Acarbonate fluoride,fluoride carbonate,fluorocarbonate orfluocarbonate is a double salt containing bothcarbonate andfluoride. The salts are usually insoluble in water, and can have more than one kind of metal cation to make more complex compounds.Rare-earth fluorocarbonates are particularly important as ore minerals for the light rare-earth elementslanthanum,cerium andneodymium.Bastnäsite is the most important source of these elements. Other artificial compounds are under investigation asnon-linear optical materials and for transparency in theultraviolet, with effects over a dozen times greater thanPotassium dideuterium phosphate.^[1]

Related to this there are alsochlorocarbonates andbromocarbonates. Along with these fluorocarbonates form the larger family of halocarbonates. In turn halocarbonates are a part ofmixed anion materials. Compounds where fluorine connects to carbon making acids are unstable,fluoroformic acid decomposes to carbon dioxide and hydrogen fluoride, andtrifluoromethyl alcohol also breaks up at room temperature.Trifluoromethoxide compounds exist but react with water to yieldcarbonyl fluoride.

The structure of the carbonate fluorides is mainly determined by the carbonate anion, as it is the biggest component. The overall structure depends on the ratio of carbonate to everything else, i.e. number (metals and fluorides)/number of carbonates. For ratios from 1.2 to 1.5 the carbonates are in a flat dense arrangement. From 1.5 to 2.3 the orientation is edge on. From 2.5 to 3.3 the arrangement is flat open. With a ratio from 4 to 11, the carbonate arrangement is flat-lacunar.^[2]

The simplest formula is LnCO₃F, where Ln has a 3+ charge.

For monocations there is A₃CO₃F, where A is a large ion such as K, Rb or Tl.^[2]

For M =alkali metal, and Ln = lanthanide: MLnCO₃F₂ 1:1:1:2; M₃Ln(CO₃)₂F₂ 3:1:2:2; M₂Ln(CO₃)₂F 2:1:2:1; M₄Ln(CO₃)₂F₃·H₂O 4:1:2:3; M₄Ln₂(CO₃)₃F₄ 2:3:3:4.^[2] M₂Ln(CO₃)F₂ 2:1:1:3.

For B =alkaline earth and Ln = lanthanide (a triple-charged ion) BLn(CO₃)₂F 1:1:2:1; BLn₂(CO₃)₃F₂ 1:2:3:2 B₂Ln₃(CO₃)₅F₃ 2:3:5:3; B₂Ln(CO₃)₂F₃ 2:1:2:3; B₂Ln(CO₃)F₅ 2:1:1:5 B₂Ln(CO₃)₃F 2:1:3:1; B₃Ln(CO₃)F₇ 3:1:1:7; B₃Ln₂(CO₃)₅F₂ 3:2:5:2.^[2]

For alkali with dication combinations: MB: MBCO₃F MB₃(CO₃)₂F₃·H₂O.^[2]

For dications A and B there is ABCO₃F₂ with a degenerate case of A = B.^[2]

KPb₂(CO₃)₂F is layered. Each layer is like a sandwich, with lead and carbonate in the outer sublayers, and potassium and fluoride in the inner layer. K_2.70Pb_5.15(CO₃)₅F₃ extends this structure with some of the layers also being a double-decker sandwich of carbonate, fluoride, carbonate, fluoride, carbonate.^[3]

In the rare-earth fluorocarbonates the environment for the rare-earth atoms is 9-coordinated. Six oxygen atoms from carbonate are at the apices of a trigonal prism, and fluoride ions cap the rectangular faces of the prism.^[4]

Carbonate fluoride compounds can be formed by a variety of related methods involving heating the precursor ingredients with or without water. Thallous fluoride carbonate was made simply by evaporating a fluoride thallium solution in ethanol and water in air. It absorbed sufficient carbon dioxide to yield the product. Most other carbonate fluorides are very insoluble and need high-temperature water to crystallise from. Supercritical water heated between 350 and 750 °C under pressures around 200 bars can be used. A sealed platinum tube can withstand the heat and pressure. Crystallisation takes about a day. With subcritical water around 200 °C, crystallisation takes about 2 days. This can happen in a teflon-coated pressure autoclave. The starting ingredients can be rare-earth fluorides and alkali carbonates. The high pressure is needed to keep the water liquid and the carbon dioxide under control, otherwise it would escape. If the fluoride levels are low, hydroxide can substitute for the fluoride. Solid-state reactions require even higher temperatures.^[2]

Bastnäsite along with lukechangite (andpetersenite) can be precipitated from a mixed solution of CeCl₃, NaF, and NaOH with carbon dioxide.^[5] Another way to make the simple rare-earth fluorocarbonates is to precipitate a rare-earth carbonate from a nitrate solution withammonium bicarbonate and then add fluoride ions withhydrofluoric acid (HF).^[6]

