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Burgos

Burgos is a city in Spain located in the autonomous community of Castile and León. It is the capital and most populated municipality of the province of Burgos.

Burgos is situated in the north of the Iberian Peninsula, on the confluence of the Arlanzón river tributaries, at the edge of the central plateau. The municipality has a population of about 180,000 inhabitants. The Camino de Santiago runs through Burgos.

Founded in 884 by the second Count of Castile, Diego Rodríguez Porcelos, Burgos soon became the leading city of the embryonic County of Castile. The 11th century chieftain Rodrigo Díaz de Vivar (El Cid) had connections with the city: born near Burgos, he was raised and educated there. In a long-lasting decline from the 17th century, Burgos became the headquarters of the Francoist proto-government (1936-1939) following the start of the Spanish Civil War. Declared in 1964 as Pole of Industrial Promotion and in 1969 as Pole of Industrial Development, the city has grown since then in terms of economic activity. At the regional level, Burgos forms part of an economic axis together with the cities of Valladolid and Palencia. In 2008, the international Burgos Airport started to service commercial flights.

The Museum of Human Evolution opened here in 2010. The museum features remains of the first hominins in Europe, who lived in this area 750,000-800,000 years ago. The Cathedral of Burgos is a World Heritage Site. Burgos was selected as the "Spanish Gastronomy Capital" of 2013. In 2015 UNESCO named it "City of Gastronomy", and it has been part of the Creative Cities Network since then.

Toponymy
There are several possible origins for the toponymy. When the city was founded, the inhabitants of the surrounding country moved into the fortified village, whose Visigothic name of Burgos signified consolidated walled villages (Gothic baurgs). The cities Burgas in Bulgaria and numerous cities containing the Germanic burg "city" such as Hamburg have a similar literal composition. The city began to be called Caput Castellae ("Cabeza de Castilla" or "Head of Castile").

History
Early humans occupied sites around Burgos as early as 800,000 years ago. When the Romans took possession of what is now the province of Burgos, the site had been a Celtic city. In Roman times, it belonged to Hispania Citerior ("Hither Spain") and then to Hispania Tarraconensis. In the 5th century, the Visigoths drove back the Suebi, then the Berbers occupied almost all of Castile in the 8th century, though only for a very brief period, and left little if any trace of their occupation. King Alfonso III the Great of León reconquered it about the middle of the 9th century, and built several castles for the defence of Christendom, which was then extended through the reconquest of lost territory. The region came to be known as Castile (Latin castella), i.e. "(land of) castles".

Burgos was founded in 884 as an outpost of this expanding Christian frontier, when Diego Rodríguez "Porcelos", count of Castile, governed this territory with orders to promote the increase of the Christian population; with this end in view he gathered the inhabitants of the surrounding country into one fortified village. The city began to be called Caput Castellae ("Cabeza de Castilla" or "Head of Castile"). The county (condado) of Castile, subject to the Kings of León, continued to be governed by counts and was gradually extended; one of these counts, Fernán González, established his independence.

In the 11th century, the city became the see of the Roman Catholic Diocese of Burgos and the capital of the Kingdom of Castile. Burgos was a major stop for pilgrims on the French Way the most popular path to Santiago de Compostela and a centre of trade between the Bay of Biscay and the south, which attracted an unusually large foreign merchant population, who became part of the city oligarchy and excluded other foreigners. Throughout the 13th and 14th centuries, Burgos was a favourite seat of the kings of León and Castile and a favoured burial site. The consejo or urban commune of Burgos was firmly in the hands of an oligarchic class of caballeros villanos, the "peasant knights" of Burgos, who provided the monarchs with a mounted contingent: in 1255 and 1266 royal charters granted relief from taxes to those citizens of Burgos who owned horses and could arm themselves, provided that they continue to live within the city walls. The merchant oligarchy succeeded the cathedral chapter as the major purchasers of land after 1250; they carried on their mercantile business in common with municipal or royal functions and sent their sons to England and Flanders to gain experience in overseas trade. A few families within the hermandades or confraternities like the Sarracín and Bonifaz succeeded in monopolising the post of alcalde, or mayor; a special court, the alcalde del rey was first mentioned at Burgos in 1281. By the reign of Alfonso X, the exemption of the non-noble knights and religious corporations, combined with exorbitant gifts and grants to monasteries and private individuals, placed great stress on the economic well-being of the realm.

In the century following the conquest of Seville (1248) on the Moors, Burgos became a testing ground for royal policies of increasing power against the consejo, in part by encouraging the right to appeal from the consejo to the king. In 1285, Sancho IV added a new body to the consejo which came to dominate it: the jurado in charge of collecting taxes and overseeing public works; the king reserved the right to select its members. The city perceived that danger to its autonomy came rather from an uncontrolled aristocracy during royal minorities: Burgos joined the hermandades of cities that leagued together for mutual protection in 1295 and 1315. In the 14th century, official royal intrusion in city affairs was perceived as a palliative against outbreaks of violence by the large excluded class of smaller merchants and artisans, on whom the tax burden fell. The alguacil was the royal official instituted to judge disagreements.

On 9 June 1345, sweeping aside the city government, Alfonso XI established direct royal rule of Burgos through the Regimiento of sixteen appointed men.

In 1574, Pope Gregory XIII made the bishopric a Metropolitan archbishopric, at the request of king Felipe II.

Burgos has been the scene of many wars: with the Moors, the struggles between León and Navarre, and between Castile and Aragon. In the Peninsular War against Napoleonic France, the siege of Burgos (between 19 September to 21 October) was a scene of a withdrawal for Arthur Wellesley, 1st Duke of Wellington. Again in the 19th-century Carlist civil wars of the Spanish succession Burgos was the scene of a battle. During the Spanish Civil War, Burgos was the base of General Franco's rebel nationalist government.

Jewish History
Historically, there was a large and thriving Jewish community in Burgos. Its first documentation dates to 974. In the 13th century, Burgos was the largest Jewish center in northern Castile. Renowned Talmudists Meir Abulafia, Todros ben Joseph Abulafia, and poet Todros ben Judah Halevi Abulafiawere born in Burgos. In the latter half of the 13th century, many kabbalists began moving to Burgos. After new restrictions were imposed on Jews, and they began to suffer from more violent acts of antisemitism, many converted to Christianity. The most well-known convert from Burgos, Paul of Burgos, was originally a Spanish Jew named Solomon Halevi. During the Edict of Expulsion in 1492, most Jews in Burgos became conversos rather than fleeing Spain. Those who did went to Portugal.

Geography and climate
At an elevation of 865 m, the city of Burgos lies in the transition zone between a Warm-summer Mediterranean climate (Köppen Csb) and an Oceanic climate (Köppen Cfb), with some continental influence resulting from its distance from the sea and higher altitude. Burgos' climate features chilly and windy winters, due to altitude and an inland location, which always include snow and temperatures below freezing. Temperature ranges can be extreme and Burgos is drier than Spain's coastal regions, although there is year-round precipitation. Average annual precipitation is 546 mm and the average annual relative humidity is 72%. In winter, temperatures very often (almost every day) drop below freezing, often reaching temperatures as low as -10 °C, and snowfalls are common, while the summer months see average high temperatures of 27.5 C. The lowest recorded temperature in Burgos was -22 C on 20 January 1885. The highest recorded temperature was 39 C on 13 August 1987.

Sights
Burgos is rich in ancient churches and convents. The three most notable are the cathedral, with its chapel of the Condestables de Castilla (Lords Constable of Castile), the monastery of Las Huelgas and the Carthusian monastery of Miraflores. Minor notable churches are San Esteban, San Gil (Sancti Aegidii), San Pedro, San Cosme y San Damián, Santiago (Sancti Jacobi), San Lorenzo and San Lesmes (Adelelmi). The Convento de la Merced, occupied by the Jesuits, and the Hospital del Rey are also of historic and architectural interest.

