Question

How many electrons can be accommodated in a p-orbital? (a) 6 electrons (b) 2 electrons (c) 4 electrons (d) None of these

Short answer

The p-orbital can accommodate 6 electrons.

Step-by-step solution

Understanding Electron Accommodation in Orbitals

Each p-orbital can hold a maximum number of electrons. Orbitals are regions of space around an atom's nucleus where electrons are likely to be found. Electron configuration rules help determine this capacity.

Exploring p-orbital Characteristics

A p-orbital can hold a specific number of electrons as part of an atom's electron configuration, contributing to the chemical properties and behavior of elements.

Analyzing Electron Spin and Capacity

According to the Pauli Exclusion Principle, each orbital can hold a maximum of two electrons, which must have opposite spins. For p-orbitals, which consist of three spatial orientations, this principle allows a total of six electrons.

Applying the Principle to p-orbital

The three p-orbitals ( px, py, pz) accommodate 2 electrons each. Hence, the total electron capacity for p-orbitals is calculated as follows: 3 ext{ orbitals} imes 2 ext{ electrons/orbital} = 6 ext{ electrons}.

Key concepts

Pauli Exclusion Principle

The Pauli Exclusion Principle is a key concept in understanding how electrons fill atomic orbitals. It states that no two electrons in an atom can have the same set of all four quantum numbers.
This means that in any given orbital, two electrons can coexist only if they have opposite spins. Spins are represented as either +1/2 or -1/2. This principle ensures that each electron remains uniquely distinguishable within an atom.
For example, in the case of a p-orbital:

• There are three possible orientations: \(p_x\), \(p_y\), and \(p_z\).
• Each orientation can hold two electrons with opposite spins.
• Thus, combining these features allows a total of six electrons to occupy the p-orbital set.

Understanding the Pauli Exclusion Principle helps clarify why electrons fill specific orbitals in their unique ways, contributing to the arrangement and overall behavior of elements on the periodic table.

Electron Configuration Rules

Electron configuration rules are guidelines that help predict how electrons distribute among the orbitals of an atom. These rules ensure that electrons occupy orbitals in the most energetically favorable arrangement.
One of the main rules is the Aufbau Principle, which suggests that electrons fill orbitals starting from the lowest energy level moving up. Besides, the Hund's Rule stipulates that electrons will first fill all orbitals with the same energy (degenerate orbitals) one by one before pairing up.
For instance:

• In the p-block of the periodic table, after the s-orbital is filled, electrons move to the p-orbitals.
• They first occupy each of the three p-orbitals singly before pairing occurs.
• This behavior ensures the electron configuration reflects the minimum energy state of an atom.

Utilizing electron configuration rules enables us to write out configurations like \([Ne] 3s^2 3p^6\) for atoms like sulfur, providing insights into their chemical properties and reactions.

Orbital Characteristics

Understanding orbital characteristics, specifically for p-orbitals, is essential in grasping how atoms bond and react with one another. Orbitals are essentially regions in space where there is a high probability of finding electrons.
The p-orbitals have distinct features:

• There are three p-orbitals: \(p_x\), \(p_y\), and \(p_z\).
• These orbitals are oriented at right angles to each other, aligning along the x, y, and z axes.
• Each orbital has a dumbbell shape, which influences their bonding characteristics in molecules.

The spatial arrangement of p-orbitals is critical in determining the geometry of molecules and their magnetic properties. It also plays a significant role in phenomena such as covalent bonding by enabling overlap with orbitals from other atoms to form bonds.
Grasping these characteristics enhances comprehension of how elements interact at an atomic level, influencing both physical and chemical behaviors.