Jump to content

Metal ammine complex

From Wikipedia, the free encyclopedia
(Redirected from Cobalt ammine complexes)
Ball-and-stick model of the tetraamminediaquacopper(II) cation, [Cu(NH3)4(H2O)2]2+

In coordination chemistry, metal ammine complexes are metal complexes containing at least one ammonia (NH3) ligand. "Ammine" is spelled this way for historical reasons;[1] in contrast, alkyl or aryl bearing ligands are spelt with a single "m". Almost all metal ions bind ammonia as a ligand, but the most prevalent examples of ammine complexes are for Cr(III), Co(III), Ni(II), Cu(II) as well as several platinum group metals.[2]

History

[edit]
Structural representations used by Alfred Werner (right) and Sophus Mads Jørgensen for one isomer of the dichloride salt of the complex [Pt(NH3)2(pyridine)2]2+.[3]

Ammine complexes played a major role in the development of coordination chemistry, specifically determination of the stereochemistry and structure. They are easily prepared, and the metal-nitrogen ratio can be determined by elemental analysis. Through studies mainly on the ammine complexes, Alfred Werner developed his Nobel Prize-winning concept of the structure of coordination compounds (see Figure).[4][2]

Originally salts of [Co(NH3)6]3+ were described as the luteo (Latin: yellow) complex of cobalt. This name has been discarded as modern chemistry considers color less important than molecular structure. Other metal ammine complexes also were labeled according to their color, such as purpureo (Latin: purple) for a cobalt pentammine complex, and praseo (Greek: green) and violeo (Latin: violet) for two isomeric tetrammine complexes.[5]

One of the first ammine complexes to be described was Magnus' green salt, which consists of the platinum tetrammine complex [Pt(NH3)4]2+.[6]

Structure and bonding

[edit]

Ammonia is a pure σ-donor, in the middle of the spectrochemical series, and shows intermediate hard–soft behaviour (see also ECW model). Its relative donor strength toward a series of acids, versus other Lewis bases, can be illustrated by C-B plots.[7][8]

An ammine ligand bound to a metal ion is markedly more acidic than a free ammonia molecule, although deprotonation in aqueous solution is still rare. One example is the reaction of mercury(II) chloride with ammonia (Calomel reaction) where the resulting mercuric amidochloride is highly insoluble.

HgCl2 + 2 NH3 → HgCl(NH2) + [NH4]Cl

Ammonia is a Lewis base and a "pure" sigma donor. It is also compact such that steric effects are negligible. These factors simplify interpretation of structural and spectroscopic results.The Co–N distances in complexes [M(NH3)6]n+ have been examined closely by X-ray crystallography.[9]

M–N distances for [M(NH3)6]n+
M n+ M–N distance (Å) d-electron configuration comment
Co 3+ 1.936 t2g6 eg0 low-spin trications are small
Co 2+ 2.114 t2g5 eg2 population of eg orbital and lower positive charge
Ru 3+ 2.104 t2g5 eg0 low spin trication, but Ru is intrinsically larger than Co
Ru 2+ 2.144 t2g6 eg0 low spin dication

Examples

[edit]

Homoleptic poly(ammine) complexes are known for many of the transition metals. Most often, they have the formula [M(NH3)6]n+ where n = 2, 3, and even 4 (M = Pt).[10]

Platinum group metals

[edit]

Platinum group metals form diverse ammine complexes. Pentaamine(dinitrogen)ruthenium(II) and the Creutz–Taube complex are well-studied examples of historic significance. The complex cis-[PtCl2(NH3)2], under the name Cisplatin, is an important anticancer drug. Pentamminerhodium chloride ([RhCl(NH3)5]2+) is an intermediate in the purification of rhodium from its ores.

Cobalt(III) and chromium(III)

[edit]

The ammines of chromium(III) and cobalt(III) are of historic significance. Both families of ammines are relatively inert kinetically, which allows the separation of isomers.[11] For example, tetraamminedichlorochromium(III) chloride, [Cr(NH3)4Cl2]Cl, has two forms - the cis isomer is violet, while the trans isomer is green. The trichloride of the hexaammine (hexamminecobalt(III) chloride, [Co(NH3)6]Cl3) exists as only a single isomer. "Reinecke's salt" with the formula [NH4]+[Cr(NCS)4(NH3)2]−·H2O was first reported in 1863.[12]

Nickel(II), zinc(II), copper(II)

[edit]
Sample of chloropentamminecobalt chloride [CoCl(NH3)5]Cl2, illustrating the vibrant colors typical of transition metal ammine complexes.

Zinc(II) forms a colorless tetraammine with the formula [Zn(NH3)4]2+.[13] Like most zinc complexes, it has a tetrahedral structure. Hexaamminenickel is violet, and the copper(II) complex is deep blue. The latter is characteristic of the presence of copper(II) in qualitative inorganic analysis.