Pb₂(CO₃)F₂ can be made by boiling a water solution oflead nitrate,sodium fluoride andpotassium carbonate in a 2:2:1 molar ratio.^[7]

The visible spectrum of fluorocarbonates is determined mainly by the cations contained. Different structures only have slight effect on the absorption spectrum of rare-earth elements.^[4] The visible spectrum of the rare-earth fluorocarbonates is almost entirely due to narrow absorption bands fromneodymium.^[4] In thenear infrared around 1000 nm there are some absorption lines due tosamarium and around 1547 nm are some absorption features due topraseodymium. Deeper into the infrared, bastnäsite has carbonate absorption lines at 2243, 2312 and 2324 nm. Parisite only has a very weak carbonate absorption at 2324 nm, and synchysite absorbs at 2337 nm.^[4]

The infrared spectrum due to vibration of carbon–oxygen bonds in carbonate is affected by how many kinds of position there are for the carbonate ions.^[4]

An important chemical reaction used to prepare rare-earth elements from their ores, is to roast concentrated rare-earth fluorocarbonates with sulfuric acid at about 200 °C. This is then leached with water. This process liberates carbon dioxide and hydrofluoric acid and yields rare-earth sulfates:

2 LnCO₃F + 3 H₂SO₄ → Ln₂(SO₄)₃ + 2 HF + 2 H₂O + 2 CO₂.

Subsequent processing precipitates a double sulfate with sodium sulfate at about 50 °C. The aim is to separate out the rare-earth elements from calcium, aluminium, iron and thorium.^[8]

At high enough temperatures the carbonate fluorides lose carbon dioxide, e.g.

KCu(CO₃)F → CuO + KF + CO₂

at 340 °C.^[2]

The processing of bastnäsite is important, as it is the most commonly minedcerium mineral. When heated in air or oxygen at over 500 °C, bastnäsite oxidises and loses volatiles to formceria (CeO₂). Lukechangite also oxidises to ceria and sodium fluoride (NaF). Ce₇O₁₂ results when heated to over 1000 °C.^[5]

2 Ce(CO₃F) + O₂ → 2 CeO₂ + 2 CO₂ + F₂^[5]

Na₃Ce₂(CO₃F)₄F +⁠1/2⁠ O₂ → 2 CeO₂ + 3 CO₂ + NaF + Na₂CO₃^[5]

At 1300 °C Na₂CO₃ loses CO₂, and between 1300 and 1600 °C NaF and Na₂O boil off.^[5]

When other rare-earth carbonate fluorides are heated, they lose carbon dioxide and form an oxyfluoride:

LaCO₃F → LaOF + CO₂^[9]

In some rare-earth extraction processes, the roasted ore is then extracted withhydrochloric acid to dissolve rare earths apart from cerium. Cerium is dissolved if the pH is under 0, and thorium is dissolved if it is under 2.^[10]

KCdCO₃F when heated yieldscadmium oxide (CdO) andpotassium fluoride (KF).^[11]

When lanthanum fluorocarbonate is heated in ahydrogen sulfide, or carbon disulfide vapour around 500 °C, lanthanum fluorosulfide forms:

LaCO₃F +⁠1/2⁠ CO₂ → LaSF + 1.5 CO₂^[12]

Note that this also works for other lanthanides apart from cerium.

When lanthanum carbonate fluoride is heated at 1000 °C with alumina,lanthanum aluminate is produced:^[13]

LaCO₃F + 2 Al₂O₃ → LaAlO₃ + CO₂ + equiv AlOF

Within the hot part of the Earth's crust, rare-earth fluorocarbonates should react withapatite to formmonazite.^[14]

Somerare-earth fluorocarbonate minerals exist. They make up most of the economic ores for light rare-earth elements (LREE). These probably result from hydrothermal liquids from granite that contained fluoride.^[15] Rare-earth fluorocarbonate minerals can form inbauxite oncarbonate rocks, as rare-earth fluoride complexes react with carbonate.^[16] Carbonate fluoride compounds of rare-earth elements also occur incarbonatites.^[17]

These are non-linear optical crystals in the AMCO₃F familyKSrCO₃FKCaCO₃FRbSrCO₃FKCdCO₃FCsPbCO₃FRbPbCO₃FRbMgCO₃FKMgCO₃FRbCdCO₃FCsSrCO₃FRbCaCO₃FKZnCO₃FCsCaCO₃FRbZnCO₃F^[30]

• Fluorides
• Carbonates
• Mixed anion compounds
• Double salts

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