Among the other interesting architectural structures, in the walls of the city are the gateway of Santa María, erected for the first entrance of the Emperor Charles V, and the arch of Fernán González.

Gothic Cathedral
Construction on Burgos' Gothic Cathedral began in 1221 and spanned mainly from the 13th to 15th centuries. It has been declared a UNESCO World Heritage Site. The Cathedral is the resting place for El Cid (the famous knight from Medieval Spain's history) and his wife Dona Jimena. The west front is flanked by towers terminating in octagonal spires covered with open stonework traceries. The middle section, which serves as an entrance, has three alabaster pilasters, the intercolumnar spaces bearing panel-pictures representing the martyrdom of saints. The façade possesses ornate and fantastic surface decoration.

The octagonal chapel of the Condestable, in florid, thus highly sculpted, Gothic design, has a roof finished with balustraded turrets, needle-pointed pinnacles and statues. In the lower portion, coats of arms, shields and crouching lions have been worked into the ensemble. The exterior of the sacristy is decorated with carved traceries, figures of angels and armoured knights. The elaborate tabernacle is composed of two octagonal sections in Corinthian style.

Monasterio de las Huelgas
The Monasterio de las Huelgas Reales (Monastery of the Royal Retreats) on the outskirts of the city, was founded in 1180 by king Alfonso VIII, and was begun in a pre-Gothic style, although almost every style has been introduced over many additions. The remarkable cloisters have been described as "unrivalled for beauty both of detail and design, and perhaps unsurpassed by anything in its age and style in any part of Europe" (1911 Encyclopædia Britannica). One cloister has semicircular arches with delicate and varied columns; the other has an ogival style of early Gothic. The interior of the church has enormous columns supporting its magnificent vault; the entrance is modern. This convent historically benefited from extraordinary privileges granted to its abbess by kings and popes.

Miraflores Charterhouse
The Carthusian monastery, Miraflores Charterhouse (Cartuja de Miraflores) is situated about four kilometres from the historic city center. Among the treasures of the Charterhouse are the wooden statue of St. Bruno, the wooden choir stalls in the church and the tombs of King Juan II and of his spouse, Queen Isabella of Portugal, constructed of marble and with their recumbent effigies sculpted in alabaster. Around the top frieze are statues of angels in miniature. The French soldiers in the Spanish War of Independence (1814) mutilated this work, cutting off some of the heads and carrying them away to France. King Juan II's daughters by his first wife, heiresses Princesses Catherine and Eleanor of Asturias, are also buried in the monastery.

Castle
Located on a hill 75 meters above the city, it is a medieval defensive fortress built during the time of the reconquest. Used by Castilian kings as accommodation and a place for celebrations, it later became an artillery fort and then a weapons factory. After its destruction by the French during the War of Independence in 1813, there are hardly any buildings left standing. From the Castle it is possible to see fantastic views of the city.

Museums
Burgos has a total of 10 museums, the newest being the Museum of Human Evolution (the 10th most visited museum in Spain).

Museum of Human Evolution
The Museum of Human Evolution was inaugurated on 13 July 2010. Its foundation is based on the archaeological site of Atapuerca located 20 km east of Burgos. The Atapuerca site has been designated a UNESCO World Heritage Site. It contains several caves, where fossils and stone tools of the earliest known hominids in West Europe have been found, near Atapuerca Mountains.

The museum also allows the visitor to travel to the archaeological site and biological park.

Museum of Burgos
It summarizes the history of the province of Burgos. It has important objects and documents from all the ages, starting from Atapuerca, passing to the Romans and Iberians, and finishing in the contemporary period. These include the traditional sword of El Cid.

The museum is located in the renaissances palaces, the House of Íñigo Angulo and the House of Miranda, which has a main patio that structures the museum.

Book Museum
Between the Main Square "Plaza Mayor" and the Promenade of the Espolón "Paseo del Espolón," on the side street, Travesía del Mercado nº3, you find the Fadrique de Basilea Book Museum "Museo del Libro Fadrique de Basilea", of the Burgos publisher, Siloé, a cosy museum that shows the history of the book from its first written form to its modern electronic form.

Education
State Education in Spain is free, and compulsory from 6 to 16 years. The current education system is called LOE (Ley Orgánica de Educación).

Universities

 * University of Burgos

Often abbreviated 'UBU', the University of Burgos is a public university with about 10,000 students studying over 30 different undergraduate degrees, over 20 PhD Programmes, as well as several Official Masters and other graduate courses. It was established in 1994 when it was divided from the University of Valladolid.

UBU cooperates with the Erasmus Project, a European Union (EU) student-exchange programme, and each semester hosts students from various countries across the 15-state European Union, as well as Iceland, Liechtenstein, Norway, Malta and associated countries in Eastern Europe. Additionally, university students from various regions around the world including China, Latin and North America are able to participate in exchange programmes to study abroad at the University of Burgos. In total, the University of Burgos has education agreements with over 100 international academic institutions.

Isabella I of Castile University

Isabella I of Castile University, established in 2008, is a private state-recognized university located in Burgos. It was fully accredited by the Spanish Government in 2011.

Parks and recreation
Most of the parks in the city of Burgos can be found along the river Arlanzón. Indeed, the banks of the river itself constitute a green corridor along the city. In total, there are nearly 3 square miles (1,878 acres) of parks in Burgos and one tree for every 3 inhabitants.

 * El Parque del Castillo (or the Castle Park) is one of the lungs of the city. The city grew from the southern foot of the hill of the castle. Gradually, some of the upper neighborhoods began disappearing. After the destruction of the castle, military installations were established there. In the 1950s, the engineer Jaquotot Mariano, who later became mayor of the city, decided to leave the hill completely bare. Besides the ruins of the castle, this park contains the Bird Conservation Center, as well as play areas for children and spaces for exercise. Also, there is a terraced area with gardens and an ornamental fountain. The highlight of the park is the viewpoint from which you get a great view of the city, especially the cathedral. A metal plaque on the rail at the vista point helps to identify the most characteristic places of the city. The park contains the unique view of the city of Burgos from the viewpoint of the Castle. During special events throughout the year, visitors can see fireworks displays from that same view point. The park can be reached by car where there is parking, and on foot by climbing a series of stairs starting from the church of San Esteban.

 * Fuentes Blancas is Burgos' most visited large park extending east from the city along the river Arlanzón. It contains a network of walking and biking paths and includes camping and various outdoor recreational activities.
 * El Paseo del Espolón is the tree-lined promenade that is the most emblematic of Burgos. It's situated along the banks of the river Arlanzón, and consists of a landscaped walk that goes from the theater square to the Arc de Santa Maria, passing shops and cafes.
 * El Paseo de la Isla is another one of Burgos' tree-lined garden walks. It passes by the Palacio de la Isla which served as one of Franco's headquarters during the Spanish Civil War and is currently the Language Institute of Castilian and Leonese.
 * Parque de El Parral is located at the old "Hospital del Rey", a former hospital for pilgrims of the "Way of Saint James", now part of the University of Burgos campus.
 * El paseo de la Quinta is located on the south bank of the Arlanzón River, on the old and new path of the Cartuja de Miraflores. In the natural area of Las Veguillas, the Municipal Nurseries and the Fuentes Blancas Park, a wooded area near the city and close to the extensive pine forests of La Cartuja. It gets its name from old property of Burgos prelates, known as Quinta del Arzobispo.

