Source: Laboratory of Dr. Neal Abrams — SUNY College of Environmental Science and Forestry
Transition metals are found everywhere from vitamin supplements to electroplating baths. Transition metals also make up the pigments in many paints and compose all minerals. Typically, transition metals are found in the cationic form since they readily oxidize, or lose electrons, and are surrounded by electron donors called ligands. These ligands do not form ionic or covalent bonds with the metal center, rather they take on a third type of bond known as coordinate-covalent. The coordinate-covalent bond between a ligand and a metal is dynamic, meaning that ligands are continuously exchanging and re-coordinating around the metal center. The identities of both the metal and the ligand dictates which ligands will bond preferentially over another. In addition, color and magnetic properties are also due to the types of complexes that are formed. The coordination compounds that form are analyzed using a variety of instruments and tools. This experiment explores why so many complexes are possible and uses a spectrochemical (color and chemical) method to help identify the type of coordination complex that forms.
Coordination complexes have at least one metal complex, which contains a metal center and is surrounded by electron-donating ligands. This is known as a complex ion. Counterions balance the charge of the complex ion to form the molecular coordination complex. Coordination complexes are soluble in water, where the counter-ion and metal-ion complex dissociate. The metal ion and ligands behave like a polyatomic ion and do not dissociate.
The geometry of a complex ion takes on standard Valence-Shell Electron-Pair Repulsion Theory (VSEPR) geometries including linear, square planar, tetrahedral, and octahedral. Octahedral complex ions are the most common geometry. Crystal field theory explains energy splitting among d-orbitals in transition metal ions and the VSEPR geometries. Energy splitting is influenced by the shape and orientation of the d-orbital lobes.
Ligands and the Spectrochemical Series
Ligands are classified by the number of bonds, or attachments, they can make with a metal center. A single attachment is known as monodentate (one-toothed). A ligand that makes two attachments is called bidentate (two-toothed), and three attachments is known as tridentate. Ligands donate electron density to the metal center to form the coordinate-covalent bond. Ligands may be charged or neutral. Ligands are classified as being strong or weak according to the spectrochemical series:
(weak) I- < Br- < Cl- < SCN- < F- < OH- < ox2-< ONO- < H2O < NCS- < EDTA4- < NH3 < en < NO2- < CN- (strong)
When six ligands approach a metal center to form an octahedral complex, the five degenerate d-orbitals split into three lower-energy degenerate t2g orbitals and two higher-energy degenerate eg orbitals. The distance of the splitting between the t2g and eg orbitals is dictated by the strength of the ligand according to the spectrochemical series.
Hund's rule still applies and electrons fill orbitals one at a time, but they fill in accordance to size of the splitting of the t2g and eg orbitals. If the split is small, electrons will fill all the orbitals singly before pairing. This maximizes the number of unpaired electrons and is called high-spin. Likewise, a strong-field causes a large t2g-eg split: electrons pair in the t2g set before filling the higher-energy eg orbitals. This minimizes the number of unpaired electrons and is called low-spin. The drive for electrons to pair is governed by the energy (or size) of the orbital splitting compared to the energy of electron pairing. If the energy of pairing is high compared to the energy of moving into the eg orbitals, then electrons are high-spin. If the energy of pairing is low compared to the energy of moving into the higher eg orbitals, then electrons are low-spin.
The distance that electrons have to move from the lower t2g state to the higher eg state in the metal center dictates the energy of electromagnetic radiation that the complex absorbs. If that energy is in the visible region (400-700 nm, 1.77 eV - 3.1 eV), the complex generally has a color. Weak-field ligands (I- → OH-) cause small splittings and complexes absorb low-energy light (i.e. red) which appear green in color. Strong-field ligands (EDTA → CN-) absorb high-energy light (i.e. blue-violet) and appear red-yellow in color. Complexes with ligands that are between strong and weak on the spectrochemical series, like ammonia, can adopt either a weak- or strong-field geometry.
The color-ligand relationship is the rationale for the name "spectrochemical series". The number of paired and unpaired electrons also gives rise to paramagnetic diamagnetic properties in metal complexes.
When four ligands coordinate around a metal center, either a square planar or tetrahedral complex can result. The orbital energies in tetrahedral complexes are flipped compared to octahedral complexes, with eg being lower in energy than t2g. This is due to the orientations of d-orbitals with respect to the coordinating ligands. In square planar complexes, there are several differences in orbital energies, with dyz and dxz being degenerate and lowest in energy (lower than dz2,), then dxy, and finally the highest energy dx2-y2 orbital.
