Cuprate

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Cuprates are a class of compounds that contain copper (Cu) atom(s) in an anion. They can be broadly categorized into two main types:

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1. Inorganic cuprates: These compounds have a general formula of XYCumOn. Some of them are non-stoichiometric. Many of these compounds are known for their superconducting properties.[ citation needed ] An example of an inorganic cuprate is the tetrachloridocuprate(II) or tetrachlorocuprate(II) ([ Cu Cl 4]2−), an anionic coordination complex that features a copper atom in an oxidation state of +2, surrounded by four chloride ions.

2. Organic cuprates: These are organocopper compounds, some of which having a general formula of [CuR2]−, where copper is in an oxidation state of +1, where at least one of the R groups can be any organic group. These compounds, characterized by copper bonded to organic groups, are frequently used in organic synthesis due to their reactivity.[ citation needed ] An example of an organic cuprate is dimethylcuprate(I) anion [Cu(CH3)2]−.

One of the most studied cuprates is Y Ba 2 Cu 3 O 7, a high-temperature superconducting material. This oxide cuprate has been the subject of extensive research due to its ability to conduct electricity without resistance at relatively high temperatures.[ citation needed ]

The term 'cuprate' originates from 'cuprum', the Latin word for copper. It is primarily used in the context of oxide materials, anionic coordination complexes, and anionic organocopper compounds, reflecting the diverse roles of copper in chemistry. The term is mainly used in three contexts: oxide materials, anionic coordination complexes, and anionic organocopper compounds.[ citation needed ]

Oxide cuprates

One of the simplest oxide-based cuprates is potassium cuprate(III) KCuO2. This species can be viewed as the K + salt of the polyanion [CuO−2]n. As such the material is classified as an oxide cuprate. This dark blue diamagnetic solid is produced by heating potassium peroxide and copper(II) oxide in an atmosphere of oxygen: [1]

K2O2 + 2 CuO → 2 KCuO2

Other cuprates(III) of alkali metals are known; in addition, the structures of KCuO2 (potassium cuprate(III)), RbCuO2 (rubidium cuprate(III)) and CsCuO2 (caesium cuprate(III)) have been determined as well. [2]

KCuO2 was discovered first in 1952 by V. K. Wahl and W. Klemm, they synthesized this compound by heating copper(II) oxide and potassium superoxide in an atmosphere of oxygen. [3]

2 KO2 + 2 CuO → KCuO2 + O2

It can also be synthesized by heating potassium superoxide and copper powder: [4]

KO2 + Cu → KCuO2

KCuO2 reacts with the air fairly slowly. It starts to decompose at 760 K (487 °C; 908 °F) and its color changes from blue to pale green at 975 K (702 °C; 1,295 °F). Its melting point is 1,025 K (752 °C; 1,385 °F). [3] [4]

RbCuO2 (blue-black) and CsCuO2 (black) can be prepared by reaction of rubidium oxide and caesium oxide with copper(II) oxide powders, at 675 K (402 °C; 755 °F) and 655 K (382 °C; 719 °F) in oxygen atmosphere, respectively. Either of them reacts with the air fast, unlike KCuO2. [4]

In fact, KCuO2 is a non-stoichiometric compound, so the more exact formula is KCuOx and x is very close to 2. This causes the formation of defects in the crystal structure, and this leads to the tendency of this compound to be reduced. [4]

Sodium cuprate(III) NaCuO2 can be produced by using hypochlorites or hypobromites to oxidize copper hydroxide under alkaline and low temperature conditions. [5]

2 NaOH + CuSO4 → Cu(OH)2↓
Cu(OH)2 + 2 NaOH + NaClO → 2 NaCuO2 + NaCl + H2O

Cuprates(III) are not stable in water, and they can oxidize water as well. [5]

4 CuO−2 + 2 H2O → 4 CuO + O2↑ + 4 OH−

Sodium cuprate(III) is reddish-brown, but turns black gradually as it decomposes to copper(II) oxide. [5] In order to prevent decomposition, it must be prepared at low temperature in the absence of light.[ citation needed ]

Coordination complexes

Copper forms many anionic coordination complexes with negatively charged ligands such as cyanide, hydroxide, and halides, as well as alkyls and aryls.

