Question

In the transition elements, the incoming electron occupies $(n-1)\mathrm{d}$ sublevel in preference to (1) $\mathrm{np}$ (2) $\mathrm{ns}$ (3) $(n-1)\mathrm{d}$ (4) $(n+1)\mathrm{s}$

Short answer

(3) $(n-1)\mathrm{d}$

Step-by-step solution

Understanding Transition Elements

Transition elements are elements whose atoms have an incomplete d subshell or can give rise to cations with an incomplete d subshell.

Electron Configuration Basics

In electron configuration, electrons fill orbitals in a specific order: first the 1s orbital, then 2s, then 2p, and so on. The general order is determined by the increasing energy levels of the orbitals.

Identifying the Sublevel for Transition Elements

For transition elements, electrons start to fill the $(n-1)\mathrm{d}$ sublevel after the $n\mathrm{s}$ orbital. This is because the $(n-1)\mathrm{d}$ sublevel is lower in energy than $n\mathrm{p}$ and $(n+1)\mathrm{s}$ sublevels.

Preferred Sublevel for Incoming Electrons

Therefore, for transition elements, after filling the $n\mathrm{s}$ orbital, the incoming electron occupies the $(n-1)\mathrm{d}$ sublevel in preference to other sublevels like $\mathrm{np}$ or $n\mathrm{p}$.

Key concepts

Transition Elements

Transition elements are fascinating components of the periodic table. Found in the d-block, they are defined by having incomplete d sublevels in their ground state or in their common oxidation states. This gives them unique properties like the ability to form colorful compounds, multiple oxidation states, and to be excellent conductors of electricity.

Transition elements are typically metals and are found in the central block of the periodic table, ranging from group 3 to group 12. Some well-known transition elements include iron (Fe), copper (Cu), and gold (Au).

These elements play a crucial role in chemistry and industry because of their ability to easily form compounds with other elements. They are also essential in biological processes; for instance, iron is a key component of hemoglobin in blood.

Electron Configuration

The concept of electron configuration is fundamental to understanding how elements behave and interact with one another. Electron configuration refers to the distribution of electrons in an atom's orbitals. The arrangement follows the Pauli Exclusion Principle, the Aufbau Principle, and Hund's Rule.

Electrons fill orbitals starting from the lowest energy level to the highest. This pattern of filling follows the sequence: 1s, 2s, 2p, 3s, 3p, 4s, and then 3d, and so on. It is essential to remember the order of filling, as it helps in predicting the chemical properties and behavior of the elements.

For transition elements, the electron configuration can be quite complex because electrons fill the 3d sublevel after the 4s orbital. For example, the electron configuration of Iron (Fe) is [Ar] 3d⁶ 4s².

(n-1)d Sublevel

A crucial aspect of transition elements is the filling of the (n-1)d sublevel. For transition elements, the outermost electrons begin to fill the d orbitals one energy level below the current period. This means that for an element in the fourth period (n=4), the 3d sublevel (n-1) is filled after the 4s orbital.

The reason for this is primarily due to the energy hierarchy of the orbitals. The (n-1)d sublevel is generally higher in energy compared to the ns orbital but lower compared to the np and (n+1)s orbitals. Thus, after filling the ns orbital, the (n-1)d sublevel is the next preferred orbital for the incoming electron.

In other words, for transition elements, once the ns orbital is occupied, the subsequent electrons will preferentially fill the (n-1)d sublevel, providing them with their unique properties and behavior. This accounts for why the transition elements exhibit similarities in chemical properties despite differing atomic numbers.