Economy
Burgos is one of Spain's richest cities, with a GDP above the Spain average. It is an important trade and tourist center with a sizeable manufacturing base.

Agriculture
Wheat is a major crop. The city is surrounded by wheat fields.

Secondary sector
Burgos is a very industrial city, with a secondary sector widely developed. The city has the biggest industrial park of north Spain, called Villalonquéjar.

The city is the headquarters of Grupo Antolin, designer and manufacturer of interior automotive components.

It is also the headquarters for south Europe Benteler International.

Tertiary sector
It employs the bigger percentage in the city and is represented by the public sector (production, delivery and allocation of goods and services), due to the capital status.

Culture
The city has a complete culture infrastructure, remarked by the Auditorio y Palacio de Congresos de Burgos, which opened in mid-2012.

Language
Spanish, also called Castilian, originated in Castile and spread throughout the Crown of Castile's possessions following the reconquista. It was eventually dispersed throughout the Spanish empire following the conquest and colonization of the Americas, and Spain after the union of the Catholic Monarchs's kingdoms being co-spoken with other languages. Of the many different variations and dialects found throughout the world, it's argued that the purest Spanish can be found in the region of Burgos due to the fact that the language first developed here in the 10th century.

The Spanish language can be traced back to the monastery of Valpuesta located 100 km north-east of the city of Burgos. The Valpuesta cartularies are significant in the history of the Spanish language, and their status as manuscripts containing the earliest words written in Spanish has been officially recognised. The first utterings of Spanish continued in the "Silos Glosses", recorded in the Burgos monastery of Santo Domingo de Silos.

Cuisine
Burgos province is famous in gastronomy for:

Queso de Burgos, a white cheese which is soft and unctuous (because it is made with whey). Although originally made with sheep's milk, now cow's milk or mixtures are more common. Each comarca (rural district) produces a minor variation, and the major dairies produce an industrial product that is acceptable for people with sensitive digestion.

Burgos is blessed with a moderate climate and this fresh cheese was able to be conserved there without the need for curing of more than 10 days. With the improvement of aseptic industrial production processes this can be extended to about 30 days at a cool 6 °C. Its production reaches 35,000 tons annually.

Morcilla de Burgos, a pig's-blood sausage (black pudding), is a staple country food known across the Iberian peninsula. Spiced with onions and herbs its most noticeable content is rice (often mistaken for fat) which makes it one of the lightest and healthiest products of its kind. Oral tradition says that it must be "salty, smooth and piquant" (see Spanish pages Burgos (desambiguación) for details). As with the Queso de Burgos, several comarcas or towns in the province (Cardeñadijo, Sotopalacios, Aranda de Duero, Briviesca, Covarrubias, Villarcayo, Trespaderne, Miranda de Ebro...) made their own morcillas, with minor variations between them.

Even though Burgos is not on a D.O. wine is a fundamental piece in local gastronomy thanks to nearby wine cellars from Ribera de Duero, Rioja and Arlanza D.O.

Festivals
The city's main festival is San Pedro y San Pablo (also referred to locally as "Sampedros") celebrated on June 29. Every year, for about two weeks, the city celebrates with fireworks, concerts, sports, folklore, games for children, theater and other activities.

Burgos Airport
Burgos Airport is located only 5 km from the city centre, with daily flights to Barcelona International Airport, and during spring and summer to Palma de Mallorca and Paris Orly.

Railway
The city is considered a first-class rail route through which one rail lines circulate, operated by Renfe: Madrid–Hendaye railway. The faster trains use the AVE line between Madrid and Valladolid.

2010 marked 150 years since the arrival of the first railway to the city; the first train was present on 25 October 1860. Madrid was joined with the French border and this was considered the main route of railway communication between Spain and Europe. Located at the halfway point, Burgos soon become an important hub of communications between the centre and the north.

Since December 2007, the city has been connected to the main provincial capitals by high-speed service Alvia. It also provides other services, middle and long distance, which connect with the main population centres of the country. 18 long distance trains and 12 middle distance trains circulate through the station each day, which results in around 330,000 yearly users.

The new railway station was opened on 12 December 2008 under the name Burgos-Rosa de Lima, belonging to ADIF. It is located in the neighbourhood of Villímar, northeast of the city, and was adapted for the arrival of AVE. Bus 25 takes travellers direct to the city centre.

In July 2022, AVE high speed trains started serving the city, connecting it to cities such as Madrid in 1 hour and 33 minutes with a stop in Valladolid, or Bilbao in 70 minutes. Starting in January 2023, AVE will connect Burgos to Valencia in 3 hours and 30 minutes without the need to transfer in Madrid.

Trams
A project to implement a tram for the city has been planned, which would run the length of the Boulevard, officially called the Avenida de Valencia. It would consist of a line of about 12 km long and will originate from the Burgos-Rosa de Lima railway station and end at the University of Burgos. Its cost is estimated between 80 and 120 million euros.

Due to the economic crisis and real estate, the project has been postponed temporarily until further funding becomes available. In February 2010 work began on the development of the Boulevard, which reserves a lane for public transport: first buses, and eventually the tram.

Bus
Burgos has an international bus station located in the city centre, just one hundred meters from the Cathedral. It connects the city with nearly every region of Spain, as well as with a wide number of European (Amsterdam, Basel, Bern, Bratislava, Brussels, Cologne, Hamburg, Frankfurt, Geneva, Gdańsk, Kyiv, Lisbon, Ljubljana, Lviv, Łódź, London, Lyon, Marseille, Moscow, Munich, Nice, Paris, Porto, Prague, Rotterdam, Sofia, The Hague, Warsaw, Zürich) and even North African (Casablanca, Marrakech, Rabat) cities.

Bicycle
The city has its own public bicycle rental system, called Bicibur. It has been designed by a local company and has more than 20 points of distribution throughout the city. A network of over 100 km of bicycle lanes net the city making it the 6th largest bicycle lane network in Spain and the first in kilometres per resident.

Twin towns – sister cities
Burgos is twinned with:
 * 🇫🇷 Loudun, France
 * 🇫🇷 Pessac, France

In media
Segundo de Chomón filmed a short documentary about Burgos in 1911.

Scenes from The Good, the Bad and the Ugly were filmed near Burgos, which doubled for the Southwestern United States. The location called Sad Hill Cemetery (specifically in Santo Domingo de Silos), was where the last sequence of the film was shot.

Burgos is the setting for a scene in the 2010 film The Way, directed by Emilio Estevez and starring his father Martin Sheen.

Sports
The first football team in the city, Burgos CF, was founded in 1936, reaching the top division for six seasons in the late 1970s, before disappearing due to serious economic debts in 1983. It was replaced as the main football team of the city by its reserve team, renamed Real Burgos CF, which itself ceased to compete in 1994, after three successful seasons at the top flight. A new incarnation of Burgos CF was immediately refounded, but didn't start to compete until the 1994 season, in the sixth-tier Primera Provincial division. The team gradually transformed into professionalism, and since 2021 it plays in the second division of the Spanish football pyramid, hosting games at the El Plantío stadium.

The city has two professional basketball teams, CB San Pablo and CB Tizona, both playing (as of 2024) in the second division of the Spanish basketball league system.

license: Creative Commons - Attribution Share-Alike - https://creativecommons.org/licenses/by-sa/4.0/

title: Burgos


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Calcium carbonate

Calcium carbonate is a chemical compound with the chemical formula CaCO3. It is a common substance found in rocks as the minerals calcite and aragonite, most notably in chalk and limestone, eggshells, gastropod shells, shellfish skeletons and pearls. Materials containing much calcium carbonate or resembling it are described as calcareous. Calcium carbonate is the active ingredient in agricultural lime and is produced when calcium ions in hard water react with carbonate ions to form limescale. It has medical use as a calcium supplement or as an antacid, but excessive consumption can be hazardous and cause hypercalcemia and digestive issues.

