Question

Most of the second row transition metals do not follow the normal orbital filling pattern. Five of them \(-\mathrm{Nb}, \mathrm{Mo}, \mathrm{Ru}, \mathrm{Rh},\) 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. a. \(\mathrm{Mo}, \mathrm{Mo}^{+}, \mathrm{Ag}, \mathrm{Ag}^{+}\) b. \(\mathrm{Ru}, \mathrm{Ru}^{3+}\) c. \(\mathrm{Rh}, \mathrm{Rh}^{2+}\) d. \(\mathrm{Pd}, \mathrm{Pd}^{+}, \mathrm{Pd}^{2+}\)

Short answer

Mo: \([\mathrm{Kr}] 5s^1 4d^5\), Mo⁺: \([\mathrm{Kr}] 5s^0 4d^5\), Ag: \([\mathrm{Kr}] 5s^1 4d^{10}\), Ag⁺: \([\mathrm{Kr}] 4d^{10}\), Ru: \([\mathrm{Kr}] 5s^1 4d^7\), Ru³⁺: \([\mathrm{Kr}] 4d^5\), Rh: \([\mathrm{Kr}] 5s^1 4d^8\), Rh²⁺: \([\mathrm{Kr}] 4d^7\), Pd: \([\mathrm{Kr}] 4d^{10}\), Pd⁺: \([\mathrm{Kr}] 4d^{10}\), Pd²⁺: \([\mathrm{Kr}] 4d^9\).

Step-by-step solution

- Determine Electron Configurations for Neutral Atoms

Find the ground state electron configuration for each neutral atom. According to the information given, five of the second row transition metals have a \([\mathrm{Kr}] 5s^1 4d^x\) configuration. The atomic numbers for Mo, Ru, Rh, and Ag are 42, 44, 45, and 47 respectively. Pd has a configuration of \([\mathrm{Kr}] 4d^{10}\) and its atomic number is 46. By filling in the 4d orbitals accordingly, we find the configurations: Mo: \([\mathrm{Kr}] 5s^1 4d^5\), Ru: \([\mathrm{Kr}] 5s^1 4d^7\), Rh: \([\mathrm{Kr}] 5s^1 4d^8\), Ag: \([\mathrm{Kr}] 5s^1 4d^{10}\), Pd: \([\mathrm{Kr}] 4d^{10}\).

- Determine Electron Configurations for Ions

For each ion, remove electrons starting from the highest energy level. For Mo⁺, remove one electron from the 5s orbital, for Ag⁺, remove one electron from the 5s orbital, and so on. Follow the same pattern for the other ions, remembering that electrons are removed from the s orbital before the d orbitals: Mo⁺: \([\mathrm{Kr}] 5s^0 4d^5\), Ag⁺: \([\mathrm{Kr}] 4d^{10}\), Ru³⁺: \([\mathrm{Kr}] 4d^5\), Rh²⁺: \([\mathrm{Kr}] 4d^7\), Pd⁺: \([\mathrm{Kr}] 4d^{10}\), Pd²⁺: \([\mathrm{Kr}] 4d^9\).

Key concepts

Transition Metals

Transition metals are a group of elements found in the center of the periodic table, which are known to have unique properties, including the ability to form various oxidation states and colorful compounds. In the context of the exercise, we focus on the second-row transition metals, including molybdenum (Mo), silver (Ag), ruthenium (Ru), rhodium (Rh), and palladium (Pd).

These elements are noteworthy because they often exhibit exceptions to the expected electronic configurations. For instance, instead of fully occupying the 4s orbital, they tend to have a single electron in the 5s orbital and a varying number of electrons in the 4d orbitals, as shown by the electron configuration patterns given in the exercise. Since these configurations can impact the chemistry and reactivity of the metals, understanding them is crucial for students studying chemistry.

Orbital Filling Pattern

The orbital filling pattern is a guideline that dictates the order in which electrons populate the orbitals of an atom. It follows the principle that orbitals are filled with electrons in order of increasing energy, as predicted by the Aufbau principle. The typical sequence is 1s, 2s, 2p, 3s, 3p, and so forth, with the 's' orbitals being filled before the 'p' orbitals, and 'd' orbitals generally filled after the 'p' orbitals of the same principle energy level.

However, for transition metals, there is often a deviation from this typical pattern. For example, instead of filling the 4s orbital completely before starting to fill the 4d orbitals, these metals might start filling the 4d orbital when the 4s orbital is only partially filled, leading to configurations like \[\mathrm{Kr}\] 5s^1 4d^x\. Understanding the reasons behind such deviations, like orbital energy closeness and electron-electron repulsions, is essential for grasping the behavior of these elements.

Neutral atoms and ions

Neutral atoms contain an equal number of protons and electrons, resulting in no overall charge. However, when atoms gain or lose electrons, they become ions – charged species with unequal numbers of protons and electrons. For transition metals, the formation of ions commonly involves the removal of electrons from the highest-energy orbitals first, typically following Hund's rule and the Aufbau principle.

In the exercise, we see this with the transition metals losing electrons from their 5s orbital when forming positive ions. For instance, the loss of an electron from Mo results in Mo⁺ with the configuration \[\mathrm{Kr}\] 5s^0 4d^5\. It is also significant that the 5s electron is removed before the 4d electrons, which is partly due to the 5s orbital having a slightly higher energy level than the 4d orbitals in these specific elements.

Energy Levels in Atoms

Energy levels in atoms refer to the fixed distances at which electrons orbit the nucleus. These levels contain sublevels, such as s, p, d, and f, where the electrons are found. When electrons populate these sublevels, they follow specific rules to minimize the energy of the atom.

An understanding of these energy levels and sublevels can explain the particular electron configurations seen in transition metals. For instance, the electron configurations of the second-row transition metals show us that there are instances where sublevel energies overlap, allowing for unique configurations like those of Pd, which has a full 4d orbital and no 5s electrons in its neutral state. Recognizing the difference in energy levels helps to clarify why electrons are removed from the s orbital before the d orbitals during ion formation.