Instruct: Given a question, retrieve passages that answer the question.
Query: Does calcium have a reaction to water?
Calcium carbonate is a chemical compound with the formula CaCO3. 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.
The calcium sulfate (CaSO) in the tube retains carbon dioxide selectively as it's heated, and thereby, removed from the solution. The drying agent absorbs any aerosolized water and/or water vapor (reaction 3.). The mix of the drying agent and NaOH absorbs the CO and any water that may have been produced as a result of the absorption of the NaOH (reaction 4.).
While, the solubility of most inorganic salts (NaCl, KCl, ...) increases with temperature (endothermic dissolution reaction) , some inorganic salts such as calcium carbonate and calcium sulfate have also a retrograde solubility, i.e., their solubility decreases with temperature. In the case of calcium carbonate, it is due to the degassing of CO whose solubility decreases with temperature as it is the case for most of the gases (exothermic dissolution reaction in water). In the case of calcium sulfate, the reason is that the dissolution reaction of calcium sulfate itself is exothermic and therefore is favored when the temperature decreases (then, the dissolution heat is more easily evacuated, see Le Chatelier's principle). In other terms, the solubility of calcium carbonate and calcium sulfate increases at low temperature and decreases at high temperature, as it is also the case for calcium hydroxide (portlandite) often referred as a didactic case study to explain the reason of retrograde solubility. 
Scolecite is a tectosilicate mineral belonging to the zeolite group; it is a hydrated calcium silicate, CaAlSiO·3HO. Only minor amounts of sodium and traces of potassium substitute for calcium. There is an absence of barium, strontium, iron and magnesium. Scolecite is isostructural (having the same structure) with the sodium-calcium zeolite mesolite and the sodium zeolite natrolite, but it does not form a continuous chemical series with either of them. It was described in 1813, and named from the Greek word, σκώληξ (sko-lecks) = "worm" because of its reaction to the blowpipe flame.
Calcium silicate hydrate (also shown as C-S-H) is a result of the reaction between the silicate phases of Portland cement and water. This reaction typically is expressed as:
Non-hydraulic cement, such as slaked lime (calcium oxide mixed with water), hardens by carbonation in the presence of carbon dioxide, which is naturally present in the air. First calcium oxide (lime) is produced from calcium carbonate (limestone or chalk) by calcination at temperatures above 825 °C (1,517 °F) for about 10 hours at atmospheric pressure:
The calcium oxide is then "spent" (slaked) mixing it with water to make slaked lime (calcium hydroxide):
Once the excess water is completely evaporated (this process is technically called "setting"), the carbonation starts:
This reaction takes time, because the partial pressure of carbon dioxide in the air is low. The carbonation reaction requires that the dry cement be exposed to air, so the slaked lime is a non-hydraulic cement and cannot be used under water. This process is called the "lime cycle".
To make lime plaster, limestone (calcium carbonate) is heated above approximately 850 °C to produce quicklime (calcium oxide). Water is then added to produce slaked lime (calcium hydroxide), which is sold as a wet putty or a white powder. Additional water is added to form a paste prior to use. The paste may be stored in airtight containers. When exposed to the atmosphere, the calcium hydroxide very slowly turns back into calcium carbonate through reaction with atmospheric carbon dioxide, causing the plaster to increase in strength.
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).
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:
The goal of the extraction process is to obtain dry, powdered, sodium alginate. The calcium and magnesium salts do not dissolve in water; the sodium salt does. The rationale behind the extraction of alginate from the seaweed is to convert all the alginate salts to the sodium salt, dissolve this in water, and remove the seaweed residue by filtration. The alginate must then be recovered from the aqueous solution. The solution is very dilute and evaporation of the water is not economic. To the sodium alginate from the initial extraction solution, a calcium salt is added. This causes calcium alginate to form with a fibrous texture; it does not dissolve in water and can be separated from it with relative ease using a metal screen.