Chemistry
Calcium carbonate shares the typical properties of other carbonates. Notably it
 * reacts with acids, releasing carbonic acid which quickly disintegrates into carbon dioxide and water:
 * CaCO3(s) + 2 H+(aq) → Ca(2+)(aq) + CO2(g) + H2O(l)

 * releases carbon dioxide upon heating, called a thermal decomposition reaction, or calcination (to above 840 °C in the case of CaCO3), to form calcium oxide, CaO, commonly called quicklime, with reaction enthalpy 178 kJ/mol:
 * CaCO3(s) → CaO(s) + CO2(g)

 * reacts with gaseous hydrogen to form methane and water vapor plus solid calcium oxide or calcium hydroxide depending on temperature and product gas composition. Various metals including palladium and nickel are catalysts for the reaction.

Calcium carbonate reacts with water that is saturated with carbon dioxide to form the soluble calcium bicarbonate.
 * CaCO3(s) + CO2(g) + H2O(l) → Ca(HCO3)2(aq)

This reaction is important in the erosion of carbonate rock, forming caverns, and leads to hard water in many regions.

An unusual form of calcium carbonate is the hexahydrate ikaite, CaCO3*6H2O. Ikaite is stable only below 8 °C.

Preparation
The vast majority of calcium carbonate used in industry is extracted by mining or quarrying. Pure calcium carbonate (such as for food or pharmaceutical use), can be produced from a pure quarried source (usually marble).

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) This process is called carbonatation:

 * CaO + H2O → Ca(OH)2
 * Ca(OH)2 + CO2 → CaCO3 + H2O

In a laboratory, calcium carbonate can easily be crystallized from calcium chloride (CaCl2), by placing an aqueous solution of CaCl2 in a desiccator alongside ammonium carbonate [NH4]2CO3. In the desiccator, ammonium carbonate is exposed to air and decomposes into ammonia, carbon dioxide, and water. The carbon dioxide then diffuses into the aqueous solution of calcium chloride, reacts with the calcium ions and the water, and forms calcium carbonate.

Structure
The thermodynamically stable form of CaCO3 under normal conditions is hexagonal β-CaCO3 (the mineral calcite). Other forms can be prepared, the denser (2.83 g/cm3) orthorhombic λ-CaCO3 (the mineral aragonite) and hexagonal μ-CaCO3, occurring as the mineral vaterite. The aragonite form can be prepared by precipitation at temperatures above 85 °C; the vaterite form can be prepared by precipitation at 60 °C. Calcite contains calcium atoms coordinated by six oxygen atoms; in aragonite they are coordinated by nine oxygen atoms. The vaterite structure is not fully understood. Magnesium carbonate (MgCO3) has the calcite structure, whereas strontium carbonate (SrCO3) and barium carbonate (BaCO3) adopt the aragonite structure, reflecting their larger ionic radii.

Polymorphs
Calcium carbonate crystallizes in three anhydrous polymorphs, of which calcite is the thermodynamically most stable at room temperature, aragonite is only slightly less so, and vaterite is the least stable.

Crystal structure
The calcite crystal structure is trigonal, with space group R$0$c (No. 167 in the International Tables for Crystallography ), and Pearson symbol hR10. Aragonite is orthorhombic, with space group Pmcn (No 62), and Pearson Symbol oP20. Vaterite is composed of at least two different coexisting crystallographic structures. The major structure exhibits hexagonal symmetry in space group P63/mmc, the minor structure is still unknown.

Crystallization
All three polymorphs crystallize simultaneously from aqueous solutions under ambient conditions. In additive-free aqueous solutions, calcite forms easily as the major product, while aragonite appears only as a minor product.

At high saturation, vaterite is typically the first phase precipitated, which is followed by a transformation of the vaterite to calcite. This behavior seems to follow Ostwald's rule, in which the least stable polymorph crystallizes first, followed by the crystallization of different polymorphs via a sequence of increasingly stable phases. However, aragonite, whose stability lies between those of vaterite and calcite, seems to be the exception to this rule, as aragonite does not form as a precursor to calcite under ambient conditions. Aragonite occurs in majority when the reaction conditions inhibit the formation of calcite and/or promote the nucleation of aragonite. For example, the formation of aragonite is promoted by the presence of magnesium ions, or by using proteins and peptides derived from biological calcium carbonate. Some polyamines such as cadaverine and Poly(ethylene imine) have been shown to facilitate the formation of aragonite over calcite.

Selection by organisms
Organisms, such as molluscs and arthropods, have shown the ability to grow all three crystal polymorphs of calcium carbonate, mainly as protection (shells) and muscle attachments. Moreover, they exhibit a remarkable capability of phase selection over calcite and aragonite, and some organisms can switch between the two polymorphs. The ability of phase selection is usually attributed to the use of specific macromolecules or combinations of macromolecules by such organisms.

Geological sources
Calcite, aragonite and vaterite are pure calcium carbonate minerals. Industrially important source rocks which are predominantly calcium carbonate include limestone, chalk, marble and travertine.

Biological sources
Eggshells, snail shells and most seashells are predominantly calcium carbonate and can be used as industrial sources of that chemical. Oyster shells have enjoyed recent recognition as a source of dietary calcium, but are also a practical industrial source. Dark green vegetables such as broccoli and kale contain dietarily significant amounts of calcium carbonate, but they are not practical as an industrial source.

Annelids in the family Lumbricidae, earthworms, possess a regionalization of the digestive track called calciferous glands, Kalkdrüsen, or glandes de Morren, that processes calcium and CO2 into calcium carbonate, which is later excreted into the dirt. The function of these glands is unknown but is believed to serve as a CO2 regulation mechanism within the animals' tissues. This process is ecologically significant, stabilizing the pH of acid soils.

Extraterrestrial
Beyond Earth, strong evidence suggests the presence of calcium carbonate on Mars. Signs of calcium carbonate have been detected at more than one location (notably at Gusev and Huygens craters). This provides some evidence for the past presence of liquid water.

Geology
Carbonate is found frequently in geologic settings and constitutes an enormous carbon reservoir. Calcium carbonate occurs as aragonite, calcite and dolomite as significant constituents of the calcium cycle. The carbonate minerals form the rock types: limestone, chalk, marble, travertine, tufa, and others. In warm, clear tropical waters corals are more abundant than towards the poles where the waters are cold. Calcium carbonate contributors, including plankton (such as coccoliths and planktic foraminifera), coralline algae, sponges, brachiopods, echinoderms, bryozoa and mollusks, are typically found in shallow water environments where sunlight and filterable food are more abundant. Cold-water carbonates do exist at higher latitudes but have a very slow growth rate. The calcification processes are changed by ocean acidification.

Where the oceanic crust is subducted under a continental plate sediments will be carried down to warmer zones in the asthenosphere and lithosphere. Under these conditions calcium carbonate decomposes to produce carbon dioxide which, along with other gases, give rise to explosive volcanic eruptions.

Carbonate compensation depth
The carbonate compensation depth (CCD) is the point in the ocean where the rate of precipitation of calcium carbonate is balanced by the rate of dissolution due to the conditions present. Deep in the ocean, the temperature drops and pressure increases. Increasing pressure also increases the solubility of calcium carbonate. Calcium carbonate is unusual in that its solubility increases with decreasing temperature. The carbonate compensation depth ranges from 4,000 to 6,000 meters below sea level in modern oceans, and the various polymorphs (calcite, aragonite) have different compensation depths based on their stability.

