Magnetic Properties of Coordination Compounds

1.0Introduction

The study of magnetic properties of coordination compounds is an important aspect of the JEE exam syllabus. This subject area gives an understanding of the electronic structure, bonding and geometry of coordination compounds. The magnetic behaviour of a species is derived from the movement of electrons. Each electron behaves like a very small magnet owing to orbital movement around the nucleus and intrinsic spin. The inherent magnetic moment is crucial because it allows us to predict how the species will behave in an external magnetic field.

2.0Basics of Magnetism in Chemistry

Many transition metal complexes contain unpaired electrons and are therefore paramagnetic. In contrast, molecules like N₂ and ions such as Na⁺ and [Fe(CN)₆]⁴⁻, which have all electrons paired, are diamagnetic. Diamagnetic substances are slightly repelled by magnetic fields.

The extent of paramagnetism in a compound is measured by its magnetic moment (μ). A larger μ indicates stronger paramagnetism.

Magnetic moment arises from both spin and orbital angular momentum. In a nonspherical environment, the contribution from orbital angular momentum may be quenched. However, the spin-only magnetic moment remains and depends solely on the total number of unpaired electrons.

When an atom or ion has unpaired electrons, the spin of these electrons generates a magnetic moment, making the species paramagnetic. The magnitude of the magnetic moment is directly proportional to the number of unpaired electrons—the more unpaired electrons present, the larger the magnetic moment.

Magnetic susceptibility measures the force a substance experiences in a magnetic field. When the weight of a sample is compared with its apparent weight in a magnetic field, paramagnetic substances appear heavier due to the attractive force of the field. This increase in apparent weight can be used to calculate the number of unpaired electrons in the substance.

3.0Types of Magnetic Behavior

Substances are classified into different categories based on their response to an external magnetic field. For coordination compounds, the most common types are diamagnetism and paramagnetism.

Magnetic Moment

The magnetic moment (µ) is calculated using the spin-only formula: $\mu =\sqrt{n(n+2)}BM$

Where:

• n = Number of unpaired electrons
• BM = Bohr Magneton

Types of Magnetic Behavior

Paramagnetism

• A magnetic field outside of a paramagnetic material pulls it towards it. This property comes from having one or more unpaired electrons. Each unpaired electron has a small magnetic moment, and these moments align with the external field, which causes the attraction. The more unpaired electrons there are, the stronger this attraction is.
• Has at least one unpaired electron.
• Most complexes of transition metals are paramagnetic.

Diamagnetism

• An external magnetic field repels diamagnetic substances. This behavior is seen in compounds where all electrons are paired. In this state, the magnetic moments of the electrons cancel each other out, resulting in a net magnetic moment of zero.
• Diamagnetic: All paired electrons.
• Examples include complexes of Zn²⁺  or Sc^3+, where the d-orbital is filled or empty. 

Ferromagnetism

Ferromagnetic substances exhibit a strong attraction to magnetic fields and retain their magnetism even when the external field is no longer present. This occurs when the magnetic moments of unpaired electrons in a solid crystal align in the same direction over large areas, known as domains. Iron (Fe), Cobalt (Co), and Nickel (Ni) are classic examples, but this doesn't happen as often in simple coordination compounds.

The extent of bending of field lines depends on the relative permeability (${\mu }_r$​) of the material:

${\mu }_r>>1(\text{ferromagnetic}),{\mu }_r\gtrsim 1(\text{paramagnetic}),{\mu }_r<1(\text{diamagnetic})$

4.0Factors Affecting Magnetic Properties

Hybridization of Central Metal Ion

• Inner orbital (low-spin) → fewer unpaired electrons → may be diamagnetic.
• Outer orbital (high-spin) → more unpaired electrons → paramagnetic.

High-spin (weak field ligand):

                             ↑↓  ↑   ↑         d-orbitals partially filled

Low-spin (strong field ligand):

                             ↑↓  ↑↓  ↑↓      d-orbitals fully paired 

Oxidation State

• Higher oxidation → more splitting of d-orbitals → may reduce the number of unpaired electrons⁻.
• Example:
• • Fe²⁺ (d⁶, weak field ligand) → paramagnetic.
  • Fe³⁺ (d⁵, strong field ligand) → can become diamagnetic in low-spin.

Number of Unpaired Electrons

• Directly decides the magnitude of the magnetic moment.
• More unpaired electrons⁻ → stronger paramagnetism.

5.0Examples