Problem 58 - Chemistry The Central Science

Question

Give the values for \(n, l,\) and \(m_{l}\) for (a) each orbital in the 2\(p\) subshell, (b) each orbital in the 5\(d\) subshell.

Step-by-Step Solution

Answer

(a) For the 2p subshell: \(n = 2\), \(l = 1\), and \(m_l\) values are -1, 0, and +1. (b) For the 5d subshell: \(n = 5\), \(l = 2\), and \(m_l\) values are -2, -1, 0, +1, and +2.

Step 1: Find the azimuthal quantum number (l) for the p subshell

The azimuthal quantum number (l) is 0 for the s subshell, 1 for the p subshell, 2 for the d subshell, and so on. Therefore, for the p subshell, l = 1.

Step 2: Find the magnetic quantum number (m_l) values for the 2p subshell

Knowing that l = 1, the magnetic quantum numbers (m_l) can take values from -l to +l. Therefore, in the 2p subshell, m_l can have the values -1, 0, +1. There are three possible orbitals in the 2p subshell. So, for the 2p subshell, we have: - n = 2 - l = 1 - m_l: -1, 0, +1 (b) The 5d subshell: As given, the energy level (n) is 5, and the subshell is d. So, our task here is to find the value of l for the d subshell and the values of m_l for each orbital in the 5d subshell.

Step 3: Find the azimuthal quantum number (l) for the d subshell

For the d subshell, l = 2.

Step 4: Find the magnetic quantum number (m_l) values for the 5d subshell

Knowing that l = 2, the magnetic quantum numbers (m_l) can take values from -l to +l. Therefore, in the 5d subshell, m_l can have the values -2, -1, 0, +1, +2. There are five possible orbitals in the 5d subshell. So, for the 5d subshell, we have: - n = 5 - l = 2 - m_l: -2, -1, 0, +1, +2

Key Concepts

Azimuthal Quantum Number

The azimuthal quantum number, denoted by the letter \( l \), is a pivotal component in understanding how electrons are arranged within an atom's electron shells. It plays a fundamental role in defining the shape of the atomic orbitals. Each value of \( l \) corresponds to a specific subshell, which contains orbitals of a particular shape. For instance:

• The \( s \) subshell has \( l = 0 \).
• The \( p \) subshell is assigned \( l = 1 \).
• The \( d \) subshell corresponds to \( l = 2 \).
• The \( f \) subshell has \( l = 3 \).

The possible values for \( l \) range from 0 to \( n-1 \), where \( n \) is the principal quantum number, or the energy level of the electron. Understanding \( l \) helps us predict the types of orbitals that exist within each energy level.
For example, in a 2\(p\) subshell, \( n = 2 \) and \( l = 1 \), indicating the presence of p orbitals.

Magnetic Quantum Number

The magnetic quantum number, denoted as \( m_{l} \), focuses on describing the orientation of an orbital within a subshell. This number arises from the quantum mechanical nature of electrons within the atom. It emerges as electrons possess angular momentum when revolving in their orbitals.

• The values of \( m_{l} \) range from \(-l \) to \(+l \), including zero. Hence, each \( l \) value leads to multiple \( m_{l} \) values, defining the number of orientations of an orbital within a particular subshell.
• For the \( p \) subshell, where \( l = 1 \), \( m_{l} \) can be -1, 0, or +1, corresponding to the three \( p \) orbitals typically seen.
• For the \( d \) subshell with \( l = 2 \), \( m_{l} \) can be -2, -1, 0, +1, or +2, resulting in five distinct orientations for the \( d \) orbitals.

This quantum number is crucial because it explains the spatial distribution of orbitals associated with magnetic properties.

Atomic Orbitals

Atomic orbitals are regions within an atom where electrons are most likely to be found. They provide a probabilistic view of electron locus rather than exact paths, thanks to the principles of quantum mechanics. Each type of orbital (s, p, d, f) offers unique shapes and characteristics.

• s orbitals: Spherical in shape. They can hold up to two electrons. For example, the 1\(s\) and 2\(s\) orbitals are simple spheres centered around the nucleus.
• p orbitals: Dumbbell-shaped and oriented in three possible directions (x, y, z), as indicated by \( m_{l} \) values: -1, 0, +1. They appear, beginning with the second energy level (2\(p\)).
• d orbitals: More complex in shape, often depicted as clovers or having more lobes. Found from the third energy level onward (3\(d\)), with five orientations due to \( m_{l} \): -2, -1, 0, +1, and +2.

These orbitals accommodate electrons according to the principles of quantum mechanics, namely the Pauli exclusion principle and Hund’s rule. By exploring the interrelationship of quantum numbers and atomic orbitals, one can gain a profound insight into the structural foundation of matter.