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 \([\mathrm{Ar}] 4 \mathrm{~s}^{2} 3 d^{6} ;\) 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) \(\mathrm{Ti}^{4+}\).

Short answer

The electron configurations of the given transition metal cations are: a) \(\mathrm{Zn}^{2+}\): \([\mathrm{Ar}] 3d^{10}\) b) \(\mathrm{Pt}^{2+}\): \([\mathrm{Xe}] 4f^{14} 5d^7\) c) \(\mathrm{Cr}^{3+}\): \([\mathrm{Ar}] 3d^1\) d) \(\mathrm{Ti}^{4+}\): \([\mathrm{Ar}]\)

Step-by-step solution

Write the Ground State Electron Configuration

First, we need to write the ground state electron configuration for each transition metal. We can do this by using the periodic table and filling the orbitals in the correct order: a) \(\mathrm{Zn}\): \([\mathrm{Ar}]~4s^2 3d^{10}\) b) \(\mathrm{Pt}\): \([\mathrm{Xe}]~6s^2 4f^{14} 5d^{9}\) c) \(\mathrm{Cr}\): \([\mathrm{Ar}]~4s^{1} 3d^{5}\) d) \(\mathrm{Ti}\): \([\mathrm{Ar}]~4s^2 3d^{2}\)

Remove Electrons to Form Cations

Now that we have the ground state electron configurations, we can remove the appropriate number of electrons to form the cations. Remember, we need to remove \(s\) electrons first, followed by the \(d\) electrons: a) For \(\mathrm{Zn}^{2+}\), remove 2 electrons: \([\mathrm{Ar}]~4s^{0} 3d^{10} \Rightarrow [\mathrm{Ar}] 3d^{10}\) b) For \(\mathrm{Pt}^{2+}\), remove 2 electrons: \([\mathrm{Xe}]~6s^{0} 4f^{14} 5d^{9} \Rightarrow [\mathrm{Xe}] 4f^{14} 5d^{9-2} = [\mathrm{Xe}] 4f^{14} 5d^7\) c) For \(\mathrm{Cr}^{3+}\), remove 3 electrons: \([\mathrm{Ar}]~4s^{1} 3d^{5} \Rightarrow [\mathrm{Ar}]~4s^{0} 3d^{2} \Rightarrow [\mathrm{Ar}] 3d^{3-2} = [\mathrm{Ar}] 3d^{1}\) d) For \(\mathrm{Ti}^{4+}\), remove 4 electrons: \([\mathrm{Ar}]~4s^{2} 3d^{2} \Rightarrow [\mathrm{Ar}]~4s^{0} 3d^{0} = [\mathrm{Ar}]\)

Write the Final Electron Configurations

Finally, we can write the electron configurations for each cation: a) \(\mathrm{Zn}^{2+}\): \([\mathrm{Ar}] 3d^{10}\) b) \(\mathrm{Pt}^{2+}\): \([\mathrm{Xe}] 4f^{14} 5d^7\) c) \(\mathrm{Cr}^{3+}\): \([\mathrm{Ar}] 3d^1\) d) \(\mathrm{Ti}^{4+}\): \([\mathrm{Ar}]\)

Key concepts

Electron Removal in Transition Metals

In transition metals, unlike other elements, electrons are not removed from the outermost shell or the d-orbitals first. Instead, electrons are removed from the s-orbitals when forming cations. This can initially be counterintuitive since their s-orbitals are filled before the d-orbitals.

For example, consider iron, which has a ground state electron configuration of \([\mathrm{Ar}]~4s^2 3d^6\). When forming the \(\mathrm{Fe}^{2+}\) ion, two electrons are removed from the 4s orbital, resulting in the electron configuration \([\mathrm{Ar}]~3d^6\). While this seems different from typical removal from the d-orbital, it is a characteristic of the special stability associated with d-orbital electrons in transition metals.

Cation Formation in d-block Elements

In the d-block elements, also known as transition metals, cations are formed by the removal of electrons primarily from the s subshells, and then from the d subshells if necessary.

This is because when the s and d subshells are in close proximity energetically and spatially, the s electrons are less tightly held compared to the d electrons. Hence, they are the first to be removed during ionization.

Example Highlights:

• For \( ext{Zn}^{2+}\), remove 2 electrons yielding: \([\mathrm{Ar}]~3d^{10}\).
• For \( ext{Cr}^{3+}\), we remove 3 electrons, resulting in: \([\mathrm{Ar}]~3d^1\).

This removal method evidences the different electron holding strengths between the s and d subshells, promoting cation stability.

Ground State Electron Configuration

The ground state electron configuration provides a depiction of the distribution of electrons among the atomic orbitals under a neutral (non-ionized) state. It serves as the base configuration from which cations are formed by removing electrons.

Understanding the ground state configuration aids in predicting how transition metals form ions. For example, in zinc, the ground configuration \([ ext{Ar}]~4s^2 3d^{10}\) helps predict its cation form (\( ext{Zn}^{2+}\)) as \([ ext{Ar}]~3d^{10}\).

By convention, electron configurations are written starting from the lower energy orbitals to higher (1s, 2s, 2p, etc.), respecting Hund's rule and the Pauli exclusion principle for electron filling. This information equips us with a comprehensive understanding of electronic rearrangements in transition metals.

Transition Metals Electron Configurations

Transition metals have unique electron configurations that arise from their position in the periodic table. They typically have partially filled d-orbitals, which contribute to their chemical properties and behavior, such as forming colored compounds and various oxidation states.

These configurations can also differ remarkably between the neutral atoms and their cations. For instance, platinum has a ground state of \([\mathrm{Xe}]~6s^2 4f^{14} 5d^{9}\), and its cation \( ext{Pt}^{2+}\) is \([\mathrm{Xe}]~4f^{14} 5d^7\) as electrons are removed from the s orbit before the d orbit.

Knowledge of these configurations is critical for understanding transition metals' complex chemistry and physical characteristics, such as magnetic properties and catalytic behavior.

Electron Configuration Notation

Electron configuration notation is a shorthand method used to represent the distribution of electrons in an atom's orbitals. This notation is essential for depicting the electron arrangement clearly and succinctly.

The notation employs numbers, letters, and superscripts. The numbers indicate the main energy levels, the letters (s, p, d, f) represent the subshell type, and the superscripts denote the number of electrons in a subshell.

For example:

• \([ ext{Ar}]~4s^2 3d^{10}\)
• \([ ext{Xe}]~6s^2 4f^{14} 5d^9\)

This form of notation makes it easier to understand and remember the arrangement and removal of electrons in transition metals when forming cations.