Question

Give the electronic configuration for (a) \(\mathrm{Ti}^{3+}\) (b) \(\mathrm{Cr}^{2+}\) (c) \(\mathrm{Ru}^{4+}\) (d) \(\mathrm{Pd}^{2+}\) (e) \(\mathrm{Mo}^{3+}\)

Short answer

Question: Provide the electronic configurations for the following ions: (a) Ti³⁺ (b) Cr²⁺ (c) Ru⁴⁺ (d) Pd²⁺ (e) Mo³⁺. Answer: (a) Ti³⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ (b) Cr²⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁴ (c) Ru⁴⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d⁴ (d) Pd²⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d⁸ (e) Mo³⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d³

Step-by-step solution

(a) Finding electronic configuration for \(\mathrm{Ti}^{3+}\)

1. Identify the atomic number of Ti Titanium (Ti) has atomic number 22. 2. Find the electronic configuration of the neutral atom (Ti) The electronic configuration for Ti (atomic number 22) is \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^2}\). 3. Determine the electronic configuration of the ion (Ti\(^{3+}\)) Remove three outermost electrons from the neutral atom configuration. The first two electrons are removed from 4s and the third electron from 3d. Thus, the electronic configuration of \(\mathrm{Ti}^{3+}\) is \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^6 3d^1}\).

(b) Finding electronic configuration for \(\mathrm{Cr}^{2+}\)

1. Identify the atomic number of Cr Chromium (Cr) has atomic number 24. 2. Find the electronic configuration of the neutral atom (Cr) The electronic configuration for Cr (atomic number 24) is \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^5}\). 3. Determine the electronic configuration of the ion (Cr\(^{2+}\)) Remove two outermost electrons from the neutral atom configuration. The first electron is removed from 4s, and the second electron comes from 3d, leaving the electron configuration of \(\mathrm{Cr}^{2+}\) as \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^6 3d^4}\).

(c) Finding electronic configuration for \(\mathrm{Ru}^{4+}\)

1. Identify the atomic number of Ru Ruthenium (Ru) has atomic number 44. 2. Find the electronic configuration of the neutral atom (Ru) The electronic configuration for Ru (atomic number 44) is \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6 4d^7 5s^1}\). 3. Determine the electronic configuration of the ion (Ru\(^{4+}\)) Remove four outermost electrons from the neutral atom configuration. The first electron is removed from 5s, and the next three are removed from 4d. Thus, the electronic configuration of \(\mathrm{Ru}^{4+}\) is \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6 4d^4}\).

(d) Finding electronic configuration for \(\mathrm{Pd}^{2+}\)

1. Identify the atomic number of Pd Palladium (Pd) has the atomic number 46. 2. Find the electronic configuration of the neutral atom (Pd) The electronic configuration for Pd (atomic number 46) is \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6 4d^10}\). 3. Determine the electronic configuration of the ion (Pd\(^{2+}\)) Remove two outermost electrons from the neutral atom configuration. Both of these electrons come from 4d. Thus, the electronic configuration of \(\mathrm{Pd}^{2+}\) is \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6 4d^8}\).

(e) Finding electronic configuration for \(\mathrm{Mo}^{3+}\)

1. Identify the atomic number of Mo Molybdenum (Mo) has the atomic number 42. 2. Find the electronic configuration of the neutral atom (Mo) The electronic configuration for Mo (atomic number 42) is \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6 4d^5 5s^1}\). 3. Determine the electronic configuration of the ion (Mo\(^{3+}\)) Remove three outermost electrons from the neutral atom configuration. The first is removed from 5s, and the next two are removed from 4d. The electronic configuration of \(\mathrm{Mo}^{3+}\) is \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6 4d^3}\).

Key concepts

Electron Configuration

Understanding electron configuration is essential when studying the properties and behaviors of elements. Electron configuration refers to the arrangement of electrons in an atom's orbitals, which are defined by the atom's quantum numbers. The distribution of electrons across different energy levels and orbitals determines how the element interacts with others.
An effective method to remember the order of the orbitals is through the use of the Aufbau principle, Hund's rule, and the Pauli exclusion principle. The Aufbau principle suggests that electrons occupy orbitals starting with the lowest energy levels before moving to higher levels. For instance, 1s is filled before 2s, and so on, leading to configurations such as \(1s^2 2s^2 2p^6\) for the first ten electrons of a neutral atom.
Hund's rule tells us that within a single subshell, electrons will occupy empty orbitals alone before pairing up. The Pauli exclusion principle further indicates that no two electrons can have the same set of four quantum numbers, effectively meaning an orbital can hold a maximum of two electrons with opposite spins.
In the context of ions, the electron configuration may change due to the loss or gain of electrons, altering the element's chemical properties. For instance, Titanium normally has an electron configuration of \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^2\), but as a \(Ti^{3+}\) ion, it loses three electrons — two from the 4s orbital and one from the 3d — resulting in \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^1\). Each element possesses a unique pattern that can be predicted using these principles.

Transition Metals

Transition metals, which include elements like Titanium, Chromium, Palladium, and others, are found in the d-block of the periodic table. They are known for their ability to form positively charged ions with various oxidation states and their complex electron configurations.
A distinctive feature of transition metals is having incomplete d-orbitals, which can participate in the formation of chemical bonds. Unlike s-block elements where the outermost s-orbital electrons are typically removed first when forming ions, transition metals often lose their s-orbital electrons before d-orbital electrons when they form cations. This behavior is observed in the earlier provided examples, where \(4s\) electrons are lost before \(3d\) when the ions \(Ti^{3+}\), \(Cr^{2+}\), and \(Mo^{3+}\) are formed.
Additionally, the electron configurations of transition metals can display exceptions to the expected order. Chromium and Copper are classic examples where the \(4s\) orbital has one electron each - \(4s^1 3d^5\) instead of \(4s^2 3d^4\) - to stabilize the atom thanks to a half-filled d-orbital which is energetically favorable. Understanding these nuances is key for students to accurately deduce the properties and reactivity of transition metal ions.

Oxidation States

The oxidation state of an atom in a compound represents the degree of oxidation or reduction an atom has undergone, indicated by the number of electrons lost, gained, or shared in comparison to its neutral state. For many elements, particularly the transition metals, oxidation states can vary widely, leading to a multitude of possible compounds and reactions.
When an atom loses electrons, it becomes positively charged, resulting in a positive oxidation state. The number of electrons lost is often represented as a superscript beside the element symbol, like \(Fe^{2+}\) or \(Fe^{3+}\), indicating iron has lost two or three electrons, respectively. For transition metals, determining the correct number of electrons removed to achieve a certain oxidation state can be complex due to their ability to lose different numbers of s and d orbital electrons.
In the examples provided, the ions \(Ti^{3+}\), \(Cr^{2+}\), \(Ru^{4+}\), \(Pd^{2+}\), and \(Mo^{3+}\) all exhibit positive oxidation states indicating loss of electrons. The actual process of finding the electronic configurations of these ions involves identifying the most stable electron arrangements, which can vary depending on factors like electron pairing, overall energy of the orbital configuration, and the specific element in question. This can alter the chemical behavior and bonding characteristics of the ion, influencing how it interacts in chemical reactions.