Question

Recall from Chapter 4 that \(C r\) and \(C u\) are exceptions to the normal orbital filling, resulting in a \([A r] 4 s^{1} 3 d^{x}\) configuration. Write the ground state electron configuration for each species. $${ a. }\ \mathrm{Cr}, \mathrm{Cr}^{+}, \mathrm{Cr}^{2+}, \mathrm{Cr}^{3+} \quad \text { b. } \mathrm{Cu}, \mathrm{Cu}^{+}, \mathrm{Cu}^{2+}$$

Short answer

\text{Cr}: [Ar] 4s^1 3d^5, \text{Cr}^+: [Ar] 3d^5, \text{Cr}^{2+}: [Ar] 3d^4, \text{Cr}^{3+}: [Ar] 3d^3, \text{Cu}: [Ar] 4s^1 3d^{10}, \text{Cu}^+: [Ar] 3d^{10}, \text{Cu}^{2+}: [Ar] 3d^9.

Step-by-step solution

Ground State Configuration of Neutral Chromium (Cr)

Chromium has an atomic number of 24. Its electron configuration is an exception to the typical filling order. Instead of fully filling the 4s before the 3d, one electron is taken from the 4s orbital to half-fill the 3d orbital, resulting in a configuration of \[\text{Cr}: [Ar] 4s^1 3d^5\].

Ground State for Chromium Ions

For ions of Chromium:\- Chromium ion (Cr+): One electron is removed, typically from the highest energy level, which is the 4s orbital, resulting in \[\text{Cr}^+: [Ar] 3d^5\].\- Chromium ion (Cr2+): A second electron is removed, also from the 4s orbital, leading to \[\text{Cr}^{2+}: [Ar] 3d^4\].\- Chromium ion (Cr3+): A third electron is removed, this time from 3d orbital, resulting in \[\text{Cr}^{3+}: [Ar] 3d^3\].

Ground State Configuration of Neutral Copper (Cu)

Copper has an atomic number of 29. Similar to Chromium, it too is an exception to the normal orbital filling order. One electron from the 4s orbital is moved to the 3d orbital to fully fill it, giving a configuration of \[\text{Cu}: [Ar] 4s^1 3d^{10}\].

Ground State for Copper Ions

For the ions of Copper:\- Copper ion (Cu+): One electron is removed from the 4s orbital, since it has a higher energy level than 3d, resulting in \[\text{Cu}^+: [Ar] 3d^{10}\]. There is no change in the 3d orbital as it's fully filled and more stable.\- Copper ion (Cu2+): A second electron is removed, now from the 3d orbital, resulting in \[\text{Cu}^{2+}: [Ar] 3d^9\].

Key concepts

Ground State Electron Configuration

Understanding the ground state electron configuration is crucial for comprehending how atoms interact and form compounds. The ground state refers to the most stable, lowest energy arrangement of electrons in an atom. Electrons fill subshells (s, p, d, f) based on the Aufbau principle, which states they occupy the lowest available energy orbital first.

The electron configuration for any given element follows a specific sequence, often represented using the periodic table as a guide. For instance, the electron configuration of hydrogen, with one electron, is simply \(1s^1\). However, as we look at heavier elements, this pattern becomes more complex, involving additional subshells with different energies.

For heavier elements like chromium (Cr) and copper (Cu), the ground state electron configuration is affected by electron repulsions and desires for increased stability, leading to seemingly anomalous configurations. This is where the typical \(4s\) before \(3d\) filling order can see exceptions. Understanding these exceptions is essential for accurately predicting the properties and chemical behavior of these transition metals.

Transition Metals

Transition metals are elements found in the d-block of the periodic table, known for their ability to form various oxidation states and colorful compounds. These metals include groups 3-12 on the periodic table and have partially filled d orbitals. This unique electron arrangement is responsible for their distinctive chemical and physical properties.

When writing electron configurations for transition metals, realize their d orbitals become more stable when half-filled or fully filled. This is due to electron-electron repulsions being minimized and the exchange energy being maximized in these configurations. This leads to unusual energy shifts between \(4s\) and \(3d\) orbitals, causing irregularities in their electron configurations. For example, chromium prefers a \(4s^1 3d^5\) configuration as a half-filled \(3d\) subshell provides added stability, and copper chooses \(4s^1 3d^{10}\) for a completely filled \(3d\) subshell.

Orbital Filling Exceptions

In the world of electron configurations, there are certain exceptions to the expected order of orbital filling. While most elements follow the Aufbau principle, transition metals such as chromium and copper exhibit notable anomalies. These exceptions occur because the energies of the \(4s\) and \(3d\) orbitals are very close, and in some specific cases, it's more energetically favorable to deviate from the norm.

This behavior results in exceptions like chromium having a \(4s^1 3d^5\) configuration instead of \(4s^2 3d^4\). Removing an electron from the \(4s\) orbital and adding it to the \(3d\) orbital allows chromium to achieve a half-filled \(3d\) subshell, which is exceptionally stable. The same rationale applies to copper's \(4s^1 3d^{10}\) configuration, where the stability of a fully-filled \(3d\) subshell overrides the expected \(4s^2 3d^9\) configuration.

When removing electrons to form ions, these transition metals usually lose the \(4s\) electrons before losing the \(3d\) electrons, despite the \(4s\) being filled after the \(3d\) in the neutral atom's electron configuration. This can seem counterintuitive but further emphasizes the need to understand the underlying principles governing electron configurations and the stability associated with certain electron arrangements.