Problem 47 - Chemistry The Central Science In Si Units

Question

Give three examples of +2 ions that have an electron configuration of \(n d^{10}(n=3,4,5 \ldots)\).

Step-by-Step Solution

Answer

Examples: \(Zn^{2+}\), \(Cd^{2+}\), \(Hg^{2+}\).

Step 1: Understand the Problem

We need to find three ions with a +2 charge that achieves the electron configuration of a completely filled d-orbital, specifically denoted by \(nd^{10}\). The subscript \(n\) represents the principal quantum number denoting energy levels beyond 3.

Step 2: Identify Elements with Filled D-Subshells

Elements that typically end in \(d^{10}\) configurations are the group 12 transition metals: Zinc (Zn), Cadmium (Cd), and Mercury (Hg). These elements naturally end with a \(d^{10}\) configuration because they fill their d-orbitals completely in their neutral state as Zn:\([Ar] 3d^{10} 4s^2\), Cd: \([Kr] 4d^{10} 5s^2\), and Hg: \([Xe] 4f^{14} 5d^{10} 6s^2\).

Step 3: Form Their +2 Ions

Upon forming a +2 ion, these elements lose two electrons usually from the s-orbital. This results in \(Zn^{2+}\), \(Cd^{2+}\), and \(Hg^{2+}\) all of whom have the electron configurations: Zn: \([Ar] 3d^{10}\), Cd: \([Kr] 4d^{10}\), Hg: \([Xe] 4f^{14} 5d^{10}\). In each +2 state, the \(d\)-orbital remains completely filled.

Key Concepts

Electron Configuration

Electron configuration is an essential concept in chemistry. It defines how electrons are arranged in an atom or ion. Electrons fill available energy levels around the nucleus from lower to higher energies. This order follows the Aufbau principle.

The s, p, d, and f orbitals represent different energy levels and sublevels. Each orbital type has a specific capacity for holding electrons:

• s-orbital: can hold up to 2 electrons.
• p-orbital: can hold up to 6 electrons.
• d-orbital: can hold up to 10 electrons.
• f-orbital: can hold up to 14 electrons.

In transition metals, the electron configuration gets intriguing. For example, in an atom like zinc (Zn), you first fill the 4s orbital before the 3d orbital. However, when forming positive ions, electrons are often removed from the s-orbital first. This is why a zinc ion ( Zn^{2+} ) ends with a 3d configuration.

D-Orbital

The d-orbital plays a crucial role in transition metals. It is responsible for some of their unique properties, such as high melting and boiling points, as well as their ability to form brightly colored compounds.

A single d-orbital can hold a maximum of 10 electrons, which leads to different electron configurations based on the number of electrons they actually contain.

When a transition metal forms an ion, typically, it will lose electrons from an outer s-orbital first, which keeps the d-orbital fully or partially filled. Having a filled d-orbital can stabilize the ion because lower energy configurations are more stable. An example of this is seen in the group 12 metal ions like Zn^{2+} or Hg^{2+}, which maintains a filled d subshell by the loss of outer shell s electrons.

Group 12 Elements

Group 12 elements are a special set of transition metals. These include Zinc (Zn), Cadmium (Cd), and Mercury (Hg). They lie in the d-block of the periodic table and are known for having a fully filled d orbital even in their elemental forms.

These metals have unique properties:

• They are typically not very reactive compared to other transition metals.
• They generally form stable +2 oxidation states.
• They have practical uses in industries - for example, mercury in thermometers and zinc for galvanizing to prevent corrosion.

The neutral atoms of Zn, Cd, and Hg have the electron configurations closing in a filled d-orbital: Zn is [Ar] 3d^{10} 4s^{2}, Cd is [Kr] 4d^{10} 5s^{2}, and Hg is [Xe] 4f^{14} 5d^{10} 6s^{2}. Upon ionization to Zn^{2+}, Cd^{2+}, and Hg^{2+}, they lose their s electrons but maintain totally filled d orbitals. This keeps their configurations stable, making these ions unique and predictable in chemical reactions.