Copper(I), silver(I), and gold(I)

[edit]

Copper(I) forms only labile complexes with ammonia, including the trigonal planar [Cu(NH3)3]+.[14] Silver gives the diammine complex [Ag(NH3)2]+ with linear coordination geometry.[15] It is this complex that forms when otherwise rather insoluble silver chloride dissolves in aqueous ammonia. The same complex is the active ingredient in Tollens' reagent. Gold(I) chloride reacts with ammonia to form [Au(NH3)2]+.[16]

Reactions

[edit]

Ligand exchange and redox reactions

[edit]

Since ammonia is a stronger ligand in the spectrochemical series than water, metal ammine complexes are stabilized relative to the corresponding aquo complexes. For similar reasons, metal ammine complexes are less strongly oxidizing than are the corresponding aquo complexes. The latter property is illustrated by the stability of [Co(NH3)6]3+ in aqueous solution and the nonexistence of [Co(H2O)6]3+ (which would oxidize water).

Acid-base reactions

[edit]

Once complexed to a metal ion, ammonia is no longer basic. This property is illustrated by the stability of some metal ammine complexes in strong acid solutions. When the M–NH3 bond is weak, the ammine ligand dissociates and protonation ensues. The behavior is illustrated by the respective non-reaction and reaction with [Co(NH3)6]3+ and [Ni(NH3)6]2+ toward aqueous acids.

The ammine ligands are more acidic than is ammonia (pKa ~ 33). For highly cationic complexes such as [Pt(NH3)6]4+, the conjugate base can be obtained. The deprotonation of cobalt(III) ammine-halide complexes, e.g. [CoCl(NH3)5]2+ labilises the Co–Cl bond, according to the Sn1CB mechanism.

Oxidation of ammonia

[edit]

Deprotonation can be combined with oxidation, allowing the conversion of ammine complexes into nitrosyl complexes:[17]

H2O + [Ru(terpy)(bipy)(NH3)]+ → [Ru(terpy)(bipy)(NO)]2+ + 5 H+ + 6 e−

H-atom transfer

[edit]

In some ammine complexes, the N–H bond is weak. Thus one tungsten ammine complex evolve hydrogen:[17]

2 W(terpy)(PMe2Ph)2(NH3)]+ → 2 [W(terpy)(PMe2Ph)2(NH2)]+ + H2

This behavior is relevant to the use of metal-ammine complexes as catalysts for the oxidation of ammonia.[18]

Applications

[edit]

Metal ammine complexes find many uses. Cisplatin (cis-[PtCl2(NH3)2]) is a drug used in treating cancer.[19] Many other amine complexes of the platinum group metals have been evaluated for this application.

In the separation of the individual platinum metals from their ore, several schemes rely on the precipitation of [RhCl(NH3)5]Cl2. In some separation schemes, palladium is purified by manipulating equilibria involving [Pd(NH3)4]Cl2, [PdCl2(NH3)2], and [Pt(NH3)4][PtCl4] (Magnus's green salt).

In the processing of cellulose, the copper ammine complex known as Schweizer's reagent ([Cu(NH3)4(H2O)2](OH)2) is sometimes used to solubilise the polymer. Schweizer's reagent is prepared by treating an aqueous solutions of copper(II) ions with ammonia. Initially, the light blue hydroxide precipitates only to redissolve upon addition of more ammonia:

[Cu(H2O)6]2+ + 2 OH− → Cu(OH)2 + 6 H2O
Cu(OH)2 + 4 NH3 + 2 H2O → [Cu(NH3)4(H2O)2]2+ + 2 OH−

Silver diammine fluoride ([Ag(NH3)2]F) is a topical medicament (drug) used to treat and prevent dental caries (cavities) and relieve dentinal hypersensitivity.[20]

See also

[edit]

