Metal ammine complex

Ball-and-stick model of the tetraamminediaquacopper(II) cation, [Cu(NH

4(H
2O)
2]2+

In coordination chemistry, metal ammine complexes are metal complexes containing at least one ammonia (NH
3) ligand. "Ammine" is spelled this way for historical reasons[citation needed]; 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.[1]

## Contents

* 1 History
* 2 Structure and bonding
* 3 Examples
* 3.1 Platinum group metals
* 3.2 Cobalt(III) and chromium(III)
* 3.3 Nickel(II), zinc(II), copper(II)
* 3.4 Copper(I), silver(I), and gold(I)
* 4 Reactions
* 4.1 Ligand exchange and redox reactions
* 4.2 Acid-base reactions
* 4.3 Oxidation of ammonia
* 4.4 H-atom transfer
* 5 Applications
* 6 See also
* 7 References

## History

Structural representations used by Alfred Werner (right) and Sophus Mads Jørgensen for one isomer of the dichloride salt of the complex [Pt(NH

2(pyridine)
2]2+.[2]

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).[3][1]

One of the first ammine complexes to be described was Magnus' green salt, which consists of the platinum tetrammine complex [Pt(NH

4]2+.[4]

## Structure and bonding

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.[5][6]

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.

HgCl
2 \+ 2 NH
3 → HgCl(NH
2) + [NH
4]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(NH

6]n+
have been examined closely by X-ray crystallography.[7]

M–N distances for [M(NH

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

Homoleptic poly(ammine) complexes are known for many of the transition metals. Most often, they have the formula [M(NH

6]n+
where n = 2, 3, and even 4 (M = Pt).[8]

### Platinum group metals

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-[PtCl
2(NH

2], under the name Cisplatin, is an important anticancer drug. Pentamminerhodium chloride ([RhCl(NH

5]2+) is an intermediate in the purification of rhodium from its ores.

* Metal-Ammine Complexes
* Carboplatin, a widely used anticancer drug.

* Pentamminerhodium chloride, the dichloride salt of a pentammine halide complex.

* Pentaamine(dinitrogen)ruthenium(II), the first metal dinitrogen complex.

* Hexamminecobalt(III) chloride, the trichloride salt of the hexammine complex [Co(NH

6]3+. It is famously stable in concentrated hydrochloric acid.

* Reinecke's salt features a very stable anionic diamine complex of Cr(III), which is used as a counteranion.

### Cobalt(III) and chromium(III)

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.[9] For example, tetraamminedichlorochromium(III) chloride, [Cr(NH

4Cl
2]Cl, has two forms - the cis isomer is violet, while the trans isomer is green. The trichloride of the hexaammine (hexamminecobalt(III) chloride, [Co(NH

6]Cl
3) exists as only a single isomer. "Reinecke's salt" with the formula [NH
4]+
[Cr(NCS)
4(NH

2]−
* H
2O was first reported in 1863.[10]

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

Sample of chloropentamminecobalt chloride [CoCl(NH

5]Cl
2, illustrating the vibrant colors typical of transition metal ammine complexes.

Zinc(II) forms a colorless tetraammine with the formula [Zn(NH

4]2+.[11] 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)

Copper(I) forms only labile complexes with ammonia, including the trigonal planar [Cu(NH3)3]+.[12] Silver gives the diammine complex [Ag(NH3)2]+ with linear coordination geometry.[13] 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(NH

2]+
.[14]

## Reactions

### Ligand exchange and redox reactions

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(NH

6]3+ in aqueous solution and the nonexistence of [Co(H
2O)
6]3+ (which would oxidize water).

### Acid-base reactions

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–NH
3 bond is weak, the ammine ligand dissociates and protonation ensues. The behavior is illustrated by the respective non-reaction and reaction with [Co(NH

6]3+ and [Ni(NH

6]2+ toward aqueous acids.

