Problem 84 - Chemistry The Molecular Science

Question

Give the electron configuration of (a) \(\mathrm{Cr}^{2+}\). (b) \(\mathrm{Zn}^{2+}\) (c) \(\mathrm{Co}^{2+}\). (d) \(\mathrm{Mn}^{4+}\).

Step-by-Step Solution

Answer

(a) Cr²⁺: [Ar] 3d⁴ (b) Zn²⁺: [Ar] 3d¹⁰ (c) Co²⁺: [Ar] 3d⁷ (d) Mn⁴⁺: [Ar] 3d³

Step 1: Find the Atomic Number

To determine the electron configuration of an ion, first find the atomic number of the neutral element: - Chromium (Cr): Atomic number 24. - Zinc (Zn): Atomic number 30. - Cobalt (Co): Atomic number 27. - Manganese (Mn): Atomic number 25.

Step 2: Write the Neutral Atom Configuration

Write the electron configuration for each neutral atom: - Cr: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹ - Zn: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² - Co: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁷ 4s² - Mn: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s²

Step 3: Adjust for Ion Charge

Determine the electron configuration by removing electrons for the given ion charge (remove from 4s before 3d): - Cr²⁺: Remove 2 electrons from 4s → 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁴ - Zn²⁺: Remove 2 electrons from 4s → 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ - Co²⁺: Remove 2 electrons from 4s and 3d → 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁷ - Mn⁴⁺: Remove 4 electrons from 4s and 3d → 1s² 2s² 2p⁶ 3s² 3p⁶ 3d³

Key Concepts

Transition Metals

Transition metals are elements found in the center of the periodic table, specifically in groups 3 to 12. They are characterized by having a partially filled d subshell in their atomic or common oxidation states. Transition metals, such as Chromium (Cr), Zinc (Zn), Cobalt (Co), and Manganese (Mn), exhibit unique properties due to this electron configuration.

Some of these unique properties include:

• Variable oxidation states, allowing them to form ions with different charges.
• Colored compounds, which arise from electronic transitions between d-orbitals.
• Potential for forming complex ions with ligands, enhancing their roles as catalysts.

The interaction of d electrons plays a crucial role in these properties. As you consider electron configurations of transition metal ions, remember that you typically remove electrons from the 4s orbital before the 3d, despite the order in which they are filled.

Ion Charge

Ion charge refers to the total charge of an ion which can be positive (cation) or negative (anion) based on whether electrons have been lost or gained. In transition metals, ions are usually formed by losing electrons. While calculating the electron configuration of an ion, it is essential to adjust for its charge.

For example:

• When forming a +2 charged ion, two electrons are usually removed.
• For \(\mathrm{Cr}^{2+}\), two electrons are removed primarily from the 4s orbital, turning its configuration into \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^4\).
• Similarly, for \(\mathrm{Mn}^{4+}\), four electrons are removed, leading to \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^3\).

Always remember, for transition metals, electrons from the 4s subshell are removed before electrons from the 3d subshell when they form cations.

Atomic Number

The atomic number of an element is the number of protons present in its nucleus. It is a fundamental identifier of elements and defines the element's position on the periodic table. For instance, Chromium (Cr), with an atomic number of 24, means that a neutral Chromium atom contains 24 protons and 24 electrons.

Here’s how the atomic number impacts electron configuration:

• The number of electrons corresponds to the atomic number in a neutral atom.
• For example, Cobalt (Co) with an atomic number of 27 has a neutral electron configuration of \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^7 4s^2\).
• In ions, the total number of electrons will differ from the atomic number due to loss or gain of electrons.

Understanding the atomic number helps in determining the starting point for writing both neutral and ionic electron configurations, providing clarity on electron arrangements and ion formations.