Structure and Color
Because the distance in orbital splitting varies with ligand strength, a coordination complex with the same metal center can have a variety of colors based upon the coordinating ligand. For example, an aqueous solution of Ni(H2O)62+ has a light green color, but Ni(NH3)62+ is deep blue. The color arises from the change in energy between the t2g-eg orbitals. NH3 is a stronger field ligand, which pushes the orbitals further apart from one another as well as displaces the H2O ligands from the metal center. We will further explore the effect of ligands on color and coordination complexes in this experiment.
1. Nickel Complexes and Colors
- Ni(H2O)62+ complex (Figure 1a)
- Prepare a 1 M solution of Ni(H2O)62+ by dissolving NiSO4 in the appropriate volume of water.
- Further dilute the Ni(H2O)62+solution by adding 70 mL of the 1 M solution to 1,000 mL of deionized water.
- Divide the Ni(H2O)62+ among seven 400-mL beakers.
- The aqueous nickel solution takes on a light green color since water is a weak-field ligand.
- The absorbance spectrum indicates red wavelengths are absorbed, justifying the opposite, green, which is observed.
- Ni(NH3)62+ complex (Figure 1b)
- Add a 5 M aqueous ammonia solution to a beaker and stir.
- The solution takes on a deep blue color, indicating the solution is absorbing orange light which is higher in energy than red light.
- The absorbance spectrum indicates yellow wavelengths are absorbed, justifying the opposite, blue, which is observed.
- Ammonia is a stronger field ligand than water, which increases the splitting between the t2g and eg orbitals.
- Ni(en)32+ complex (Figure 1c)
- Add a 30% ethylenediamine (en) solution to the aqueous Ni(H2O)62+ complex and stir.
- The solution progressively turns from light blue to blue to purple as ethylenediamine molecules progressively coordinate around the metal center to eventually form Ni(en)3+.
- Ethylenediamine is a stronger ligand than water or ammonia and it is bidentate. The purple color indicates the solution is absorbing yellow light which is higher in energy than orange or red light.
- The absorbance spectrum indicates yellow wavelengths are absorbed, justifying the opposite, purple, which is observed.
- Ni(dmg)22+ complex (Figure 1d)
- Dimethylgloxine (dmg) is a bidentate ligand that chelates a large number of metals. Only two dmg molecules are required per metal center because Ni(dmg)22+ has a square-planar geometry.
- Add 1% dmg to the aqueous complex.
- A solid pink/red precipitate forms, the insoluble Ni(dmg)22+ complex.
- A visible transmission spectrum of the complex is not possible, but the red color indicates green light is being absorbed. Green is higher energy than yellow, orange, and red.
- Ni(CN)42- complex (Figure 1e)
- The cyanide ion (CN-) is monodentate, but a very strong field ligand, which also forms square-planar complexes with nickel (II).
- Add a 1 M KCN solution.
- A yellow Ni(CN)42- complex forms almost immediately.
- Note: Working with cyanide salts must be done with great care. Addition of acid may result in the formation of cyanide gas.
- Cyanide is a stronger ligand than any of the other ligands because there is σ-bonding from the ligand to the metal and π-back bonding from the metal to the ligand. The yellow color indicates the solution is absorbing blue light, which is higher in energy than green, yellow, orange, and red.
- The absorbance spectrum indicates yellow wavelengths are absorbed, justifying the opposite, purple, which is observed.
2. Ligand Strength
- According to the spectrochemical series, some ligands are stronger-field than others, which correspond to the size of the splitting of the d-orbitals of the central metal ion.
- Stronger field ligands replace weaker field ligands in solution.
- An aqueous solution of nickel sulfate appears light green because the Ni(H2O)62+ complex forms.
- Sequentially add solutions of ammonia, ethylenediamine, dimethylglyoxime, and cyanide to the nickel-containing solution while stirring.
- After each addition, the previous color disappears and the new color appears.
- The color change indicates the formation of a new coordination complex driven by the strength of the ligand. These can be quantified by the equilibrium constant for each reaction:
Ni(H2O)62+(aq) + 6 NH3 (aq) → Ni(NH3)62+ (aq) + 6 H2OKeq = 1.2 x 109
Ni(NH3)62+ (aq) + 3 en(aq) → Ni(en)32+ (aq) + 6 NH3 (aq)Keq = 1.1 x 109
Ni(en)32+ (aq) + 2 Hdmg(aq) → Ni(dmg)2 (s) + 3 en(aq) + 2 H+ (aq)Keq = 1.35 x 105
Ni(dmg)2 (s) + 4 CN- (aq) - → Ni(CN)4-2 (aq) + 2 dmg- (aq)Keq = 6.3 x 107
- The equilibrium constant in each reaction is very large (>1), indicating that the reactions are all product driven.
Figure 1. Structures of nickel (II) coordination complexes a-e.
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.
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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.
- 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).