. All of the following chemical reactions are reversible and several may occur simultaneously at a specific location within a concrete structure, influenced by leachate solution pH.
Calcium cycle
The naturally occurring calcium cycle has been altered by human intervention. Calcium is predominantly extracted from limestone deposits to be utilised by many industrial processes. Purification of iron ore and aluminium, replacing asbestos break lining and some coatings for electric cables. Furthermore, calcium is used within the household to maintain alkaline pH of swimming pools, counteracting acidic disinfectants and in the food production industry to produce bicarbonate soda, some wines and dough.
Calcium metabolism
Calcium is the most abundant mineral in the human body. The average adult body contains in total approximately 1 kg, 99% in the skeleton in the form of calcium phosphate salts . The extracellular fluid (ECF) contains approximately 22 mmol, of which about 9 mmol is in the plasma. Approximately 10 mmol of calcium is exchanged between bone and the ECF over a period of twenty-four hours. The concentration of calcium ions inside the cells (in the intracellular fluid) is more than 7,000 times lower than in the blood plasma (i.e. at <0.0002 mmol/L, compared with 1.4 mmol/L in the plasma)
Calcium
Some calcium compounds were known to the ancients, though their chemistry was unknown until the seventeenth century. Pure calcium was isolated in 1808 via electrolysis of its oxide by Humphry Davy, who named the element. Calcium compounds are widely used in many industries: in foods and pharmaceuticals for calcium supplementation, in the paper industry as bleaches, as components in cement and electrical insulators, and in the manufacture of soaps. On the other hand, the metal in pure form has few applications due to its high reactivity; still, in small quantities it is often used as an alloying component in steelmaking, and sometimes, as a calcium–lead alloy, in making automotive batteries.
Acid rain
Acid rain can damage buildings, historic monuments, and statues, especially those made of rocks, such as limestone and marble, that contain large amounts of calcium carbonate. Acids in the rain react with the calcium compounds in the stones to create gypsum, which then flakes off.