The most common stalactites are speleothems, which occur in limestone caves. They form through deposition of calcium carbonate and other minerals, which is precipitated from mineralized water solutions. Limestone is the chief form of calcium carbonate rock which is dissolved by water that contains carbon dioxide, forming a calcium bicarbonate solution in underground caverns. The chemical formula for this reaction is:
Reaction 5:
"Hydrogen + calcium carbonate → methane + calcium oxide + water "
The following equilibrium reaction describes the dissolving and formation of calcium carbonate and calcium bicarbonate (on the right):
Calcium carbonate (limestone) is decomposed by heat to produce calcium oxide (quicklime) and carbon dioxide gas. Then calcium oxide reacts with water to form calcium hydroxide (slaked lime) which is accompanied by the release of heat, a reaction known as exothermic.
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.
The reaction of calcium carbide with water, producing acetylene and calcium hydroxide, was discovered by Friedrich Wöhler in 1862.
A pozzolanic material is by definition capable of binding calcium hydroxide in the presence of water. Therefore, the chemical measurement of this pozzolanic activity represents a way of evaluating pozzolanic materials. This can be done by directly measuring the amount of calcium hydroxide a pozzolan consumes over time. At high water to binder ratio (suspended solutions), this can be measured by titrimety or by spectroscopic techniques. At lower water to binder ratios (pastes), thermal analysis or X-ray powder diffraction techniques are commonly used to determine remaining calcium hydroxide contents. Other direct methods have been developed that aim to directly measure the degree of reaction of the pozzolan itself. Here, selective dissolutions, X-ray powder diffraction or scanning electron microscopy image analysis methods have been used.
Calcium oxide (lime) is reacted with water to regenerate the calcium hydroxide used in Reaction 2:
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. 
One significant application of calcium hydroxide is as a flocculant, in water and sewage treatment. It forms a fluffy charged solid that aids in the removal of smaller particles from water, resulting in a clearer product. This application is enabled by the low cost and low toxicity of calcium hydroxide. It is also used in fresh water treatment for raising the pH of the water so that pipes will not corrode where the base water is acidic, because it is self-regulating and does not raise the pH too much.
The main use of calcium sulfate is to produce plaster of Paris and stucco. These applications exploit the fact that calcium sulfate which has been powdered and calcined forms a moldable paste upon hydration and hardens as crystalline calcium sulfate dihydrate. It is also convenient that calcium sulfate is poorly soluble in water and does not readily dissolve in contact with water after its solidification.
The haloform reaction is one of the oldest organic reactions known. In 1822, Georges-Simon Serullas added potassium metal to a solution of iodine in ethanol and water to form potassium formate and iodoform, called in the language of that time "hydroiodide of carbon". In 1831, Justus von Liebig reported the reaction of chloral with calcium hydroxide to form chloroform and calcium formate. The reaction was rediscovered by Adolf Lieben in 1870. The iodoform test is also called the Lieben haloform reaction. A review of the Haloform reaction with a history section was published in 1934.
The cause of calciphylaxis is unknown. It does not seem to be an immune type reaction. In other words, calciphylaxis is not a hypersensitivity reaction (i.e., allergic reaction) leading to sudden local calcification. Clearly, additional factors are involved in calciphylaxis. It is also known as calcific uremic arteriolopathy; however, the disease is not limited to patients with kidney failure. The current belief is that in end-stage kidney disease, abnormal calcium and phosphate homeostasis result in the deposition of calcium in the vessels, also known as metastatic calcification. Once the calcium has been deposited, a thrombotic event occurs within the lumen of these vessels, resulting in tissue infarction. It is unknown what the triggers are that cause the thrombotic and ischemic event. Reported risk factors include female sex, obesity, elevated calcium-phosphate product, medications such as warfarin, vitamin D derivatives e.g. calcitriol, calcium-based binders, or systemic steroids, protein C or S deficiency, low blood albumin levels, and diabetes mellitus.
Commercial heat sources for self-heating food packaging use an exothermic (heat releasing) reaction between quicklime, or calcium oxide, and water. Quicklime, inexpensive and readily available, is generally recognized by the FDA as safe. The product of the reaction is calcium hydroxide.
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.