Copper(I)

Cuprates containing copper(I) tend to be colorless, reflecting their d10 configuration. Structures range from linear 2-coordinate, trigonal planar, and tetrahedral molecular geometry. Examples include linear [CuCl 2]− and trigonal planar [CuCl3]2−. [6] Cyanide gives analogous complexes but also the trianionic tetracyanocuprate(I), [Cu(CN)4]3−. [7] Dicyanocuprate(I), [Cu(CN)2]−, exists in both molecular or polymeric motifs, depending on the countercation. [8]

Copper(II)

Caesium salt of hexafluorocuprate(IV) CCs2CuF6.svg
Caesium salt of hexafluorocuprate(IV)

Cuprates containing copper(II) include trichlorocuprate(II), [CuCl3]−, which is dimeric, and square-planar tetrachlorocuprate(II), [CuCl4]2−, and pentachlorocuprate(II), [CuCl5]3−. [9] [10] 3-Coordinate chlorocuprate(II) complexes are rare. [11]

Tetrachlorocuprate(II) complexes tend to adopt flattened tetrahedral geometry with orange colors. [12] [13] [14] [15]

Sodium tetrahydroxycuprate(II) (Na2[Cu(OH)4]) is an example of a homoleptic (all ligands being the same) hydroxide complex. [16]

Cu(OH)2 + 2 NaOH → Na2[Cu(OH)4]

Copper(III) and copper(IV)

Hexafluorocuprate(III) [CuF6]3− and hexafluorocuprate(IV) [CuF6]2− are rare examples of copper(III) and copper(IV) complexes. They are strong oxidizing agents.

Organic cuprates

Structure of lithium diphenylcuprate(I) etherate, 2[Ph2Cu] Li*2OEt2. Lithium-diphenylcuprate-etherate-dimer-from-xtal-2D-skeletal.png
Structure of lithium diphenylcuprate(I) etherate, 2[Ph2Cu] Li·2OEt 2.

Cuprates have a role in organic synthesis. They are invariably Cu(I), although Cu(II) or even Cu(III) intermediates are invoked in some chemical reactions. Organic cuprates often have the idealized formulas [CuR2]− and [CuR3]2−, both of which contain copper in an oxidation state of +1, where R is an alkyl or aryl. These reagents find use as nucleophilic alkylating reagents. [18]

