Question

When concentrated hydrochloric acid is added to a red solution containing the \(\operatorname{Co}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{2+}\) complex ion, the solution turns blue as the tetrahedral \(\mathrm{CoCl}_{4}^{2-}\) complex ion forms. Explain this color change.

Short answer

The color change from red to blue when concentrated hydrochloric acid is added to a solution containing the \(\operatorname{Co}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{2+}\) complex is due to the replacement of water molecules with chloride ions, forming the \(\mathrm{CoCl}_{4}^{2-}\) complex. This new ligand environment changes the splitting of the cobalt ion's d-orbitals and the electronic transitions within the complex. Consequently, the complexes absorb light at different wavelengths, causing the observed color change; the initial complex appears red while the final complex appears blue.

Step-by-step solution

Identify the initial complex and its reaction with hydrochloric acid

The starting complex is \(\operatorname{Co}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{2+}\), which has a red color. When concentrated hydrochloric acid is added, it results in the formation of the \(\mathrm{CoCl}_{4}^{2-}\) complex, which has a blue color.

Describe the reaction between the cobalt complex and hydrochloric acid

When concentrated hydrochloric acid is added to the solution, its high concentration of chloride ions (Cl-) will replace the water molecules surrounding the cobalt ion. This results in the following reaction: \[ \operatorname{Co}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{2+} + 4\mathrm{Cl}^{-} \rightarrow \mathrm{CoCl}_{4}^{2-} + 6\mathrm{H}_{2}\mathrm{O} \] In this reaction, the hexaaquacobalt(II) complex ion reacts with chloride ions to form the tetrachlorocobaltate(II) complex ion and releases water as a product.

Explain the color change based on electronic transitions

The observed color change is due to the electronic transition occurring within the ligand field of the cobalt complex. In the original \(\operatorname{Co}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{2+}\) complex, the cobalt ion is in an octahedral environment surrounded by water molecules. This arrangement causes the d-orbitals to split into two levels: t2g and eg. When hydrochloric acid is added and the complex changes to \(\mathrm{CoCl}_{4}^{2-}\), the cobalt ion now finds itself in a different, tetrahedral ligand field environment surrounded by chloride ions. This change in the ligand environment also changes the splitting of the d-orbitals, resulting in different energy gaps between the orbitals. As a consequence, the absorbed light's wavelength - and therefore the color that our eyes perceive - is different for the two complexes. The \(\operatorname{Co}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{2+}\) complex absorbs blue light and appears red, while the \(\mathrm{CoCl}_{4}^{2-}\) complex absorbs yellow-orange light and appears blue.

Summary

To summarize, the concentrated hydrochloric acid leads to a replacement of water molecules in the \(\operatorname{Co}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{2+}\) complex by chloride ions, forming the \(\mathrm{CoCl}_{4}^{2-}\) complex. The new ligand environment results in a change of the electronic transitions happening in the cobalt complex, leading to a different color being absorbed by the complex. The result is a color change from red to blue in the solution.

Key concepts

Transition metal complexes

Transition metal complexes play a crucial role in coordination chemistry. They are formed when metal ions bond with a set of molecules or ions known as ligands. Transition metals are unique due to their d-orbitals, which allow these complex formations. In the case of the cobalt complex in our exercise, the metal ion is cobalt, and it initially coordinates with water molecules to form \(\operatorname{Co}(\mathrm{H}_{2} \mathrm{O})_{6}^{2+}\).

This complex arrangement allows transition metals to exhibit a variety of colors. These colors result from the wavelengths of light absorbed due to electronic transitions within the complex. What's fascinating is that when hydrochloric acid is added, the nature and arrangement of the complex change due to the replacement of water ligands with chloride ions, forming the new tetrachlorocobaltate(II) complex, \(\mathrm{CoCl}_{4}^{2-}\). The altered coordination sphere showcases the diverse capabilities of transition metal chemistry.

Key aspects include:

• Cobalt can exist in different coordination environments, affecting its electronic structure.
• The nature of ligands (water vs. chloride) can drastically change the properties of the transition metal complex, including color.

Ligand field theory

Ligand field theory is fundamental in understanding the behavior of transition metal complexes. This theory explains how ligands cause the d-orbitals of a transition metal to split into groups with different energy levels when they form a complex. The octahedral and tetrahedral complexes have different splitting patterns. For instance, in our exercise, the \(\operatorname{Co}(\mathrm{H}_{2} \mathrm{O})_{6}^{2+}\) complex is octahedral, leading to a specific d-orbital splitting pattern.

When chloride ligands replace water, converting the complex into \(\mathrm{CoCl}_{4}^{2-}\), the crystal field changes from octahedral to tetrahedral. This change alters the energy levels and splitting of the d-orbitals. In tetrahedral complexes, the orbitals split into a lower energy set and a higher energy set, but in a different manner than what is observed in octahedral complexes. The degree of splitting and the relative energies define the color we perceive in transition metal complexes.

Essentials to consider:

• Octahedral and tetrahedral configurations result in distinct energy level diagrams.
• Ligand type and arrangement significantly impact the electronic structure of the metal ion.
• The difference in energy levels corresponds to absorbed light wavelengths, affecting visible color.

Electronic transitions

Electronic transitions in coordination complexes are the movement of electrons between different d-orbitals. The color change observed in our cobalt complex reaction is due to these transitions. In the \(\operatorname{Co}(\mathrm{H}_{2} \mathrm{O})_{6}^{2+}\) complex, electrons move between d-orbitals split by the presence of water ligands.

When the complex switches to \(\mathrm{CoCl}_{4}^{2-}\) upon adding hydrochloric acid, the ligands' change modifies the energy levels of the d-orbitals. As a result, different wavelengths of light are absorbed compared to the original complex in water. Initially, the complex absorbs blue light, appearing red. After forming the \(\mathrm{CoCl}_{4}^{2-}\) complex, it absorbs yellow-orange light, making it appear blue.

Important points:

• Electronic transitions occur between different energy levels influenced by the ligand environment.
• The frequency of absorbed light determines the color of the complex.
• The transition from one ligand field environment to another leads to a visible change in color.