Write the general electron configuration of a transition element (a) in Period \(5 ;\) (b) in Period 6 .

The Periodic Table is a valuable tool that arranges all known elements in a systematic way based on their atomic number, electron configurations, and recurring chemical properties. Each row in the Periodic Table is called a 'period,' while each column is referred to as a 'group.' Elements are broadly classified into metals, nonmetals, and metalloids.

Transition metals are located in groups 3-12 and periods 4-7. These elements play a significant role in various chemical reactions due to their ability to form multiple oxidation states and complex ions. Understanding the layout of the Periodic Table is crucial for identifying element properties and predicting behaviors in reactions.

Electron configuration describes the distribution of electrons in an atom's orbitals. It's essential to understand this concept to predict an element's chemical properties and how it will interact with other elements. The general format for electron configuration is writing out the electrons in their respective energy levels and sublevels (s, p, d, f).

For example, the electron configuration for Carbon (atomic number 6) is 1s² 2s² 2p². Orbital filling follows specific rules, such as the Aufbau principle (electrons occupy the lowest energy orbitals first), Hund's Rule (every orbital in a subshell is singly occupied before any orbital is doubly occupied), and the Pauli Exclusion Principle (no two electrons can have the same set of quantum numbers). Understanding these rules is fundamental for grasping more complex configurations, especially for transition metals.

Transition metals are a unique group of elements found in the d-block of the Periodic Table, specifically in groups 3-12. They are characterized by partially filled d-orbitals, which allow for complex electron configurations and a variety of oxidation states. This flexibility is why transition metals are often used as catalysts and in alloy formation.

In Period 5 (n=5), the general electron configuration for transition elements is \[ \text{[Kr]} 4d^{1-10} 5s^{0-2} \]. For Period 6 (n=6), including the lanthanides, the configuration extends to \[ \text{[Xe]} 4f^{0-14} 5d^{1-10} 6s^{0-2} \]. The presence of empty or partially filled d-orbitals allows transition metals to form colored compounds, exhibit magnetic properties, and engage in unique bonding scenarios essential for industrial applications.

The Lanthanides, often referred to as rare earth elements, are a group of 15 metallic elements from lanthanum (La) to lutetium (Lu) in the Periodic Table. They are found in Period 6 and follow the f-block, characterized by filling the 4f orbitals. Lanthanides usually have similar chemical properties due to their similar electron configurations.

The general electron configuration for lanthanides is \[ \text{[Xe]} 4f^{0-14} 5d^{0-1} 6s^{2} \]. These elements are known for their magnetic and phosphorescent properties, making them essential in various high-tech applications, such as permanent magnets, phosphors in display screens, and catalysts. Understanding the properties and configurations of lanthanides is crucial for advancements in technology and materials science.