Question

When writing electron configurations for elements in period 4 , why are electrons placed in the \(4 s\) sublevel before the \(3 d\) sublevel?

Short answer

Electrons are placed in the \(4 s\) sublevel before the \(3 d\) sublevel when writing electron configurations for elements in period 4 because the \(4 s\) sublevel has lower energy according to the Aufbau’s principle. Thus, it gets filled first.

Step-by-step solution

Understanding Electron Configuration

In electron configuration, we fill the electron shells and subshells in the order from lowest to highest energy following Aufbau’s principle. Energy levels determine where an electron resides in an atom. The lower the energy level, the closer the electron is to the nucleus and the less energy it has.

Understanding the 4s and 3d sublevels

According to quantum mechanics and some experimental observations, the \(4 s\) sublevel is of lower energy than the \(3 d\) sublevel. Hence, it becomes logical to fill up the \(4 s\) sublevel first before moving to the \(3 d\) sublevel.

Applying the Rule

Therefore, when writing electron configurations for elements in period 4, electrons are placed in the \(4 s\) sublevel before the \(3 d\) sublevel.

Key concepts

Aufbau principle

The Aufbau principle is a foundational concept used in chemistry to determine the electron configuration of atoms. This principle helps to explain why electrons fill orbitals in a specific order when an atom is at its ground state. According to the Aufbau principle, electrons occupy the lowest energy orbitals available before moving to higher energy orbitals. This minimizes the energy of the atom, leading to a more stable electron arrangement.

Consider this analogy: filling up seats in a bus. If all the seats are empty, you would naturally fill up the seats starting from the front because it's easier to hop on and off. Similarly, electrons fill the lowest energy level orbitals first. This principle is guided by the Pauli exclusion principle, which ensures that no two electrons can have the same set of quantum numbers, and Hund's rule, which prefers maximum multiplicity or unpaired spins first.

The determination of which orbital has lower energy follows the n+l rule, where n is the principal quantum number and l is the azimuthal quantum number. The sum (n+l) defines the energy order and electrons are added according to this until subshells reach their capacity.

Quantum Mechanics

Quantum mechanics provides the mathematical framework needed for understanding the behavior and arrangement of electrons around an atom. The electrons are located in atomic orbitals, which are regions of space where an electron is likely to be found.

Quantum mechanics introduces the concept of quantum numbers, which specify the properties and energies of atomic orbitals. These quantum numbers include:

• The principal quantum number (n), indicating the main energy level.
• The azimuthal quantum number (l), indicating the shape of the orbital.
• The magnetic quantum number (m_l), specifying the orientation of the orbital.
• The spin quantum number (m_s), denoting the spin direction of the electron.

These quantized energy levels lead to certain subshells being filled before others, due to variations in energy, such as the filling of the 4s subshell before the 3d subshell. This is because, despite being a higher principal quantum number, the 4s subshell has a lower energy state than the 3d subshell when an atom is in its lowest energy configuration.

Quantum mechanics thus provides a robust explanation for the observed electron configuration patterns and underlies the systematic filling order represented by the Aufbau principle.

Sublevels

In atomic structure, *sublevels* are divisions within principal energy levels that contain orbitals with specified shapes and capacities. These sublevels can house a particular number of electrons based on their shape. For instance, the s sublevel holds up to 2 electrons, the p sublevel accommodates 6, the d sublevel can contain 10, and the f sublevel can host 14 electrons.

The sequence of sublevel filling is crucial in determining an element's electron configuration, which can be seen in the periodic pattern in the periodic table itself. This arrangement of electrons across available sublevels determines the chemical properties and reactivity of each element. As elements progress across a period in the periodic table, electrons fill these sublevels in a manner that minimizes the atom's energy, as indicated by the Aufbau principle.

In the case of elements in period 4, electrons first fill the 4s sublevel before the 3d sublevel, even though the 3d sublevel is part of a lower principal energy level. This counterintuitive behavior occurs because of the slight energy differences between these sublevels as predicted by quantum mechanics. Understanding sublevel filling is essential for predicting electron configurations and deriving insights into the chemical behavior of different elements.