Orientations of D Orbitals: Videos & Practice Problems

Understanding the orientation of d orbitals is crucial in chemistry, as it describes where electrons are most likely to be found around an atom's nucleus. There are five d orbitals, each with a unique orientation in three-dimensional space relative to the x, y, and z axes. Three of these orbitals (dxy, dyz, dxz) are oriented between the axes, while the remaining two (dx2-y2 and dz2) lie on or along the axes. The dx2-y2 orbital has lobes lying on the x and y axes, whereas the d_z2 orbital has a distinctive alignment along the z-axis, piercing through the nucleus. This spatial arrangement of d orbitals is fundamental to understanding molecular bonding and the electronic structure of transition metals.

d Orbital Orientations

Understanding the orientation of d orbitals is essential in grasping the behavior of electrons in atoms. An orbital represents the region around the nucleus where an electron is most likely to be found. There are five distinct d orbitals, each with unique orientations that can be categorized into two groups: those that lie between the axes and those that lie on or along the axes.

To visualize these orientations, consider the three-dimensional coordinate system with the x, y, and z axes. The first three d orbitals, which lie between the axes, are designated as follows:

• d_xy: This orbital is oriented between the x and y axes, with a highlighted lobe positioned in that space.
• d_yz: This orbital is situated between the y and z axes, featuring a lobe that occupies that region.
• d_xz: This orbital lies between the x and z axes, with its lobe highlighted in that area.

The remaining two d orbitals are aligned with the axes:

• d_x²-y²: This orbital has lobes that extend along the x and y axes, effectively intersecting them.
• d_z²: This orbital is oriented along the z axis, with lobes that penetrate through this axis.

In summary, the five d orbitals—d_xy, d_yz, d_xz, d_x²-y², and d_z²—each have specific orientations that play a crucial role in the chemical bonding and properties of elements. The first three orbitals are positioned between the axes, while the last two are aligned with them, highlighting the diverse spatial arrangements of electrons in d orbitals.

Orientations of d Orbitals Example

In quantum mechanics, the behavior of electrons in atoms is described by orbitals, which are regions in space where there is a high probability of finding an electron. When considering the likelihood of an electron being found along the x and y axes, we focus on specific d orbitals.The d orbital designated as \(d_{x^2-y^2}\) is particularly relevant here. This orbital is oriented along the x and y axes, making it the most probable location for an electron when we are looking for positions along these axes. In contrast, the \(d_{z^2}\) orbital is primarily oriented along the z-axis and does not contribute to the probability of finding an electron along the x and y axes.The other d orbitals, such as \(d_{xy}\), \(d_{xz}\), and \(d_{yz}\), are situated between the axes and do not align directly with the x and y axes. Therefore, they are not suitable choices for this scenario.In summary, when determining the most likely orbital for an electron to be found along the x and y axes, the \(d_{x^2-y^2}\) orbital is the correct answer, as it is specifically designed to reflect that orientation.

In which of the following orbitals an electron is the most likely to be found along the z axis?

• A. d_xy
• B. d_yz
• C. d_xz
• D. d_x²–y²
• E. d_z²

Answer: (E)

Here’s what students ask on this topic:

What is the maximum amount of electrons that all d orbital orientations can hold?

The d orbitals can hold a maximum of 10 electrons. In an atom, there are five d orbitals—d_xy, d_yz, d_xz, d_x²-y², and d_z²—and each orbital can contain up to two electrons. These orbitals are part of the third shell and higher (n≥3) of an atom. Electrons in these orbitals are filled according to Hund's rule, which states that electrons will fill empty orbitals singly before pairing up. The electron configuration for d orbitals starts to fill after the s orbitals of the same principal quantum number (n-1). For example, in the fourth period of the periodic table, the 3d orbitals begin to fill after the 4s orbital is filled. The filling of these orbitals is important for understanding the chemical behavior of transition metals, where the d orbitals play a crucial role.

What do d orbitals look like?

D orbitals have a more complex shape than the simpler s and p orbitals. There are five d orbitals, designated as d_xy, d_yz, d_zx, d_x²-y², and d_z². The first three (d_xy, d_yz, d_zx) have similar cloverleaf shapes with four lobes each; they lie in the planes indicated by their subscripts. For example, the d_xy orbital lies in the xy-plane. The d_x²-y² orbital also has a cloverleaf shape but with lobes lying along the x and y axes.

The d_z² orbital is unique; it has two lobes along the z-axis and a doughnut-shaped region in the xy-plane. It is sometimes depicted with the lobes shaped like a p orbital with a ring around the middle. These shapes are important because they determine how d orbitals overlap with other orbitals and thus, how they participate in bonding within molecules.