Problem 75 - Chemistry The Molecular Science

Question

Nearly all first-row transition elements form \(2+\) ions. (a) For which of these elements are the \(2+\) ions paramagnetic? (b) For which element do compounds containing \(2+\) ions and chloride ions have the greatest paramagnetism? (Chloride ions have no unpaired electrons.)

Step-by-Step Solution

Answer

(a) All except Zn²⁺ are paramagnetic. (b) Mn²⁺ has the greatest paramagnetism.

Step 1: List the first-row transition elements

The first-row transition elements range from Scandium (Sc) to Zinc (Zn) in the periodic table. The full list is: Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn.

Step 2: Determine electron configurations of 2+ ions

Transition elements typically lose electrons first from their outermost shell. For first-row transition elements, this means removing electrons from the 4s sublevel. For example, Fe (with an electron configuration of [Ar] 4s² 3d⁶) becomes Fe²⁺ ([Ar] 3d⁶) after losing two electrons from the 4s orbital.

Step 3: Identify paramagnetic ions

An ion is paramagnetic if it has unpaired electrons in its electron configuration. Evaluate the electron configurations of the 2+ ions to determine which have unpaired electrons: - Sc²⁺: [Ar] 3d¹ - Ti²⁺: [Ar] 3d² - V²⁺: [Ar] 3d³ - Cr²⁺: [Ar] 3d⁴ - Mn²⁺: [Ar] 3d⁵ - Fe²⁺: [Ar] 3d⁶ - Co²⁺: [Ar] 3d⁷ - Ni²⁺: [Ar] 3d⁸ - Cu²⁺: [Ar] 3d⁹ - Zn²⁺: [Ar] 3d¹⁰ Zinc (Zn) is the only element in this group whose 2+ ion ([Ar] 3d¹⁰) has no unpaired electrons. All others are paramagnetic.

Step 4: Identify the element with greatest paramagnetism

Paramagnetism is linked to the number of unpaired electrons. Mn²⁺ has the electron configuration [Ar] 3d⁵, with five unpaired electrons, which is the maximum among first-row transition metals' 2+ ions. Therefore, Mn²⁺ shows the greatest paramagnetism.

Key Concepts

Paramagnetism

Paramagnetism is a fascinating property found in certain materials that are attracted by an external magnetic field. This is a result of unpaired electrons within the atom or ion. Each unpaired electron can be depicted like a tiny magnet, thanks to its spin. Even though the effect might seem small, it becomes noticeable when these unpaired electrons are numerous.

When analyzing paramagnetism in transition metal ions, the primary interest is in the electron configuration and the presence of these unpaired electrons. Ions with unpaired electrons are termed paramagnetic, and the strength of this magnetism increases with the number of unpaired electrons.

In the case of first-row transition metals forming 2+ ions, we observe that all form paramagnetic ions, except for zinc (Zn). Zn²⁺ has a configuration of \([\text{Ar}] \, 3d^{10}\), resulting in paired electrons and, therefore, no paramagnetism. The others, like Mn²⁺, with five unpaired electrons, display significant paramagnetic properties. This explains why materials like manganese(II) chloride are notably magnetic.

Electron Configuration

Understanding electron configurations is crucial when studying transition metal chemistry. Electron configuration describes the distribution of electrons in an atom's orbitals and is represented systematically using quantum numbers and subshell notations.

For transition metals, electron configurations follow the Aufbau principle, where electrons fill the lowest energy orbitals first, but with an exception for the 3d and 4s orbitals. For the first-row transition metals, electrons are removed from the 4s orbitals first before the 3d when forming ions. For instance, the electron configuration for a neutral iron atom (Fe) is \([\text{Ar}] \, 4s^2 \, 3d^6\), but after losing two electrons to form Fe²⁺, it becomes \([\text{Ar}] \, 3d^6\).

The pattern of filling these orbitals plays a significant role in determining the properties of the metal ions, including whether they are paramagnetic. The number and arrangement of unpaired electrons, dictated by these configurations, directly influence magnetic behavior in these ions.

Transition Elements

Transition elements, or transition metals, occupy the d-block of the periodic table, ranging from scandium (Sc) to zinc (Zn). They are unique because they can exhibit multiple oxidation states and form colorful compounds due to their ability to create complex ions.

These metals display a variety of interesting chemical properties thanks to their partially filled d-orbitals. Their electron configurations end typically in a 3d subshell, following the filling of the 4s orbitial. When transition elements form 2+ ions, as seen in this exercise, they typically lose the two 4s electrons first before involving their d-electrons. This influences their chemistry, making them essential in biological systems and catalysts in industrial processes.

Additionally, transition metals' ability to be involved in complex bonding and coordinate compounds further defines their chemical behavior and application. Understanding these properties helps explain why transition elements are so versatile and widely used across different fields.