Role in taphonomy
Calcium carbonate can preserve fossils through permineralization. Most of the vertebrate fossils of the Two Medicine Formation—a geologic formation known for its duck-billed dinosaur eggs—are preserved by CaCO3 permineralization. This type of preservation conserves high levels of detail, even down to the microscopic level. However, it also leaves specimens vulnerable to weathering when exposed to the surface.

Trilobite populations were once thought to have composed the majority of aquatic life during the Cambrian, due to the fact that their calcium carbonate-rich shells were more easily preserved than those of other species, which had purely chitinous shells.

Construction
The main use of calcium carbonate is in the construction industry, either as a building material, or limestone aggregate for road building, as an ingredient of cement, or as the starting material for the preparation of builders' lime by burning in a kiln. However, because of weathering mainly caused by acid rain, calcium carbonate (in limestone form) is no longer used for building purposes on its own, but only as a raw primary substance for building materials.

Calcium carbonate is also used in the purification of iron from iron ore in a blast furnace. The carbonate is calcined in situ to give calcium oxide, which forms a slag with various impurities present, and separates from the purified iron.

In the oil industry, calcium carbonate is added to drilling fluids as a formation-bridging and filtercake-sealing agent; it is also a weighting material which increases the density of drilling fluids to control the downhole pressure. Calcium carbonate is added to swimming pools, as a pH corrector for maintaining alkalinity and offsetting the acidic properties of the disinfectant agent.

It is also used as a raw material in the refining of sugar from sugar beet; it is calcined in a kiln with anthracite to produce calcium oxide and carbon dioxide. This burnt lime is then slaked in fresh water to produce a calcium hydroxide suspension for the precipitation of impurities in raw juice during carbonatation.

Calcium carbonate in the form of chalk has traditionally been a major component of blackboard chalk. However, modern manufactured chalk is mostly gypsum, hydrated calcium sulfate CaSO4*2H2O. Calcium carbonate is a main source for growing biorock. Precipitated calcium carbonate (PCC), pre-dispersed in slurry form, is a common filler material for latex gloves with the aim of achieving maximum saving in material and production costs.

Fine ground calcium carbonate (GCC) is an essential ingredient in the microporous film used in diapers and some building films, as the pores are nucleated around the calcium carbonate particles during the manufacture of the film by biaxial stretching. GCC and PCC are used as a filler in paper because they are cheaper than wood fiber. Printing and writing paper can contain 10–20% calcium carbonate. In North America, calcium carbonate has begun to replace kaolin in the production of glossy paper. Europe has been practicing this as alkaline papermaking or acid-free papermaking for some decades. PCC used for paper filling and paper coatings is precipitated and prepared in a variety of shapes and sizes having characteristic narrow particle size distributions and equivalent spherical diameters of 0.4 to 3 micrometers.

Calcium carbonate is widely used as an extender in paints, in particular matte emulsion paint where typically 30% by weight of the paint is either chalk or marble. It is also a popular filler in plastics. Some typical examples include around 15–20% loading of chalk in unplasticized polyvinyl chloride (uPVC) drainpipes, 5–15% loading of stearate-coated chalk or marble in uPVC window profile. PVC cables can use calcium carbonate at loadings of up to 70 phr (parts per hundred parts of resin) to improve mechanical properties (tensile strength and elongation) and electrical properties (volume resistivity). Polypropylene compounds are often filled with calcium carbonate to increase rigidity, a requirement that becomes important at high usage temperatures. Here the percentage is often 20–40%. It also routinely used as a filler in thermosetting resins (sheet and bulk molding compounds) and has also been mixed with ABS, and other ingredients, to form some types of compression molded "clay" poker chips. Precipitated calcium carbonate, made by dropping calcium oxide into water, is used by itself or with additives as a white paint, known as whitewashing.

Calcium carbonate is added to a wide range of trade and do it yourself adhesives, sealants, and decorating fillers. Ceramic tile adhesives typically contain 70% to 80% limestone. Decorating crack fillers contain similar levels of marble or dolomite. It is also mixed with putty in setting stained glass windows, and as a resist to prevent glass from sticking to kiln shelves when firing glazes and paints at high temperature.

In ceramic glaze 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, 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.

Health and diet
Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement for gastric antacid (such as Tums and Eno). It may be used as a phosphate binder for the treatment of hyperphosphatemia (primarily in patients with chronic kidney failure). It is used in the pharmaceutical industry as an inert filler for tablets and other pharmaceuticals.

Calcium carbonate is used in the production of calcium oxide as well as toothpaste and has seen a resurgence as a food preservative and color retainer, when used in or with products such as organic apples.

Calcium carbonate is used therapeutically as phosphate binder in patients on maintenance haemodialysis. It is the most common form of phosphate binder prescribed, particularly in non-dialysis chronic kidney disease. Calcium carbonate is the most commonly used phosphate binder, but clinicians are increasingly prescribing the more expensive, non-calcium-based phosphate binders, particularly sevelamer.

Excess calcium from supplements, fortified food, and high-calcium diets can cause milk-alkali syndrome, which has serious toxicity and can be fatal. In 1915, Bertram Sippy introduced the "Sippy regimen" of hourly ingestion of milk and cream, and the gradual addition of eggs and cooked cereal, for 10 days, combined with alkaline powders, which provided symptomatic relief for peptic ulcer disease. Over the next several decades, the Sippy regimen resulted in kidney failure, alkalosis, and hypercalcaemia, mostly in men with peptic ulcer disease. These adverse effects were reversed when the regimen stopped, but it was fatal in some patients with protracted vomiting. Milk-alkali syndrome declined in men after effective treatments for peptic ulcer disease arose. Since the 1990s it has been most frequently reported in women taking calcium supplements above the recommended range of 1.2 to 1.5 grams daily, for prevention and treatment of osteoporosis, and is exacerbated by dehydration. Calcium has been added to over-the-counter products, which contributes to inadvertent excessive intake. Excessive calcium intake can lead to hypercalcemia, complications of which include vomiting, abdominal pain and altered mental status.

As a food additive it is designated E170, and it has an INS number of 170. Used as an acidity regulator, anticaking agent, stabilizer or color it is approved for usage in the EU, US and Australia and New Zealand. It is "added by law to all UK milled bread flour except wholemeal". It is used in some soy milk and almond milk products as a source of dietary calcium; at least one study suggests that calcium carbonate might be as bioavailable as the calcium in cow's milk. Calcium carbonate is also used as a firming agent in many canned and bottled vegetable products.

Several calcium supplement formulations have been documented to contain the chemical element lead, posing a public health concern. Lead is commonly found in natural sources of calcium.

Agriculture and aquaculture
Agricultural lime, powdered chalk or limestone, is used as a cheap method of neutralising acidic soil, making it suitable for planting, also used in aquaculture industry for pH regulation of pond soil before initiating culture. There is interest in understanding whether or not it can affect pesticide adsorption and desorption in calcareous soil.

Household cleaning
Calcium carbonate is a key ingredient in many household cleaning powders like Comet and is used as a scrubbing agent.

Pollution mitigation
In 1989, a researcher, Ken Simmons, introduced CaCO3 into the Whetstone Brook in Massachusetts. His hope was that the calcium carbonate would counter the acid in the stream from acid rain and save the trout that had ceased to spawn. Although his experiment was a success, it did increase the amount of aluminium ions in the area of the brook that was not treated with the limestone. This shows that CaCO3 can be added to neutralize the effects of acid rain in river ecosystems. Currently calcium carbonate is used to neutralize acidic conditions in both soil and water. Since the 1970s, such liming has been practiced on a large scale in Sweden to mitigate acidification and several thousand lakes and streams are limed repeatedly.

