Question

How many unpaired electrons are present in each of the firstrow transition metals in the ground state?

Short answer

In the ground state, the first-row transition metals have the following unpaired electrons: Scandium (Sc) - 1, Titanium (Ti) - 2, Vanadium (V) - 3, Chromium (Cr) - 5, Manganese (Mn) - 5, Iron (Fe) - 4, Cobalt (Co) - 3, Nickel (Ni) - 2, Copper (Cu) - 1, and Zinc (Zn) - 0.

Step-by-step solution

Determine Electron Configurations

For each element, we will write down its electron configuration, focusing on the d orbital. Sc: [Ar] 4s2 3d1 Ti: [Ar] 4s2 3d2 V: [Ar] 4s2 3d3 Cr: [Ar] 4s1 3d5 Mn: [Ar] 4s2 3d5 Fe: [Ar] 4s2 3d6 Co: [Ar] 4s2 3d7 Ni: [Ar] 4s2 3d8 Cu: [Ar] 4s1 3d10 Zn: [Ar] 4s2 3d10

Count Unpaired Electrons

We will now count the unpaired electrons in the d orbital for each of these elements. Sc: 1 unpaired electron Ti: 2 unpaired electrons V: 3 unpaired electrons Cr: 5 unpaired electrons Mn: 5 unpaired electrons Fe: 4 unpaired electrons Co: 3 unpaired electrons Ni: 2 unpaired electrons Cu: 1 unpaired electron Zn: 0 unpaired electrons In summary, the first-row transition metals have the following unpaired electrons in their ground state: - Scandium (Sc): 1 - Titanium (Ti): 2 - Vanadium (V): 3 - Chromium (Cr): 5 - Manganese (Mn): 5 - Iron (Fe): 4 - Cobalt (Co): 3 - Nickel (Ni): 2 - Copper (Cu): 1 - Zinc (Zn): 0

Key concepts

Electron Configuration

Electron configuration is a method to describe the distribution of electrons across an atom's various orbitals. This helps in understanding an element’s chemical properties. For transition metals, the electron configuration typically fills the 3d sublevel or d orbitals, following the 4s orbital. The reason behind this order is balancing energy levels to ensure lower energy states are filled first.
For the first-row transition metals, their electron configurations often appear as:

• Scandium (Sc): [Ar] 4s² 3d¹
• Titanium (Ti): [Ar] 4s² 3d²
• Vanadium (V): [Ar] 4s² 3d³
• Chromium (Cr): [Ar] 4s¹ 3d⁵
• Manganese (Mn): [Ar] 4s² 3d⁵
• Iron (Fe): [Ar] 4s² 3d⁶
• Cobalt (Co): [Ar] 4s² 3d⁷
• Nickel (Ni): [Ar] 4s² 3d⁸
• Copper (Cu): [Ar] 4s¹ 3d¹⁰
• Zinc (Zn): [Ar] 4s² 3d¹⁰

In these examples, [Ar] stands for Argon, representing the filled electron shells that come before the transition metals.

Unpaired Electrons

Unpaired electrons are those electrons in an atom that do not have a partner with opposite spins in their orbitals. The presence of unpaired electrons is a key factor in determining the magnetic properties of a substance. Atoms with unpaired electrons are often paramagnetic, meaning they are attracted to a magnetic field. For transition metals, particularly those in the first-row, the number of unpaired electrons can change from element to element due to the complex filling of d orbitals. Here’s a quick breakdown:

• Scandium (Sc): 1 unpaired electron
• Titanium (Ti): 2 unpaired electrons
• Vanadium (V): 3 unpaired electrons
• Chromium (Cr): 5 unpaired electrons - Chromate’s unusual electron distribution maximizes unpaired electrons.
• Manganese (Mn): 5 unpaired electrons
• Iron (Fe): 4 unpaired electrons
• Cobalt (Co): 3 unpaired electrons
• Nickel (Ni): 2 unpaired electrons
• Copper (Cu): 1 unpaired electron
• Zinc (Zn): 0 unpaired electrons - a fully filled d subshell results in diamagnetic properties.

Understanding unpaired electrons is crucial for predicting how these metals interact with magnetic fields and other chemicals.

d Orbitals

The d orbitals are crucial in understanding the chemistry of transition metals. These sublevels hold vital importance in determining an atom's configuration and consequent chemical behavior. In a standard set, there are five d orbitals, labeled as:

• \(d_{xy}, d_{xz}, d_{yz}, d_{x^{2}-y^{2}}, d_{z^{2}}\)

These labels define orbitals based on their orientation in three-dimensional space. Each d orbital can hold a maximum of two electrons, each with opposite spins.
For transition metals, these d orbitals start filling after the s orbitals of the previous shell. They play a fundamental role in the following:

• Determining an element’s electron configuration.
• Influencing the magnetic and spectral properties due to unpaired electrons.
• Impacting the chemical reactivity and bonding due to their energy states.

When evaluating the electron configuration, it's essential to remember that filling these orbitals affects an atom’s placement in the periodic table and its chemical behavior.