Question

Most of the second-row transition metals do not follow the normal orbital filling pattern. Five of them - Nb, Mo, Ru, Rh, and Ag-have a \([\mathrm{Kr}] 5 s^{1} 4 d^{x}\) configuration and \(\mathrm{Pd}\) has a \([\mathrm{Kr}] 4 d^{10}\) configuration. Write the ground state electron configuration for each species. $$\begin{array}{ll}{\text { a. } \mathrm{Mo}, \mathrm{Mo}^{+}, \mathrm{Ag}, \mathrm{Ag}^{+}} & {\text { b. } \mathrm{Ru}, \mathrm{Ru}^{3+}} \\ {\text { c. } \mathrm{Rh}, \mathrm{Rh}^{2+}} & {\text { d. Pd, Pd }^{+}, \mathrm{Pd}^{2+}}\end{array}$$

Short answer

a. Mo: \([\mathrm{Kr}] 5s^{1} 4d^{5}\), Mo+: \([\mathrm{Kr}] 4d^{5}\)Ag: \([\mathrm{Kr}] 5s^{1} 4d^{10}\), Ag+: \([\mathrm{Kr}] 4d^{10}\)b. Ru: \([\mathrm{Kr}] 5s^{1} 4d^{7}\), Ru3+: \([\mathrm{Kr}] 4d^{6}\)c. Rh: \([\mathrm{Kr}] 5s^{1} 4d^{8}\), Rh2+: \([\mathrm{Kr}] 4d^{7}\)d. Pd: \([\mathrm{Kr}] 4d^{10}\), Pd+: \([\mathrm{Kr}] 4d^{9}\), Pd2+: \([\mathrm{Kr}] 4d^{8}\)

Step-by-step solution

Write Electron Configurations for Neutral Atoms

Start by writing the ground state electron configurations for the neutral atoms based on the given information. Molybdenum (Mo), whose atomic number is 42, has an outer shell electron configuration of \([\mathrm{Kr}] 5s^{1} 4d^{5}\). Silver (Ag), with atomic number 47, has \([\mathrm{Kr}] 5s^{1} 4d^{10}\). Ruthenium (Ru), atomic number 44, has \([\mathrm{Kr}] 5s^{1} 4d^{7}\). Rhodium (Rh), with atomic number 45, has \([\mathrm{Kr}] 5s^{1} 4d^{8}\). Palladium (Pd), with atomic number 46, is special in that it has filled d orbitals: \([\mathrm{Kr}] 4d^{10}\).

Determine Electron Configuration for Cations

When an atom loses electrons to form a cation, it typically loses them from the outermost (highest energy) orbitals first. For Mo, which has one 5s electron, Mo+ would lose this electron, leaving \([\mathrm{Kr}] 4d^{5}\) for Mo+. For Ag, Ag+ would similarly lose its 5s electron, leaving \([\mathrm{Kr}] 4d^{10}\) for Ag+. For Ru, the Ru3+ ion would lose three electrons, two from the 5s orbital and one from the 4d orbital, resulting in \([\mathrm{Kr}] 4d^{6}\) for Ru3+. Rh2+ loses two electrons from the 5s and one of the 4d orbitals, resulting in \([\mathrm{Kr}] 4d^{7}\) for Rh2+. For Pd, Pd+ would lose an electron from the 4d level (since there is no 5s electron), leaving \([\mathrm{Kr}] 4d^{9}\) for Pd+. For Pd2+, it would lose another 4d electron, resulting in \([\mathrm{Kr}] 4d^{8}\).

Compile the Electron Configurations

List the electron configurations for each species:a. Mo: \([\mathrm{Kr}] 5s^{1} 4d^{5}\), Mo+: \([\mathrm{Kr}] 4d^{5}\)Ag: \([\mathrm{Kr}] 5s^{1} 4d^{10}\), Ag+: \([\mathrm{Kr}] 4d^{10}\)b. Ru: \([\mathrm{Kr}] 5s^{1} 4d^{7}\), Ru3+: \([\mathrm{Kr}] 4d^{6}\)c. Rh: \([\mathrm{Kr}] 5s^{1} 4d^{8}\), Rh2+: \([\mathrm{Kr}] 4d^{7}\)d. Pd: \([\mathrm{Kr}] 4d^{10}\), Pd+: \([\mathrm{Kr}] 4d^{9}\), Pd2+: \([\mathrm{Kr}] 4d^{8}\)

Key concepts

Transition Metals

Transition metals comprise a large and important category of chemical elements located at the center of the periodic table. Characterized by their partially filled d orbitals, these metals often exhibit variable oxidation states and a propensity for forming colorful compounds. Notably, transition metals like iron, copper, and nickel play pivotal roles in biological processes and industrial applications.

Understanding the distinctive behaviors of transition metals, including exceptions in their electron configurations, is fundamental to grasping the complex chemistry they exhibit. While most elements follow a predictable pattern of orbital filling, these metals sometimes deviate due to the energy stability gained by having half-filled or completely filled d orbitals, as seen in certain second-row transition metals.

Atomic Structure

Atomic structure determines the physical and chemical properties of elements. At its core is the nucleus, surrounded by a cloud of electrons arranged in orbitals of varying shapes and energies. Electrons occupy these orbitals according to the Pauli exclusion principle, Hund's rule, and the Aufbau principle, which together guide the filling of orbitals in order of increasing energy.

Each electron is uniquely defined by four quantum numbers: principal (n), azimuthal (l), magnetic (m_l), and spin (m_s). The outermost electrons, particularly those in the s and d orbitals for transition metals, are crucial in determining an atom's chemical behavior.

Orbital Filling Pattern

The orbital filling pattern is the sequence in which electrons populate the orbitals of an atom. Typically, this pattern adheres to the Aufbau principle, which states that each electron occupies the lowest energy orbital available. However, this pattern can be more complex for transition metals.

Partially filled subshells, specifically d orbitals, can influence the energy levels. As a result, the expected order can be altered, leading to exceptions. For example, in elements like chromium and copper, the 4s orbital is filled before the 3d, despite the 3d being lower in energy, because half-filled or fully filled d orbitals confer extra stability.

Cation Electron Configuration

The electron configuration of a cation, an ion with a positive charge, is derived from the configuration of its neutral atom by removing electrons. Electrons are typically removed from the highest-energy orbitals, usually the outermost shells, in accordance with Hund's rule. For transition metals, this often means electrons are lost from s orbitals before d orbitals, even when the s orbitals were filled later.

For instance, when silver (Ag) forms a cation (Ag⁺), it loses the outermost 5s electron, retaining the \(\mathrm{Kr}\) 4d¹⁰ configuration. Knowledge of how cations form is essential for predicting the chemical behavior of elements, especially in the context of ionic compounds and redox reactions.