Calcium carbonate is also used in flue-gas desulfurization applications eliminating harmful SO2 and NO2 emissions from coal and other fossil fuels burnt in large fossil fuel power stations.

Plastics
Calcium carbonate is commonly used in the plastic industry as a filler. When it is incorporated in a plastic material, it can improve the hardness, stiffness, dimensional stability and processability of the material.

Calcination equilibrium
Calcination of limestone using charcoal fires to produce quicklime has been practiced since antiquity by cultures all over the world. The temperature at which limestone yields calcium oxide is usually given as 825 °C, but stating an absolute threshold is misleading. Calcium carbonate exists in equilibrium with calcium oxide and carbon dioxide at any temperature. At each temperature there is a partial pressure of carbon dioxide that is in equilibrium with calcium carbonate. At room temperature the equilibrium overwhelmingly favors calcium carbonate, because the equilibrium CO2 pressure is only a tiny fraction of the partial CO2 pressure in air, which is about 0.035 kPa.

At temperatures above 550 °C the equilibrium CO2 pressure begins to exceed the CO2 pressure in air. So above 550 °C, calcium carbonate begins to outgas CO2 into air. However, in a charcoal fired kiln, the concentration of CO2 will be much higher than it is in air. Indeed, if all the oxygen in the kiln is consumed in the fire, then the partial pressure of CO2 in the kiln can be as high as 20 kPa.

The table shows that this partial pressure is not achieved until the temperature is nearly 800 °C. For the outgassing of CO2 from calcium carbonate to happen at an economically useful rate, the equilibrium pressure must significantly exceed the ambient pressure of CO2. And for it to happen rapidly, the equilibrium pressure must exceed total atmospheric pressure of 101 kPa, which happens at 898 °C.

 * {| class="wikitable"

!P (kPa) !T (°C)
 * + Equilibrium pressure of CO2 over CaCO3 (P) versus temperature (T).
 * 0.055||0.13||0.31||1.80||5.9||9.3||14||24||34||51||72 ||80||91||101||179||901||3961
 * 550||587||605||680||727||748||777||800||830||852||871||881||891||898||937||1082||1241
 * }

With varying CO2 pressure
Calcium carbonate is poorly soluble in pure water (47 mg/L at normal atmospheric CO2 partial pressure as shown below).

The equilibrium of its solution is given by the equation (with dissolved calcium carbonate on the right):
 * {| width="500"

 * style="width:50%; height:30px;"| CaCO3 ⇌ Ca(2+) + CO3(2−)
 * Ksp = $-0 cm^{3}/mol$ to $\overline{3}$ at 25 °C
 * }

where the solubility product for [Ca(2+)][CO3(2−)] is given as anywhere from Ksp = $0$ to Ksp = $0$ at 25 °C, depending upon the data source. What the equation means is that the product of molar concentration of calcium ions (moles of dissolved Ca(2+) per liter of solution) with the molar concentration of dissolved CO3(2−) cannot exceed the value of Ksp. This seemingly simple solubility equation, however, must be taken along with the more complicated equilibrium of carbon dioxide with water (see carbonic acid). Some of the CO3(2−) combines with H+ in the solution according to

 * {| width="500"

 * style="width:50%; height:25px;"| HCO3− ⇌ H+ + CO3(2−)
 * Ka2 = $0$ at 25 °C
 * }

HCO3− is known as the bicarbonate ion. Calcium bicarbonate is many times more soluble in water than calcium carbonate—indeed it exists only in solution.

Some of the HCO3− combines with H+ in solution according to

 * width=250|H2CO3 ⇌ H+ + HCO3−
 * Ka1 = $0$ at 25 °C
 * }

Some of the H2CO3 breaks up into water and dissolved carbon dioxide according to

 * width=250|H2O + CO2(aq) ⇌ H2CO3
 * Kh = $0$ at 25 °C
 * }

And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to

 * width=250|$0$ = kH
 * where kH = 29.76 atm/(mol/L) at 25 °C (Henry constant), and PCO2 is the CO2 partial pressure.
 * }

For ambient air, PCO2 is around $0.002$ atm (or equivalently 35 Pa). The last equation above fixes the concentration of dissolved CO2 as a function of PCO2, independent of the concentration of dissolved CaCO3. At atmospheric partial pressure of CO2, dissolved CO2 concentration is $P_⁄[CO2]$ moles per liter. The equation before that fixes the concentration of H2CO3 as a function of CO2 concentration. For [CO2] = $0$, it results in [H2CO3] = $0$ moles per liter. When [H2CO3] is known, the remaining three equations together with

 * {| width="450"

 * style="width:50%; height:25px;"|H2O ⇌ H+ + OH−
 * K = 10−14 at 25 °C
 * }

(which is true for all aqueous solutions), and the constraint that the solution must be electrically neutral, i.e., the overall charge of dissolved positive ions [Ca(2+)] + 2 [H+] must be cancelled out by the overall charge of dissolved negative ions [HCO3−] + [CO3(2−)] + [OH−], make it possible to solve simultaneously for the remaining five unknown concentrations (the previously mentioned form of the neutrality 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 balance is not neutral).

The adjacent table shows the result for [Ca(2+)] and [H+] (in the form of pH) as a function of ambient partial pressure of CO2 (Ksp = $0$ has been taken for the calculation).
 * At atmospheric levels of ambient CO2 the table indicates that the solution will be slightly alkaline with a maximum CaCO3 solubility of 47 mg/L.
 * As ambient CO2 partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely low PCO2, dissolved CO2, bicarbonate ion, and carbonate ion largely evaporate from the solution, leaving a highly alkaline solution of calcium hydroxide, which is more soluble than CaCO3. For PCO2 = 10−12 atm, the [Ca(2+)][OH−]2 product is still below the solubility product of Ca(OH)2 ($0$). For still lower CO2 pressure, Ca(OH)2 precipitation will occur before CaCO3 precipitation.
 * As ambient CO2 partial pressure increases to levels above atmospheric, pH drops, and much of the carbonate ion is converted to bicarbonate ion, which results in higher solubility of Ca(2+).

The effect of the latter is especially evident in day-to-day life of people who have hard water. Water in aquifers underground can be exposed to levels of CO2 much higher than atmospheric. As such water percolates through calcium carbonate rock, the CaCO3 dissolves according to one of the trends above. When that same water then emerges from the tap, in time it comes into equilibrium with CO2 levels in the air by outgassing its excess CO2. The calcium carbonate becomes less soluble as a result, and the excess precipitates as lime scale. This same process is responsible for the formation of stalactites and stalagmites in limestone caves.

Two hydrated phases of calcium carbonate, monohydrocalcite CaCO3*H2O and ikaite CaCO3*6H2O, may precipitate from water at ambient conditions and persist as metastable phases.

With varying pH, temperature and salinity: CaCO3 scaling in swimming pools
In contrast to the open equilibrium scenario above, many swimming pools are managed by addition of sodium bicarbonate (NaHCO3) to the concentration of about 2 mmol/L as a buffer, then control of pH through use of HCl, NaHSO4, Na2CO3, NaOH or chlorine formulations that are acidic or basic. In this situation, dissolved inorganic carbon (total inorganic carbon) is far from equilibrium with atmospheric CO2. Progress towards equilibrium through outgassing of CO2 is slowed by 1. the slow reaction
 * H2CO3 ⇌ CO2(aq) + H2O;

2. limited aeration in a deep water column; and

3. periodic replenishment of bicarbonate to maintain buffer capacity (often estimated through measurement of total alkalinity).

