Magnetic Properties of Complex Ions: Octahedral Complexes: Videos & Practice Problems

In octahedral complexes, the nature of the ligand influences the energy of d orbitals. Weak field ligands create a small crystal field splitting energy (Δ), leading to degenerate orbitals and high spin complexes, which are typically paramagnetic. Conversely, strong field ligands result in a large Δ, causing low spin complexes that may be diamagnetic. Understanding these concepts is crucial for predicting the magnetic properties of coordination compounds and their behavior in various chemical reactions.

For octahedral complexes, Weak-Field Ligands create High-spin complexes and Strong-Field Ligands create Low-spin complexes.

In octahedral complexes, the nature of the ligands attached to the central metal ion significantly influences the arrangement of electrons in the d orbitals. When weak field ligands are present, they create a small crystal field splitting energy, denoted as Δ. This small Δ means that the d orbitals remain nearly degenerate, allowing electrons to fill the lower energy orbitals before moving to the higher energy ones. As a result, these complexes are classified as high spin complexes, which typically exhibit paramagnetism due to the presence of unpaired electrons.

Conversely, strong field ligands lead to a larger Δ, causing a greater energy difference between the lower and higher d orbitals. In this scenario, electrons are less likely to occupy the higher energy orbitals, resulting in the formation of low spin complexes. It is important to note that when dealing with strong field ligands in octahedral complexes, one must assess the overall electron configuration to determine whether the complex is diamagnetic (with all electrons paired) or paramagnetic (with unpaired electrons). This distinction is crucial for understanding the magnetic properties of the complex based on the type of ligands involved.

Example

To determine the spin and magnetism of the cobalt complex ion with six ammonia ligands and an overall charge of +2, we start by analyzing the cobalt ion's electron configuration. Cobalt, with an atomic number of 27, has a neutral electron configuration of Ar 4s²3d⁷. When cobalt is in the +2 oxidation state, it loses two electrons from the highest energy level, which is the 4s orbital. This results in the cobalt ion having 7 d electrons in the 3d orbitals.

Next, we identify the nature of the ligand. Ammonia (NH₃) is a neutral ligand and is classified as a strong field ligand. This classification is important because strong field ligands create a large crystal field splitting energy (Δ), which influences how the d electrons are arranged in the presence of the ligand. According to the crystal field theory, strong field ligands tend to favor low spin configurations, meaning that electrons will fill the lower energy orbitals before pairing up in higher energy orbitals.

To visualize this, we can draw the octahedral crystal field splitting diagram. In this diagram, the d orbitals split into two sets: the lower energy orbitals (t_2g) and the higher energy orbitals (e_g). For our cobalt complex, we fill the 7 d electrons according to the low spin configuration. Following Hund's rule, we fill the lower energy orbitals first:

1. Fill the three t_2g orbitals with one electron each (3 electrons total).
2. Pair up the next three electrons in the t_2g orbitals (3 paired electrons).
3. Place the final 7th electron in one of the t_2g orbitals.

This results in a configuration where there is one unpaired electron remaining. Since the presence of unpaired electrons is a characteristic of paramagnetic species, we conclude that this cobalt complex is paramagnetic.

In summary, the cobalt complex ion with six ammonia ligands exhibits a low spin configuration and is classified as paramagnetic due to the presence of one unpaired electron in its d orbitals.

Determine the spin and number of unpaired electrons in the following complex ion: [Mn(en)₃]³⁺.

• A. high-spin; 1 unpaired electron
• B. low-spin; 1 unpaired electron
• C. high-spin; 2 unpaired electrons
• D. low-spin; 2 unpaired electrons

Answer: (D)

Determine the spin and number of unpaired electrons in the following complex ion: [Cd(H­₂O)₄]²⁺.

• A. high-spin; 2 unpaired electrons
• B. low-spin; 2 unpaired electrons
• C. high-spin; 0 unpaired electrons
• D. low-spin; 4 unpaired electrons

Answer: (C)

Which of the following complex ions is/are diamagnetic in nature?
I. [Mn(Br₄)]^2–                              II. [V(NO₂)₄]^4–                                 III. [Zn(NH₃)₄]²⁺                          IV. [Sc(H₂O)₆]³⁺

• A. I, II, and IV
• B. II and III
• C. I and II
• D. I, III, and IV
• E. iii and Iv

Answer: (E)

Here’s what students ask on this topic:

What is the difference between high spin and low spin octahedral complexes?

In octahedral complexes, the difference between high spin and low spin relates to the arrangement of electrons in the d orbitals, influenced by the type of ligands. High spin complexes occur with weak field ligands, which result in a small crystal field splitting energy (Δ). This allows electrons to occupy higher energy orbitals, leading to unpaired electrons and paramagnetism. Conversely, low spin complexes form with strong field ligands, creating a large Δ. Electrons pair in lower energy orbitals, often resulting in diamagnetism due to the absence of unpaired electrons. The spin state affects the magnetic properties and color of the complex.

How do weak field and strong field ligands affect the magnetic properties of octahedral complexes?

Weak field ligands in octahedral complexes lead to a small crystal field splitting energy (Δ), resulting in high spin complexes. In these complexes, electrons occupy higher energy orbitals, often leaving unpaired electrons, which makes the complex paramagnetic. Strong field ligands, on the other hand, cause a large Δ, leading to low spin complexes. Here, electrons pair in lower energy orbitals, potentially resulting in diamagnetism if all electrons are paired. The type of ligand thus directly influences whether the complex is paramagnetic or diamagnetic, based on the electron arrangement in the d orbitals.

Why are high spin complexes usually paramagnetic?

High spin complexes are usually paramagnetic because they contain unpaired electrons in their d orbitals. In these complexes, weak field ligands cause a small crystal field splitting energy (Δ), allowing electrons to occupy higher energy orbitals without pairing. This results in unpaired electrons, which contribute to the magnetic moment of the complex, making it paramagnetic. Paramagnetism arises from the presence of these unpaired electrons, which align with an external magnetic field, enhancing the magnetic properties of the complex.

What determines whether an octahedral complex is paramagnetic or diamagnetic?

The magnetic properties of an octahedral complex, whether paramagnetic or diamagnetic, are determined by the electron configuration in the d orbitals, influenced by the type of ligands. Weak field ligands lead to high spin complexes with unpaired electrons, resulting in paramagnetism. Strong field ligands create low spin complexes, where electrons pair in lower energy orbitals, potentially leading to diamagnetism if all electrons are paired. The specific arrangement of electrons, dictated by the ligand field strength, ultimately determines the magnetic nature of the complex.

How does the crystal field splitting energy (Δ) influence the spin state of octahedral complexes?

The crystal field splitting energy (Δ) is crucial in determining the spin state of octahedral complexes. A small Δ, associated with weak field ligands, results in high spin complexes where electrons occupy higher energy orbitals, leading to unpaired electrons and paramagnetism. Conversely, a large Δ, caused by strong field ligands, leads to low spin complexes. In these, electrons pair in lower energy orbitals, potentially resulting in diamagnetism if all electrons are paired. Thus, Δ directly influences whether a complex is high spin or low spin, affecting its magnetic properties.