Orientations of D Orbitals - Online Tutor, Practice Problems & Exam Prep

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

In this video, we're going to take a look at the orientation of the orbitals. Now recall that an orbital is the area around nucleus where an electron is most likely to be found. And we're going to say that there are 5D orbitals with different orientations. These orientations themselves can be grouped into two sets, so we have orientations that lie in between our axes and those that lie on or along the axis.

Your professor, you might hear them say on the axes, and in another instance you might hear them say along the axis. These are synonymous with one another. They mean the same thing. Now, before we talk about these different types of orientations, remember we have our axes. We typically have our X axis oriented this way, our Y is oriented this way, and then our Z is vertical.

For the first three axes, we're going to say these are the ones that lie in between the axes. And in each of these three, we've highlighted a particular orbital lobe. For the first one, the highlighted lobe is up front and it's in between X&Y. That's why this orbital is designated as dxy. The next one is in between Y&Z, so this one is designated dyz. And then finally we have this one that's highlighted in between X&Z and that's why it's called dxz.

For the last two, they are on the axes or along the axes. And if you notice, you can see that they that the axes themselves kind of go through two of the lobes and those lobes lie on the X&Y axes. That's why this one is dx2−y2. And then finally this last one here it's along the Y axis and it pierces both all the way through. So that's why this one is dz2. So these are our five different types of orbitals, each with their orientation. The first three are in between the axes and the last two are on or along the axis.

Orientations of d Orbitals Example

Here it says an electron in which of the following orbital is the most likely to be found along the X&Y axis. So they said along the axis or on the axis. So there's only two possibilities and those are D sub X² -, y², or DZ squared.

We can see from the designations here we're talking about along the X&Y axis. The first one is the one that contains those two variables of X&Y. So here that would mean that answer or option D would have to be our correct answer.

The first three do not work because they are in between the axes and then DZ squared. Here we're talking about the one that's predominantly on X&Y and that have to be option D over option East.

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

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.