Complexes de coordination
English
Diviser
Vue d'ensemble
Source : Laboratoire du Dr Neal Abrams — SUNY College of Environmental Science and Forestry
Métaux de transition sont partout de suppléments vitaminiques pour galvanoplastie. Aussi, métaux de transition forment des pigments dans les nombreuses peintures et composer tous les minéraux. En règle générale, les métaux de transition se trouvent sous la forme cationique puisqu’ils facilement oxydent, ou perdent des électrons et sont entourés par les donneurs d’électrons appelés ligands. Ces ligands ne pas forme ionique ou covalent liaisons avec le centre métallique, plutôt qu’ils prennent sur un troisième type de liaison connue comme covalentes coordonnées. La coordonnée-covalente entre un ligand et d’un métal est dynamique, ce qui signifie que les ligands sont continuellement échanger et re-coordination autour du centre métallique. Les identités des fois le métal et le ligand dicte les ligands adhère préférentiellement sur l’autre. En outre, couleur et propriétés magnétiques sont également en raison des types de complexes qui sont forment. Les composés de coordination qui se forment sont analysés en utilisant une variété d’instruments et d’outils. Cette expérience examine pourquoi tant de complexes sont possibles et utilise une méthode de spectrochimiques (couleur ou chimiques) pour aider à identifier le type de complexe qui se forme.
Principles
Procédure
Applications and Summary
From pigments to people, transitional metals are found throughout fields of chemistry, biology, geology, and engineering. Understanding the behavior of transition metals under different chemical states can be as simple as monitoring color or magnetic behavior. Nearly every 3d (4th row) transition metal is vital to physiological function and, in all cases, these metals are bound by ligands to form coordination complexes. For example, iron is vital to oxygen transport in all vertebrates. Hemoglobin, a complex protein, contains four heme subunits with Fe2+ in the center of each. In hemoglobin, the Fe2+ is chelated by a tetradentate ring and a histidine residue, making it square pyramidal (five-sided). When oxygen is present, the subunits become octahedral. O2 is considered a strong-field ligand, which causes large d-orbital t2g-eg splitting, making it low-spin. Relatively high-energy light is required to promote an electron to the eg state, so blue light is absorbed making oxygenated (arterial) blood appear bright red. In contrast, deoxygenated (venous) blood has a smaller d-orbital splitting and lower energy light red light is absorbed, making deoxygenated blood appear dark, purplish-red. In the same respect, carbon monoxide, CO, is a strong-field ligand and will displace oxygen. It gives blood an even brighter red appearance due to strong-field splitting. The preferential binding for CO over O2 in blood is often fatal.
Another application of coordination chemistry is in paints and pigments. While many pigments are simple metal oxides, others like Prussian Blue and Phthalocyanine Blue are coordination complexes whose color arises from the splitting in d-orbitals (Figure 2). In Prussian Blue, iron is surrounded by six cyanide ligands, creating the high-spin iron (III) hexacyanoferrate complex, Fe(CN)63-. Another compound, Phthalocyanine Blue, is a square planar complex with a copper (II) ion in the center surrounded by a tetradentate phthalocyanine molecule.
Figure 2. Prussian Blue, an iron-centered coordination complex and Phthalocyanine Blue, a copper-centered coordination complex.
Coordination compounds have a metal ion center with surrounding ligands and a counterion to balance charge. The ligands can be monodentate or chelating with two-four attachment sites. Ligands are also categorized by the spectrochemical series, which classifies the relative strength of the ligands to split a metal's d-orbitals. Both color and magnetic properties are influenced by the metal and the ligands. Large d-orbital splitting requires large energies to promote electrons into the higher energy orbitals and absorbs high-energy light (short wavelength). These are low spin-complexes and have the maximum number of paired electrons. In contrast, a small d-orbital splitting is known as weak-field and absorbs low energy light as well as has the maximum number of unpaired electrons. The charge and identity of the metal ion as well as the bound ligands define both the observed color and magnetic properties in coordination compounds.
References
- Shakhashiri, B. Z.; G. E. Dirreen, G. E; Juergens, F. Color, Solubility, and Complex Ion Equilibria of Nickel (II) Species in Aqueous Solution. J. Chem. Ed. 52 (12), 900-901 (1980).
Transcription
Coordination complexes consist of a central metal atom or ion bound to some number of functional groups known as ligands.
Electrons are found in predictable locations around an atom’s nucleus, called orbitals. Most metals have a large number of accessible electrons compared to light main group elements such as nitrogen, oxygen, or carbon. Ligands interact with, or coordinate to, metals in complex ways facilitated by these many accessible electrons.
Ligands coordinate to metals in many different arrangements, or geometries, which can have a significant effect on the reactivity at the metal center. The orientations that ligands adopt are affected by the electronic nature of both the ligands and the metal.