In this situation, the dissociation constants for the much faster reactions
 * H2CO3 ⇌ H+ + HCO3− ⇌ 2 H+ + CO3(2−)

allow the prediction of concentrations of each dissolved inorganic carbon species in solution, from the added concentration of HCO3− (which constitutes more than 90% of Bjerrum plot species from pH 7 to pH 8 at 25 °C in fresh water). Addition of HCO3− will increase CO3(2−) concentration at any pH. Rearranging the equations given above, we can see that [Ca(2+)] = $0$, and [CO3(2−)] = $0.005$. Therefore, when HCO3− concentration is known, the maximum concentration of Ca(2+) ions before scaling through CaCO3 precipitation can be predicted from the formula:

 * [Ca(2+)]max = $0.001$ × $0$

The solubility product for CaCO3 (Ksp) and the dissociation constants for the dissolved inorganic carbon species (including Ka2) are all substantially affected by temperature and salinity, with the overall effect that [Ca(2+)]max increases from freshwater to saltwater, and decreases with rising temperature, pH, or added bicarbonate level, as illustrated in the accompanying graphs.

The trends are illustrative for pool management, but whether scaling occurs also depends on other factors including interactions with Mg(2+), [B(OH)4]− and other ions in the pool, as well as supersaturation effects. Scaling is commonly observed in electrolytic chlorine generators, where there is a high pH near the cathode surface and scale deposition further increases temperature. This is one reason that some pool operators prefer borate over bicarbonate as the primary pH buffer, and avoid the use of pool chemicals containing calcium.

Solubility in a strong or weak acid solution
Solutions of strong (HCl), moderately strong (sulfamic) or weak (acetic, citric, sorbic, lactic, phosphoric) acids are commercially available. They are commonly used as descaling agents to remove limescale deposits. The maximum amount of CaCO3 that can be "dissolved" by one liter of an acid solution can be calculated using the above equilibrium equations.
 * In the case of a strong monoacid with decreasing acid concentration [A] = [A−], we obtain (with CaCO3 molar mass = 100 g/mol):
 * {| class="wikitable"

! [A] (mol/L) ! Initial pH !Final pH ! Dissolved CaCO3 (g/L of acid)
 * 1
 * 10−1
 * 10−2
 * 10−3
 * 10−4
 * 10−5
 * 10−6
 * 10−7
 * 10−10
 * 0.00||1.00||2.00||3.00||4.00||5.00||6.00||6.79||7.00
 * 6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27
 * 50.0||5.00||0.514||0.0849||0.0504||0.0474||0.0471||0.0470||0.0470
 * }

 * where the initial state is the acid solution with no Ca(2+) (not taking into account possible CO2 dissolution) and the final state is the solution with saturated Ca(2+). For strong acid concentrations, all species have a negligible concentration in the final state with respect to Ca(2+) and A− so that the neutrality equation reduces approximately to 2[Ca(2+)] = [A−] yielding [Ca(2+)] ≈ 0.5 [A−]. When the concentration decreases, [HCO3−] becomes non-negligible so that the preceding expression is no longer valid. For vanishing acid concentrations, one can recover the final pH and the solubility of CaCO3 in pure water.

 * In the case of a weak monoacid (here we take acetic acid with pKa = 4.76) with decreasing total acid concentration [A] = [A−] + [AH], we obtain:

 * {| class="wikitable"

![A] (mol/L) !Initial pH !Final pH !|Dissolved CaCO3 (g/L of acid)
 * [Ca(2+)] ≈ 0.5 [A−]
 * 10−1
 * 10−2
 * 10−3
 * 10−4
 * 10−5
 * 10−6
 * 10−7
 * 10−10
 * 2.38||2.88||3.39||3.91||4.47||5.15||6.02||6.79||7.00
 * 6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27
 * 49.5||4.99||0.513||0.0848||0.0504||0.0474||0.0471||0.0470||0.0470
 * }
 * For the same total acid concentration, the initial pH of the weak acid is less acid than the one of the strong acid; however, the maximum amount of CaCO3 which can be dissolved is approximately the same. This is because in the final state, the pH is larger than the pKa, so that the weak acid is almost completely dissociated, yielding in the end as many H+ ions as the strong acid to "dissolve" the calcium carbonate.

 * The calculation in the case of phosphoric acid (which is the most widely used for domestic applications) is more complicated since the concentrations of the four dissociation states corresponding to this acid must be calculated together with [HCO3−], [CO3(2−)], [Ca(2+)], [H+] and [OH−]. The system may be reduced to a seventh degree equation for [H+] the numerical solution of which gives

 * {| class="wikitable"

! [A] (mol/L) ! Initial pH ! Final pH ! Dissolved CaCO3 (g/L of acid)
 * 1
 * 10−1
 * 10−2
 * 10−3
 * 10−4
 * 10−5
 * 10−6
 * 10−7
 * 10−10
 * 1.08||1.62||2.25||3.05||4.01||5.00||5.97||6.74||7.00
 * 6.71||7.17||7.63||8.06||8.24||8.26||8.26||8.26||8.27
 * 62.0||7.39||0.874||0.123||0.0536||0.0477||0.0471||0.0471||0.0470
 * }

 * where [A] = [H3PO4] + [H2PO4−] + [HPO4(2−)] + [PO4(3−)] is the total acid concentration. Thus phosphoric acid is more efficient than a monoacid since at the final almost neutral pH, the second dissociated state concentration [HPO4(2−)] is not negligible (see phosphoric acid).

title: Calcium carbonate


text:
Emmanuelle Seigner

Emmanuelle Seigner (born 22 June 1966) is a French actress and singer. She is known for her roles in The Diving Bell and the Butterfly (2007), The Ninth Gate (1999) and Frantic (1988). She has been nominated for a César Award for Best Actress for Venus in Fur (2013), and for two César Awards for Best Supporting Actress in Place Vendôme (1998) and La Vie en Rose (2007). She has been married to Polish film director Roman Polanski since 1989.

Early life
Seigner is the daughter of Jean-Louis Seigner (1941-2020), photographer, and Aline Ponelle, journalist. She is the older sister of Mathilde Seigner, actress, and Marie-Amélie Seigner, singer. Through their father, they are the granddaughters of the actor Louis Seigner (1903-1991) and the nieces of the actress Françoise Seigner (1928-2008), both deans and members of the Comédie-Française. She is the niece of doctor Véronique Vasseur.

She was educated at a Catholic convent school, and began modelling at the age of fourteen.

Films
Her husband directed her in Frantic (1988), Bitter Moon (1992), The Ninth Gate (1999), Venus in Fur (2013), Based On A True Story (2017) and An Officer and a Spy (2019). In 2010, Seigner was featured in Jerzy Skolimowski's Essential Killing, which went on to win the Special Jury Prize at the Venice Film Festival. In 2012, she played the principal role of Vanda in Polanski's French film adaptation of David Ives's two-character play Venus in Fur, based on the Austrian novel of the near-same name, for which she received praise as Vanda, an actress, playing against writer and theatre director Thomas played by Mathieu Amalric.

Music
Seigner appears as the main character in the music video "Hands Around My Throat" by Death in Vegas. In 2006, she became the lead singer of the pop rock band Ultra Orange, and the group's name was changed to Ultra Orange & Emmanuelle. They released a self-titled album in 2007. Seigner released a solo album called  Distant Lover in 2014.

Other activities
From 2008, she was an ambassadress of the Polish brand Dr Irena Eris.

Personal life
Seigner married Roman Polanski on August 30, 1989 in the 8th arrondissement of Paris. They have two children, Morgane (born 1993) and Elvis (born 1998). She and Polanski live in Paris.