The ammine ligands are more acidic than is ammonia (pKa ~ 33). For highly cationic complexes such as [Pt(NH

6]4+, the conjugate base can be obtained. The deprotonation of cobalt(III) ammine-halide complexes, e.g. [CoCl(NH

5]2+ labilises the Co–Cl bond, according to the Sn1CB mechanism.

### Oxidation of ammonia

Deprotonation can be combined with oxidation, allowing the conversion of ammine complexes into nitrosyl complexes:[15]

2O + [Ru(terpy)([[Chemistry:

2′−
Bipyridine|bipy]])(NH
3)]+
→ [Ru(terpy)(bipy)(NO)]2+ \+ 5 H+
\+ 6 e−

### H-atom transfer

In some ammine complexes, the N–H bond is weak. Thus one tungsten ammine complex evolve hydrogen:[15]

2 W(terpy)(PMe
2Ph)
2(NH
3)]+
→ 2 [W(terpy)(PMe
2Ph)
2(NH
2)]+
\+ H

This behavior is relevant to the use of metal-ammine complexes as catalysts for the oxidation of ammonia.[16]

## Applications

Metal ammine complexes find many uses. Cisplatin (cis-[PtCl
2(NH

2]) is a drug used in treating cancer.[17] 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(NH

5]Cl
2. In some separation schemes, palladium is purified by manipulating equilibria involving [Pd(NH

4]Cl
2, [PdCl
2(NH

2], and [Pt(NH

4][PtCl
4] (Magnus's green salt).

In the processing of cellulose, the copper ammine complex known as Schweizer's reagent ([Cu(NH

4(H
2O)
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(H
2O)
6]2+ \+ 2 OH−
→ Cu(OH)
2 \+ 6 H

Cu(OH)
2 \+ 4 NH
3 \+ 2 H
2O → [Cu(NH

4(H
2O)
2]2+ \+ 2 OH−

Silver diammine fluoride ([Ag(NH

2]F) is a topical medicament (drug) used to treat and prevent dental caries (cavities) and relieve dentinal hypersensitivity.[18]

## See also

* Ligand field theory

## References

1. ↑ 1.0 1.1 A. von Zelewsky "Stereochemistry of Coordination Compounds" John Wiley: Chichester, 1995. ISBN:0-471-95599-X.
2. ↑ Alfred Werner "Beitrag zur Konstitution anorganischer Verbindungen" Zeitschrift für anorganische Chemie 1893, Volume 3, pages 267–330.doi:10.1002/zaac.18930030136
3. ↑ "Werner Centennial" George B. Kauffman, Ed. Adv. Chem. Ser., 1967, Volume 62. ISBN:978-0-8412-0063-0
4. ↑ 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.
5. ↑ 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
6. ↑ 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.
7. ↑ 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.
8. ↑ 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.
9. ↑ Basolo, F.; Pearson, R. G. "Mechanisms of Inorganic Reactions." John Wiley and Son: New York: 1967. ISBN:0-471-05545-X
10. ↑ Reinecke, A. "Über Rhodanchromammonium-Verbindungen" Annalen der Chemie und Pharmacie, volume 126, pages 113-118 (1863). doi: 10.1002/jlac.18631260116.
11. ↑ 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. doi:10.1016/0022-2860(95)08957-W. Bibcode: 1995JMoSt.356..201E.
12. ↑ 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.
13. ↑ 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.
14. ↑ 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.
15. ↑ 15.0 15.1 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. PMID 32977718.
16. ↑ 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.
17. ↑ S. J. Lippard, J. M. Berg "Principles of Bioinorganic Chemistry" University Science Books: Mill Valley, CA; 1994. ISBN:0-935702-73-3.
18. ↑ 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.

* v
* t
* e

Coordination complexes

Inorganic ligands

* H2O
* NH3
* N3-
* N2
* O2
* CN−
* CO
* CO2
* F−
* Cl−
* Br−
* I−

Organic ligands
* C5H5
* acac
* CH2＝CH-CH2−
* PR3
* (Me3Si)2N-
* S2C2R22-
* RCN
* =C
* ≡C

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