The magnetic properties of a compound can be determined from its electron configuration and the size of its atoms. Because magnetism is generated by electronic spin, the number of unpaired electrons in a specific compound indicates how magnetic the compound is. In this section, the magnetism of the d-block elements (or transition metals) are evaluated. These compounds tend to have a large number of unpaired electrons. See the Periodic Trends page for more information.
Diamagnetism and paramagnetism are often affected by the presence of coordination complexes, which the transition metals (d-block) readily form. This article mainly discusses paramagnetism.
The spin of a single electron is denoted by the quantum number ms as +(1/2) or –(1/2). This spin is negated when the electron is paired with another, but creates a weak magnetic field when the electron is unpaired. More unpaired electrons increase the paramagnetic effects. The electron configuration of a transition metal (d-block) changes in a coordination compound; this is due to the repulsive forces between electrons in the ligands and electrons in the compound. Depending on the strength of the ligand, the compound may be paramagnetic or diamagnetic.
Some paramagnetic compounds are capable of becoming ferromagnetic. This means the compound shows permanent magnetic properties rather than exhibiting them only in the presence of a magnetic field. In a ferromagnetic element, electrons of atoms are grouped into domains in which each domain has the same charge. In the presence of a magnetic field, these domains line up so that charges are parallel throughout the entire compound. Whether a compound can be ferromagnetic or not depends on its number of unpaired electrons and on its atomic size.
- Small atoms pair up too easily and their charges cancel.
- Large atoms are difficult to keep together; their charge interaction is too weak.
Therefore, only correctly-sized atoms can form domains. Elements with this size include Fe, Co, and Ni. That means that Fe, Co and Ni are paramagnetic with the capability of permanent magnetism; in other words, they are ferromagnetic.
Ligand field theory
For a full explanation, please see the article on Ligand Field Theory. An element can have up to 10 d electrons in 5 d-orbitals: dxy, dxz, dyz, dz2, and dx2-y2. During the formation of a complex, the degeneracy (equal energy) of these orbitals is broken; the orbitals are at different energy levels.
For a six-ligand complex in an octahedral configuration, the ligands approach along the x, y, and z axes; the repulsion is therefore strongest in the orbitals along these axes (dz2 and dx2-y2). As a result, the dz2 and dx2-y2 orbitals are higher in energy than the dxy, dxz, and dyz orbitals. In a tetrahedral complex, the splitting is opposite, with the dxy, dxz, and dyz orbitals higher in energy to avoid the repulsion from the ligands approaching between the axes. The splitting in a square planar complex has four levels; they are arranged here in order of increasing energy: dyz and dxz, dxy, dz2, dx2-y2.
Depending on the strength of the ligand, the splitting energy between d-orbitals may be large or small. Ligands producing a smaller splitting energy are called "weak field" ligands, and those with a larger splitting energy are called "strong field" ligands.
Filling of d-orbitals in a complex
Hunds' Rule states that electrons fill all available orbitals with single electrons before pairing up, while maintaining parallel spins (paired electrons have opposing spins). For a set of five degenerate d-orbitals in an uncomplexed metal atom, electrons fill all orbitals before pairing to conserve pairing energy. With the addition of ligands, the situation is more complicated. The splitting energy between the d-orbitals increases the energy required to place single electrons into the higher-energy orbitals. Once the lower-energy orbitals have been half-filled (one electron per orbital), an electron can either be placed in a higher-energy orbital or paired with an electron in a lower-energy orbital. The strength of the ligands determine which option is chosen. If the splitting energy is greater than the pairing energy, the electrons will pair up; if the pairing energy is greater, unpaired electrons will occupy higher energy orbitals. In other words, with a strong-field ligand, low-spin complexes are usually formed; with a weak-field ligand, a high-spin complex is formed.
Application to Magnetism
Low-spin complexes contain more paired electrons because the splitting energy is larger than the pairing energy. These complexes, such as [Fe(CN)6]3-, are more often diamagnetic or weakly paramagnetic. Likewise, high-spin complexes usually contain more unpaired electrons because the pairing energy is larger than the splitting energy. With more unpaired electrons, high-spin complexes are often paramagnetic.
The unpaired electrons in paramagnetic compounds create tiny magnetic fields, similar to the domains in ferromagnetic materials. The strength of the paramagnetism of a coordination complex increases with the number of unpaired electrons; a higher-spin complex is more paramagnetic. The occurrence and relative strength of paramagnetism can be predicted by determining whether the compound is coordinated to a weak field ligand or a strong field ligand.
Measuring magnetism in a compound?
The Gouy balance is used to measure paramagnetism by suspending the complex in question against an equivalent weight with access to a magnetic field. We first weigh the complex without a magnetic field in its presence, then, we weigh it again in the presence of a magnetic field. If the compound is paramagnetic, it will be pulled visibly towards the electromagnet, which is the distance proportional to the magnitude of the compound's paramagnetism. If the compound, however, is diamagnetic, it will not be pulled towards the electromagnet, instead, it might even slightly be repelled by it. This will be proven by the decreased weight or the no change in weight. The change in weight directly corresponds to the amount of unpaired electrons in the compound.
"Real world" applications of magnetic properties
Ferromagnetism, the permanent magnetism associated with nickel, cobalt, and iron, is a common occurrence in everyday life. Examples of the knowledge and application of ferromagnetism include Aristotle's discussion in 625 BC, the use of the compass in 1187, and the modern-day refrigerator. Einstein declared that electricity and magnetism are inextricably linked in his theory of special relativity. He also demonstrated examples of magnets disturbed by electricity.
Paramagnetism and diamagnetism explain some of the properties of commonly-used elements and complexes. In the early days of complex-compound chemistry, paramagnetism was used to identify the shape of complexes. A technique known as electron paramagnetic resonance has been used in systems with certain para- and dia- magnetic properties to distinguish between bond types and identify the probable location of an individual element within a compound.
- Petrucci [ chapter 23. p. 968 and chapter 24 section 24-5].
- Johnson, Ronald C. and Basolo, Fred, "Coordination Chemistry: The Chemistry of Metal Complexes, W. A. Benjamin, Inc. pp40-44. 1964
- Jones, Mark M. Elementary Coordination Chemistry Prentice-Hall, Inc. 1964
- Neele Holzenkaempfer, Jesse Gipe