Problem 70 - Chemistry: The Central Science

Question

In the transition metals (the \(d\) -block), the electron configuration of cations is different than what you might expect. Instead of the \(\mathrm{d}\) electrons being lost first, \(s\) electrons are lost first. For example, the electron configuration of iron, \(\mathrm{Fe}\), is \([\mathrm{Ar}] 4 s^{2} 3 d^{6}\); but the electron configuration of \(\mathrm{Fe}^{2+}\) is \([\mathrm{Ar}] 3 d^{6} ;\) the \(4 \mathrm{~s}\) electrons are eliminated to make the cation. Write out the electron configurations of (a) \(\mathrm{Zn}^{2+}\) (b) \(\mathrm{Pt}^{2+}\) (c) \(\mathrm{Cr}^{3+}\) (d) \(\mathrm{Ti}^{4+}\).

Step-by-Step Solution

Answer

(a) Zn²⁺: \([Ar] 3d^{10}\) (b) Pt²⁺: \([Xe] 4f^{14} 5d^{9}\) (c) Cr³⁺: \([Ar] 3d^{4}\) (d) Ti⁴⁺: \([Ar]\)

Step 1: Write the ground-state electron configurations for neutral atoms

First, let's write the ground state electron configurations for the neutral atoms using the periodic table as a reference. 1. Zinc (Zn): \([Ar] 4s^2 3d^{10}\) 2. Platinum (Pt): \([Xe] 6s^2 4f^{14} 5d^{9}\) 3. Chromium (Cr): \([Ar] 4s^1 3d^{5}\) 4. Titanium (Ti): \([Ar] 4s^2 3d^{2}\)

Step 2: Determine the electron configurations for the cationic species

Now, let's write out the electron configurations for the cationic species by removing electrons from 's' orbitals before 'd' orbitals. 1. Zinc (Zn²⁺): Remove 2 electrons from 4s orbital. So, the electron configuration is: \([Ar] 3d^{10}\) 2. Platinum (Pt²⁺): Remove 2 electrons from 6s orbital. So, the electron configuration is: \([Xe] 4f^{14} 5d^{9}\) 3. Chromium (Cr³⁺): Remove 2 electrons from 4s orbital first and then remove one more electron from 3d orbital. So, the electron configuration is: \([Ar] 3d^{4}\) 4. Titanium (Ti⁴⁺): Remove 2 electrons from 4s orbital first and then remove 2 more electrons from 3d orbital. So, the electron configuration is: \([Ar]\) So, the electron configurations for the cationic species are: (a) Zn²⁺: \([Ar] 3d^{10}\) (b) Pt²⁺: \([Xe] 4f^{14} 5d^{9}\) (c) Cr³⁺: \([Ar] 3d^{4}\) (d) Ti⁴⁺: \([Ar]\)

Key Concepts

Electron Configuration

Electron configuration is like a roadmap showing how electrons are arranged in an atom. Electrons orbit the nucleus in different energy levels, which are further divided into subshells: \(s\), \(p\), \(d\), and \(f\). For transition metals, this configuration often ends in the \(d\) subshell as they are located in the \(d\)-block of the periodic table.

A key rule is that electrons fill subshells starting with the lowest energy level. However, it's important to note that when transition metals become ions, electrons are removed from the outermost \(s\) subshell before the \(d\) subshell.

This might seem counterintuitive since \(d\) electrons are added last, but it's because the \(s\) subshell is energetically higher when these atoms are ionized.

• Example: \([\mathrm{Ar}] 4s^2 3d^6\) is the configuration for neutral iron (Fe).
• When ionized to \(\mathrm{Fe}^{2+}\), it becomes \([\mathrm{Ar}] 3d^6\) because the \(s\) electrons are lost first.

Cations

Cations are ions with a positive charge, formed when an atom loses one or more electrons. Transition metals often form cations with various possible charges, influencing their chemical behavior.

When a transition metal forms a cation, the process involves the removal of electrons, primarily affecting the number of \(s\) and \(d\) electrons. This electron removal determines the new electron configuration and is essential for predicting properties like color, magnetism, and reactivity.

Key points:

• Loss of \(s\) electrons happens before \(d\) electrons in most transition metal cations.
• This affects the metal's oxidation state and its ability to form complex ions.

For example, zinc (Zn) loses two \(s\) electrons to form \(\mathrm{Zn}^{2+}: [\mathrm{Ar}] 3d^{10}\).

d-block Elements

d-block elements, also known as transition metals, occupy the central strip of the periodic table, from groups 3 to 12. These metals are unique due to the presence of electrons in the \(d\) subshell, which influences their chemical and physical properties.

The \(d\)-block is renowned for a couple of specific characteristics:

• They usually have multiple oxidation states. This enables a wide array of chemical reactions and color variations.
• The \(d\) electrons can participate in bonding, allowing these metals to form complex structures like coordination compounds.

Subtle electronic changes in the \(d\) subshell can also control properties like electrical conductivity and catalytic activity. In platinum (Pt), for example, removing \(s\) electrons forms a cation \(\mathrm{Pt}^{2+} : [\mathrm{Xe}] 4f^{14} 5d^9\), altering its chemical behavior.