See also

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References

  1. ↑ G. Brauer, ed. (1963). "Potassium Cuprate (III)". Handbook of Preparative Inorganic Chemistry. Vol. 2 (2nd ed.). NY: Academic Press. p. 1015.
  2. ↑ Hestermann, Klaus; Hoppe, Rudolf (1969). "Die Kristallstruktur von KCuO2, RbCuO2 und CsCuO2". Zeitschrift für Anorganische und Allgemeine Chemie. 270 (1–4): 69–75. doi:10.1002/zaac.19693670506.
  3. 1 2 Wahl, Von Kurt; Klemm, Wilhelm (1952). "Über Kaliumcuprat(III)". Zeitschrift für Anorganische und Allgemeine Chemie. 270 (1–4): 69–75. doi:10.1002/zaac.19522700109 . Retrieved January 20, 2023.
  4. 1 2 3 4 Costa, Giorgio A.; Kaiser, Elena (1995). "Structural and thermal properties of the alkaline cuprate KCuO2". Thermochimica Acta. 269–270: 591–598. doi:10.1016/0040-6031(95)02575-8 . Retrieved January 20, 2023.
  5. 1 2 3 Magee, J. S.; Wood, R. H. (1965). "Studies of Sodium Cuprate(III) Stability". Canadian Journal of Chemistry. 43 (5): 1234–1237. doi:10.1139/v65-164.
  6. ↑ Stricker, Marion; Linder, Thomas; Oelkers, Benjamin; Sundermeyer, Jörg (2010). "Cu(I)/(II) based catalytic ionic liquids, their metallo-laminate solid state structures and catalytic activities in oxidative methanol carbonylation". Green Chemistry. 12 (9): 1589. doi:10.1039/c003948a.
  7. ↑ Kroeker, Scott; Wasylishen, Roderick E. (1999). "A multinuclear magnetic resonance study of crystalline tripotassium tetracyanocuprate". Canadian Journal of Chemistry. 77 (11): 1962–1972. doi:10.1139/v99-181.
  8. ↑ Bowmaker, Graham A.; Hartl, Hans; Urban, Victoria (2000). "Crystal Structures and Vibrational Spectroscopy of [NBu4][Cu(CN)X] (X = Br, I) and [NBu4][Cu3(CN)4]·CH3CN". Inorganic Chemistry. 39 (20): 4548–4554. doi:10.1021/ic000399s.
  9. ↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN   978-0-08-037941-8.
  10. ↑ Willett, Roger D.; Butcher, Robert E.; Landee, Christopher P.; Twamley, Brendan (2006). "Two Halide Exchange in Copper(II) Halide Dimers: (4,4′-Bipyridinium)Cu2Cl6−xBRX". Polyhedron. 25 (10): 2093–2100. doi:10.1016/j.poly.2006.01.005.
  11. ↑ Hasselgren, Catrin; Jagner, Susan; Dance, Ian (2002). "Three-Coordinate [CuIIX3]− (X = Cl, Br), Trapped in a Molecular Crystal". Chemistry – A European Journal. 8 (6): 1269–1278. doi:10.1002/1521-3765(20020315)8:6<1269::AID-CHEM1269>3.0.CO;2-9. PMID   11921210.
  12. ↑ Mahoui, A.; Lapasset, J.; Moret, J.; Saint Grégoire, P. (1996). "Tetraethylammonium Tetramethylammonium Tetrachlorocuprate(II), [(C2H5)4N][(CH3)4N][CuCl4]". Acta Crystallographica Section C. 52 (11): 2674–2676. doi:10.1107/S0108270196009031.
  13. ↑ Guillermo Mínguez Espallargas; Lee Brammer; Jacco van de Streek; Kenneth Shankland; Alastair J. Florence; Harry Adams (2006). "Reversible Extrusion and Uptake of HCl Molecules by Crystalline Solids Involving Coordination Bond Cleavage and Formation". J. Am. Chem. Soc. 128 (30): 9584–9585. doi:10.1021/ja0625733. PMID   16866484.
  14. ↑ Kelley, A.; Nalla, S.; Bond, M. R. (2015). "The square-planar to flattened-tetrahedral CuX42− (X = Cl, Br) structural phase transition in 1,2,6-trimethylpyridinium salts". Acta Crystallogr. B . 71 (Pt 1): 48–60. doi:10.1107/S205252061402664X. PMID   25643715.
  15. ↑ Egon Wiberg; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic Chemistry. Academic Press. pp. 1252–1264. ISBN   0-12-352651-5.
  16. ↑ Brauer, G., ed. (1963). "Sodium Tetrahydroxocuprate(II)". Handbook of Preparative Inorganic Chemistry. Vol. 1 (2nd ed.). New York, NY: Academic Press. p. 1015.
  17. ↑ Lorenzen, Nis Peter; Weiss, Erwin (1990). "Synthesis and Structure of a Dimeric Lithium Diphenylcuprate:[{Li(OEt2)}(CuPh2)]2". Angewandte Chemie International Edition in English. 29 (3): 300. doi:10.1002/anie.199003001.
  18. ↑ Louis S. Hegedus (1999). Transition metals in the synthesis of complex organic molecules. University Science Books. pp. 61–65. ISBN   1-891389-04-1.
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