Problem 68 - General Chemistry

Question

The tetraaquazinc(II) ion, \(\left[\mathrm{Zn}\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}\right]^{2+}\), is colorless. Explain why this might be expected.

Step-by-Step Solution

Answer

The ion is colorless because the Zn(II) ion has completely filled d-orbitals, preventing d-d transitions.

Step 1: Understanding the Compound

The tetraaquazinc(II) ion is composed of a zinc ion surrounded by four water molecules, with zinc having a +2 charge. In its chemical notation, it is represented as \([\mathrm{Zn}(\mathrm{H}_2\mathrm{O})_4]^{2+}\).

Step 2: Transition Metals and Color

Transition metals typically have compounds that are colored due to d-d electron transitions within the d-orbitals when they absorb visible light.

Step 3: Electronic Configuration of Zn(II)

The zinc ion, Zn(II), in the compound has lost two electrons, making its electron configuration: \([\mathrm{Ar}]\,3d^{10}\). In this state, all of the d-orbitals are completely filled.

Step 4: Explaining the Lack of Color

Since the d-orbitals in Zn(II) are fully occupied, there are no empty d-orbitals available for d-d transitions. Without these transitions, there’s no absorption of visible light, resulting in a colorless appearance of the compound.

Key Concepts

Tetraaquazinc(II) Ion

The tetraaquazinc(II) ion, represented as \(\left[\mathrm{Zn}(\mathrm{H}_2\mathrm{O})_4\right]^{2+}\), is a complex ion formed by zinc and water molecules. In this complex, the zinc ion is surrounded by four water molecules acting as ligands. Ligands are molecules that can donate a pair of electrons to the central metal ion, in this case, zinc.

The zinc in this ion carries a 2+ charge, which means it has two fewer electrons than a neutral zinc atom. These water molecules are neutrally charged and form coordinate covalent bonds with the zinc ion, completing the structure of the complex.

The overall charge on the complex ion is derived from the positive charge of the zinc ion, balanced by the negative charges it attracts, but the neutral water molecules help stabilize this charge configuration.

d-d Electron Transitions

When discussing transition metals, color plays a fascinating role. Typically, many transition metal complexes are colored due to "d-d transitions."

Now, what are these d-d transitions? In transition metals, electrons can jump between different d-orbitals when visible light is absorbed. This absorption is specific to the energy difference between the d-orbitals. The specific light energy absorbed determines the complex's color.

For example, if a complex absorbs light in the blue spectrum, it will appear orange, which is the complementary color. This property makes transition metal complexes quite diverse in color depending on their ligands and metal ions. However, when dealing with zinc(II), things change a bit due to its electron configuration.

Zn(II) Electron Configuration

Understanding zinc's electron configuration sheds light on why some zinc compounds are colorless. In its neutral state, the electron configuration for zinc is \([\mathrm{Ar}] 4s^2 3d^{10}\), meaning it has completely filled d-orbitals along with two 4s electrons.

When zinc forms a \(2+\) ion, it loses the two 4s electrons resulting in the configuration \([\mathrm{Ar}] 3d^{10}\). Notice that the 3d orbitals remain fully filled even after ionization.

This full d-orbital configuration is crucial for understanding the colorlessness of \(\left[\mathrm{Zn}(\mathrm{H}_2\mathrm{O})_4\right]^{2+}\). Since all the d-orbitals are filled, there are no vacant d-orbitals available for an electron to jump into if it were to absorb visible light. Without such jumps or d-d transitions, no visible light absorption occurs, and hence the complex remains colorless.