Question

Write the ground state electron configuration for each atom and ion pair. a. \(\mathrm{Zr}, \mathrm{Zr}^{2+}\) b. \(\mathrm{Co}, \mathrm{Co}^{2+}\) c. \(\mathrm{Tc}, \mathrm{Tc}^{3+}\) d. Os, \(\mathrm{Os}^{4+}\)

Short answer

a. Zr: [Kr]4d^25s^2, Zr^2+: [Kr]4d^2; b. Co: [Ar]3d^74s^2, Co^2+: [Ar]3d^7; c. Tc: [Kr]4d^55s^2, Tc^3+: [Kr]4d^4; d. Os: [Xe]4f^145d^66s^2, Os^4+: [Xe]4f^145d^4.

Step-by-step solution

Finding the Atomic Numbers

First, look up the atomic numbers for Zirconium (Zr), Cobalt (Co), Technetium (Tc), and Osmium (Os). The atomic number corresponds to the number of protons and, in a neutral atom, also the number of electrons.

Writing Electron Configurations for Neutral Atoms

Using the atomic numbers from Step 1, write the electron configuration for each neutral atom. Place electrons in orbitals following the Aufbau principle, Hund's rule, and the Pauli exclusion principle.

Writing Electron Configurations for Ions

Remove electrons from the highest energy level (usually the outermost shell) to form the ions. For the 2+ ions, remove two electrons, and for the 3+ and 4+ ions, remove three and four electrons, respectively.

Checking Electron Configurations

Verify that the electron configurations for the ions follow the expected order based on the periodic table, considering the number of electrons removed.

Key concepts

Aufbau Principle

Understanding how electrons are arranged in an atom is vital for predicting chemical behaviors. The Aufbau principle is the guideline that scientists use to determine this arrangement. It states that electrons fill up atomic orbitals in the order of increasing energy levels, starting with the lowest. Think of it like a hotel where electrons check in to the lowest energy rooms first before moving to higher ones.

When applying the Aufbau principle, we list electron orbitals from least to most energy and populate them with electrons until we reach the atom's total electron count, which equals its atomic number for neutral atoms. For instance, the element Zirconium (Zr) starts filling electrons from the 1s orbital, then follows in ascending energy to the 4d orbital.

Hund's Rule

Hund's rule deals with how electrons are distributed across orbitals of the same energy, known as 'degenerate orbitals'. It's like assigning seats on a bus; according to Hund's rule, electrons will sit alone in an orbital with their spin unpaired until they have to double up. In essence, every orbital in a subshell gets one electron before any gets a second.

This means that if we're filling the d orbitals, each of the five d orbitals gets one electron first. This rule minimizes electron repulsion and makes atoms more stable. For Cobalt (Co), this rule is crucial when populating the 3d orbitals with electrons.

Pauli Exclusion Principle

The Pauli exclusion principle is a quantum rule that states no two electrons in an atom can have the same set of four quantum numbers. In simpler terms, it's like a strict dress code - no two electrons can be identical in their energy, orbit, and spin direction. Each electron orbital can hold two electrons maximum, and they must have opposite spins.

When writing electron configurations, we indicate the spin of each electron with arrows pointing up or down (such as ↑↓). This is key for understanding the configurations of ions too. In elements like Technetium (Tc), the Pauli exclusion principle helps explain why electron removal starts from the highest energy level when forming Tc^3+.

Atomic Numbers

The atomic number of an element, unique to each element, tells us the number of protons in the nucleus of an atom. Furthermore, in a neutral atom, the atomic number also equals the number of electrons. So, if we know the atomic number, we know how many electrons we need to place in orbitals when writing ground state electron configurations.

For example, Osmium (Os) has an atomic number of 76, which means a neutral osmium atom has 76 electrons. Knowing this allows us to fill electron orbitals accordingly, starting from 1s going up to the 6s and beyond into the f and d orbitals, following the other principles.

Electron Orbitals

Electron orbitals are regions within an atom where electrons are most likely to be found. They come in various shapes (s, p, d, f) and are distributed across different energy levels. Each orbital can contain a set number of electrons: s orbitals can hold 2, p orbitals can hold 6, d orbitals can hold 10, and f orbitals can hold 14.

When we write electron configurations, we are essentially mapping out which orbitals the electrons occupy in an atom for both neutral atoms and ions. This mapping relies on the other principles we discussed to ensure we're placing electrons correctly in a way that reflects the atom's most stable and ground-state configuration.