References

[edit]
  1. ^ "Definition of AMMINE". www.merriam-webster.com. Retrieved 2024-06-27.
  2. ^ a b A. von Zelewsky "Stereochemistry of Coordination Compounds" John Wiley: Chichester, 1995. ISBN 0-471-95599-X.
  3. ^ Alfred Werner "Beitrag zur Konstitution anorganischer Verbindungen" Zeitschrift für anorganische Chemie 1893, Volume 3, pages 267–330.doi:10.1002/zaac.18930030136
  4. ^ "Werner Centennial" George B. Kauffman, Ed. Adv. Chem. Ser., 1967, Volume 62. ISBN 978-0-8412-0063-0
  5. ^ Huheey, James E. (1983). Inorganic Chemistry (3rd ed.). p. 360.
  6. ^ Atoji, M.; Richardson, J. W.; Rundle, R. E. (1957). "On the Crystal Structures of the Magnus Salts, Pt(NH3)4PtCl4". J. Am. Chem. Soc. 79 (12): 3017–3020. doi:10.1021/ja01569a009.
  7. ^ Laurence, C. and Gal, J-F. Lewis Basicity and Affinity Scales, Data and Measurement, (Wiley 2010) pp 50–51 ISBN 978-0-470-74957-9
  8. ^ Cramer, R. E.; Bopp, T. T. (1977). "Graphical display of the enthalpies of adduct formation for Lewis acids and bases". Journal of Chemical Education. 54: 612–613. doi:10.1021/ed054p612. The plots shown in this paper used older parameters. Improved E&C parameters are listed in ECW model.
  9. ^ Hair, Neil J.; Beattie, James K. (1977). "Structure of Hexaaquairon(III) Nitrate Trihydrate. Comparison of Iron(II) and Iron(III) Bond Lengths in High-Spin Octahedral Environments". Inorganic Chemistry. 16 (2): 245–250. doi:10.1021/ic50168a006.
  10. ^ Eßmann, Ralf; Kreiner, Guido; Niemann, Anke; Rechenbach, Dirk; Schmieding, Axel; Sichla, Thomas; Zachwieja, Uwe; Jacobs, Herbert (1996). "Isotype Strukturen einiger Hexaamminmetall(II)-halogenide von 3d-Metallen: [V(NH3)6]I2, [Cr(NH3)6]I2, [Mn(NH3)6]Cl2, [Fe(NH3)6]Cl2, [Fe(NH3)6]Br2, [Co(NH3)6]Br2, und [Ni(NH3)6]Cl2". Zeitschrift für anorganische und allgemeine Chemie. 622 (7): 1161–1166. doi:10.1002/zaac.19966220709.
  11. ^ Basolo, F.; Pearson, R. G. "Mechanisms of Inorganic Reactions." John Wiley and Son: New York: 1967. ISBN 0-471-05545-X
  12. ^ Reinecke, A. "Über Rhodanchromammonium-Verbindungen" Annalen der Chemie und Pharmacie, volume 126, pages 113-118 (1863). doi:10.1002/jlac.18631260116.
  13. ^ Essmann, R. (1995). "Influence of coordination on N-H...X- hydrogen bonds. Part 1. [Zn(NH3)4]Br2 and [Zn(NH3)4]I2". Journal of Molecular Structure. 356 (3): 201–6. Bibcode:1995JMoSt.356..201E. doi:10.1016/0022-2860(95)08957-W.
  14. ^ Nilsson, Kersti B.; Persson, Ingmar (2004). "The coordination chemistry of copper(I) in liquid ammonia, trialkyl and triphenyl phosphite, and tri-n-butylphosphine solution". Dalton Transactions (9): 1312–1319. doi:10.1039/B400888J. PMID 15252623.
  15. ^ Nilsson, K. B.; Persson, I.; Kessler, V. G. (2006). "Coordination Chemistry of the Solvated AgI and AuI Ions in Liquid and Aqueous Ammonia, Trialkyl and Triphenyl Phosphite, and Tri-n-butylphosphine Solutions". Inorganic Chemistry. 45 (17): 6912–6921. doi:10.1021/ic060175v. PMID 16903749.
  16. ^ Scherf, L. M.; Baer, S. A.; Kraus, F.; Bawaked, S. M.; Schmidbaur, H. (2013). "Implications of the Crystal Structure of the Ammonia Solvate [Au(NH3)2]Cl·4NH3". Inorganic Chemistry. 52 (4): 2157–2161. doi:10.1021/ic302550q. PMID 23379897.
  17. ^ a b Dunn, Peter L.; Cook, Brian J.; Johnson, Samantha I.; Appel, Aaron M.; Bullock, R. Morris (2020). "Oxidation of Ammonia with Molecular Complexes". Journal of the American Chemical Society. 142 (42): 17845–17858. doi:10.1021/jacs.0c08269. OSTI 1706682. PMID 32977718. S2CID 221938378.
  18. ^ Zott, Michael D.; Peters, Jonas C. (2023). "Improving Molecular Iron Ammonia Oxidation Electrocatalysts via Substituent Effects that Modulate Standard Potential and Stability". ACS Catalysis. 13 (21): 14052–14057. doi:10.1021/acscatal.3c03772. S2CID 264338937.
  19. ^ S. J. Lippard, J. M. Berg "Principles of Bioinorganic Chemistry" University Science Books: Mill Valley, CA; 1994. ISBN 0-935702-73-3.
  20. ^ Rosenblatt, A.; Stamford, T. C. M.; Niederman, R. (2009). "Silver diamine fluoride: a caries "silver-fluoride bullet"". Journal of Dental Research. 88 (2): 116–125. doi:10.1177/0022034508329406. PMID 19278981. S2CID 30730306.
Retrieved from "https://en.wikipedia.org/w/index.php?title=Metal_ammine_complex&oldid=1235468255"