Question

Without referring to the text, write the ground-state electron configurations of the first-row transition metals. Explain any irregularities.

Short answer

Cr: [Ar] 3d⁵ 4s¹ and Cu: [Ar] 3d¹⁰ 4s¹ are irregular due to subshell stability.

Step-by-step solution

Understanding First-Row Transition Metals

The first-row transition metals are the elements from Scandium (Sc) to Zinc (Zn) in the periodic table. These elements have partially filled d subshells.

Identify Atomic Numbers

List the atomic numbers for Sc to Zn: Sc (21), Ti (22), V (23), Cr (24), Mn (25), Fe (26), Co (27), Ni (28), Cu (29), Zn (30).

Writing Electron Configurations

For each metal, write the electron configuration by filling the 4s and then the 3d orbitals: - Sc: [Ar] 3d¹ 4s² - Ti: [Ar] 3d² 4s² - V: [Ar] 3d³ 4s² - Cr: [Ar] 3d⁵ 4s¹ (irregular) - Mn: [Ar] 3d⁵ 4s² - Fe: [Ar] 3d⁶ 4s² - Co: [Ar] 3d⁷ 4s² - Ni: [Ar] 3d⁸ 4s² - Cu: [Ar] 3d¹⁰ 4s¹ (irregular) - Zn: [Ar] 3d¹⁰ 4s²

Explain Irregularities

Chromium (Cr) and Copper (Cu) have irregular configurations. For Cr, having 3d⁵ 4s¹ instead of 3d⁴ 4s² is due to the stability of a half-filled d subshell. For Cu, 3d¹⁰ 4s¹ provides a full d subshell, enhancing stability.

Key concepts

Transition Metals

Transition metals are unique elements found in the middle of the periodic table. They are located in groups 3-12 and are known for their ability to form various oxidation states. This ability comes from their partially filled d orbitals.
Transition metals include familiar elements such as iron, copper, and gold. These metals are known for their distinctive properties:

• They often form colorful compounds due to d–d electron transitions.
• Many transition metals are good conductors of electricity and heat.
• They are typically hard and have high melting and boiling points.

Transition metals play an essential role in many industrial applications and biological processes, making their study crucial for both chemistry and materials science.

Ground-State

The ground state of an atom refers to the lowest energy state or the most stable configuration of its electrons.
In their ground state, atoms have electrons occupying the lowest available energy orbitals, starting from the innermost shell and moving outward.
When writing electron configurations for atoms in their ground state, the Aufbau principle is used, which means "building up." According to this principle, electrons fill atomic orbitals in an order that minimizes energy:

• The process typically begins with the 1s orbital.
• Next is the 2s, followed by 2p, and so on.

Sometimes, as with certain transition metals, the predicted electron configuration might change slightly. For instance, gaining added stability, certain elements like chromium and copper will have electrons in different orbits than expected in their ground states. This adjustment provides these atoms with a more stable electron configuration.

Periodic Table

The periodic table is an organized arrangement of all known chemical elements. It was designed to highlight the recurring "periodic" properties of elements.
By understanding the layout of the periodic table, we can predict the properties of elements and their compounds.

• The table is divided into horizontal rows known as periods and vertical columns called groups or families.
• Transition metals are located in the center of the table, in groups 3-12.
• Elements are arranged by increasing atomic number, which reflects the number of protons in an atom's nucleus.

The periodic table has an important role in chemistry because it not only organizes information but also predicts chemical behavior. This insight is vital for everything from developing new materials to understanding processes in living organisms.

3d Orbital

The 3d orbital is a set of orbitals within an atom where electrons reside. These orbitals are part of the third energy level and become relevant starting from the transition metals.
Orbitals are regions in an atom where electrons are most likely to be found. The d orbital, in particular, can hold up to 10 electrons in five sub-orbitals, each with two electrons.

• The filling of the 3d orbital occurs after the 4s orbital, despite being part of the lower principal energy level.
• This order happens because the 4s orbital is lower in energy than the 3d orbital when it starts to fill.
• However, after the 3d orbital begins to take in electrons, it becomes lower in energy compared to the 4s orbital, which explains why elements like chromium and copper have irregular configurations.

Understanding the 3d orbital is crucial for predicting the chemical properties and reactivity of transition metals.