Studio albums

 * Ultra Orange & Emmanuelle (2007) Sony
 * Dingue (2010)
 * Distant Lover (2014)
 * Diabolique (as L'épée) (2019) (with Anton Newcombe & The Limiñanas)

Singles

 * "Les Mots Simples" (2008) (with Brett Anderson)

title: Emmanuelle Seigner


text:
Chalk

Chalk is a soft, white, porous, sedimentary carbonate rock. It is a form of limestone composed of the mineral calcite and originally formed deep under the sea by the compression of microscopic plankton that had settled to the sea floor. Chalk is common throughout Western Europe, where deposits underlie parts of France, and steep cliffs are often seen where they meet the sea in places such as the Dover cliffs on the Kent coast of the English Channel.

Chalk is mined for use in industry, such as for quicklime, bricks and builder's putty, and in agriculture, for raising pH in soils with high acidity. It is also used for "blackboard chalk" for writing and drawing on various types of surfaces, although these can also be manufactured from other carbonate-based minerals, or gypsum.

Description
Chalk is a fine-textured, earthy type of limestone distinguished by its light colour, softness, and high porosity. It is composed mostly of tiny fragments of the calcite shells or skeletons of plankton, such as foraminifera or coccolithophores. These fragments mostly take the form of calcite plates ranging from 0.5 to 4 microns in size, though about 10% to 25% of a typical chalk is composed of fragments that are 10 to 100 microns in size. The larger fragments include intact plankton skeletons and skeletal fragments of larger organisms, such as molluscs, echinoderms, or bryozoans.

Chalk is typically almost pure calcite, CaCO3, with just 2% to 4% of other minerals. These are usually quartz and clay minerals, though collophane (cryptocrystalline apatite, a phosphate mineral) is also sometimes present, as nodules or as small pellets interpreted as fecal pellets. In some chalk beds, the calcite has been converted to dolomite, CaMg(CO3)2, and in a few cases the dolomitized chalk has been dedolomitized back to calcite.

Chalk is highly porous, with typical values of porosity ranging from 35 to 47 per cent. While it is similar in appearance to both gypsum and diatomite, chalk is identifiable by its hardness, fossil content, and its reaction to acid (it produces effervescence on contact).

Formation
In Western Europe, chalk was formed in the Late Cretaceous Epoch and the early Palaeocene Epoch (between 100 and 61 million years ago). It was deposited on extensive continental shelves at depths between 100 and 600 m, during a time of nonseasonal (likely arid) climate that reduced the amount of erosion from nearby exposed rock. The lack of nearby erosion explains the high purity of chalk. The coccolithophores, foraminifera, and other microscopic organisms from which the chalk came mostly form low-magnesium calcite skeletons, so the sediments were already in the form of highly stable low-magnesium calcite when deposited. This is in contrast with most other limestones, which formed from high-magnesium calcite or aragonite that rapidly converted to the more stable low-magnesium calcite after deposition, resulting in the early cementation of such limestones. In chalk, absence of this calcium carbonate conversion process prevented early cementation, which partially accounts for chalk's high porosity. Chalk is also the only form of limestone that commonly shows signs of compaction.

Flint (a type of chert) is very common as bands parallel to the bedding or as nodules in seams, or linings to fractures, embedded in chalk. It is probably derived from sponge spicules or other siliceous organisms as water is expelled upwards during compaction. Flint is often deposited around larger fossils such as Echinoidea which may be silicified (i.e. replaced molecule by molecule by flint).

Geology and geographic distribution
Chalk is so common in Cretaceous marine beds that the Cretaceous Period was named for these deposits. The name Cretaceous was derived from Latin creta, meaning chalk. Some deposits of chalk were formed after the Cretaceous.

The Chalk Group is a European stratigraphic unit deposited during the late Cretaceous Period. It forms the famous White Cliffs of Dover in Kent, England, as well as their counterparts of the Cap Blanc Nez on the other side of the Dover Strait. The Champagne region of France is mostly underlain by chalk deposits, which contain artificial caves used for wine storage. Some of the highest chalk cliffs in the world occur at Jasmund National Park in Germany and at Møns Klint in Denmark.

Chalk deposits are also found in Cretaceous beds on other continents, such as the Austin Chalk, Selma Group, and Niobrara Formations of the North American interior. Chalk is also found in western Egypt (Khoman Formation) and western Australia (Miria Formation).

Chalk of Oligocene to Neogene age has been found in drill cores of rock under the Pacific Ocean at Stewart Arch in the Solomon Islands.

There are layers of chalk, containing Globorotalia, in the Nicosia Formation of Cyprus, which formed during the Pliocene.

Mining
Chalk is mined from chalk deposits both above ground and underground. Chalk mining boomed during the Industrial Revolution, due to the need for chalk products such as quicklime and bricks.

Uses
Most people first encounter chalk in school where it refers to blackboard chalk, which was originally made of mineral chalk, since it readily crumbles and leaves particles that stick loosely to rough surfaces, allowing it to make writing that can be readily erased. Blackboard chalk manufacturers now may use mineral chalk, other mineral sources of calcium carbonate, or the mineral gypsum (calcium sulfate). While gypsum-based blackboard chalk is the lowest cost to produce, and thus widely used in the developing world, use of carbonate-based chalk produces larger particles and thus less dust, and it is marketed as "dustless chalk".

Coloured chalks, pastel chalks, and sidewalk chalk (shaped into larger sticks and often coloured), used to draw on sidewalks, streets, and driveways, are primarily made of gypsum rather than calcium carbonate chalk.

Magnesium carbonate chalk is commonly used as a drying agent to obtain better grip by gymnasts and rock climbers.

Glazing putty mainly contains chalk as a filler in linseed oil.

Chalk and other forms of limestone may be used for their properties as a base. Chalk is a source of quicklime by thermal decomposition, or slaked lime following quenching of quicklime with water. In agriculture, chalk is used for raising pH in soils with high acidity. Small doses of chalk can also be used as an antacid. Additionally, the small particles of chalk make it a substance ideal for cleaning and polishing. For example, toothpaste commonly contains small amounts of chalk, which serves as a mild abrasive. Polishing chalk is chalk prepared with a carefully controlled grain size, for very fine polishing of metals.

French chalk (also known as tailor's chalk) is traditionally a hard chalk used to make temporary markings on cloth, mainly by tailors. It is now usually made of talc (magnesium silicate).

Chalk beds form important petroleum reservoirs in the North Sea and along the Gulf Coast of North America.

Previous uses
In southeast England, deneholes are a notable example of ancient chalk pits. Such bell pits may also mark the sites of ancient flint mines, where the prime object was to remove flint nodules for stone tool manufacture. The surface remains at Cissbury are one such example, but perhaps the most famous is the extensive complex at Grimes Graves in Norfolk.

Chalk was traditionally used in recreation. In field sports, such as tennis played on grass, powdered chalk was used to mark the boundary lines of the playing field or court. If a ball hits the line, a cloud of chalk or pigment dust will be visible. In recent years, powdered chalk has been replaced with titanium dioxide. In gymnastics, rock-climbing, weightlifting and tug of war, chalk — now usually magnesium carbonate — is applied to the hands and feet to remove perspiration and reduce slipping.

Chalk may also be used as a house construction material instead of brick or wattle and daub: quarried chalk was cut into blocks and used as ashlar, or loose chalk was rammed into blocks and laid in mortar. There are still houses standing which have been constructed using chalk as the main building material. Most are pre-Victorian though a few are more recent.

A mixture of chalk and mercury can be used as fingerprint powder. However, because of the toxicity of the mercury, the use of such mixtures for fingerprinting was abandoned in 1967.

title: Chalk