This video will introduce the principles of metal complexes and ligands, demonstrate a procedure for exchanging ligands at a metal center, and introduce a few applications of metal complexes in chemistry and medicine.
Ligands range from simple ions such as chloride to complex molecules such as porphyrins. The overall charge of a metal complex depends on the net charges of the metal and each ligand. Metals are frequently cationic, or positive, and ligands are often neutral or anionic.
Ligands coordinate to metals through one or more donor atoms bound to the metal. The number of non-adjacent donor groups within a ligand is called denticity. A bidentate ligand occupies two coordination sites on a metal, so a complex with three bidentate ligands can adopt the same geometry as a complex with six monodentate ligands.
Ions or solvent molecules can interact with a coordination complex without directly interfacing with the metal, often acting as counter-ions. These can also be involved in reactions in which at least one ligand is replaced with another, or substituted.
In associative substitution, the new ligand coordinates to the metal, and then one of the original ligands leaves, or dissociates. In dissociative substitution, a ligand first dissociates from the metal, after which the new ligand coordinates. Ligands may also associate or dissociate without substitution, changing the number of donor atoms around the metal.
Metal complexes usually possess orbitals that are close enough in energy to allow electronic transitions between them. The energy gap between these orbitals is correlated with certain ligand properties. These properties are often defined in the “spectrochemical series of ligands”, which ranks them from ‘weak’ to ‘strong’, where stronger ligands are associated with a larger energy difference.
It is more favorable for electrons to be in orbitals with the lowest possible energy. These stabilized orbitals are found in systems with the widest energy gap. Thus, simple exchange reactions favor complexes with strong ligands.
Coordination complexes absorb photons corresponding to the energy needed for electronic transitions across energy gaps, often in the visible spectrum. The wavelength of the absorbed light is the complementary color of the observed color of the complex. Thus, the increased energy gap from exchanging a weaker ligand for a stronger one may change the color of the complex.
Now that you understand the principles of metal complexes, let’s go through a procedure for examining changes in orbital energies by a series of ligand exchange reactions.
To begin the procedure, obtain the appropriate ligand solutions and glassware. Then, prepare a solution of 1.84 g of solid nickel sulfate hexahydrate and 100 mL deionized water. The green hexaaquanickel cation will form in solution.
In a fume hood, begin stirring the hexaaquanickel solution using a stir bar and stir plate. Then, add 15 mL of 5 M aqueous ammonia and wait for the solution color to change to deep blue, indicating the formation of the hexaamminenickel cation.
Next, add 10 mL of 30% ethylenediamine. The solution color change to purple indicates that ethylenediamine has displaced the ammonia, forming the tris(ethylenediamine)nickel cation.
Then, add 200 mL of 1% dimethylglyoxime in ethanol to the same beaker. The solution color change from purple to a suspension of the red powder indicates the formation of the poorly-soluble bis(dimethylglyoximato)nickel complex.
Finally, add 30 mL of 1 M potassium cyanide solution. The dissolution of the red solid and the solution color change to yellow indicates that the cyano ligands have displaced the dimethylglyoximato ligands, forming the tetracyanonickelate anion.
The substitution reactions were all spontaneous, following the predictions of the spectrochemical series.
The energy needed to cause electronic transitions within these complexes is predicted by the series to be lowest for water and highest for cyanide.
The complementary colors associated with each solution are red, orange, yellow, green, and blue. The energy of visible light increases from red to blue, suggesting that the absorbed photons also increase in energy as ligand strength increases, which corresponds to a larger gap between orbital energy levels.
Metal complexes are used in a wide range of domains, from chemical synthesis, to the medical field.
Many metal complexes are used as catalysts or as reagents in stoichiometric quantities in organic synthesis. Development of new catalysts with various ligands and metal centers is ongoing, allowing access to new chemical compounds. Many of the mechanisms by which these reactions occur involve ligand exchange at the metal center. A small variation in ligands can have a large effect on the reactivity of a metal complex in organic synthesis. An understanding of relative ligand strength and the steric and electronic effects of ligands on the metal complex is therefore essential when designing new catalysts.
Metal complexes are often used in chemotherapy. Development of new anti-cancer drugs often involves evaluation of complexes similar to existing drugs, but using different ligands or metals. Here, titanium and vanadium complexes were found to show similar efficacies in preliminary evaluations to cisplatin, a platinum complex widely used. These compounds may interact with cancer cells in different ways from cisplatin because of the differences, and thus may be effective against different types of cancer cells.
Contrast agents are usually metal complexes that, when introduced to the body, interact with the water in nearby tissues to either enhance or diminish MRI imaging. The development of new contrast agents focuses on minimizing the toxicity posed while retaining the properties of an effective agent.
You’ve just watched JoVE’s introduction to coordination chemistry. You should now be familiar with the principles of coordination chemistry, a procedure for performing ligand exchange at a metal center, and some applications of metal complexes.
Thanks for watching!