Skip to main content
Chemistry LibreTexts

4.8: Orbital Shape

  • Page ID
  • \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    \( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)

    ( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\id}{\mathrm{id}}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\kernel}{\mathrm{null}\,}\)

    \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\)

    \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\)

    \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)

    \( \newcommand{\vectorA}[1]{\vec{#1}}      % arrow\)

    \( \newcommand{\vectorAt}[1]{\vec{\text{#1}}}      % arrow\)

    \( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vectorC}[1]{\textbf{#1}} \)

    \( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)

    \( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)

    \( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)

    \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    The shell diagram example shown in the previous section, also known as a Bohr model, is a useful way to begin understanding how electrons fill orbitals. However, it is not quite this simple. Think back to section 4.4 on atomic theory, and how it discussed the electrons being in a cloud. Figure 4.4.1 showed no defined orbitals circling the nucleus in rings, but rather an overall area around the nucleus where the electrons were located. In an electron cloud model the electrons still exist within orbitals, but the shapes of the orbitals are more like a probability map of the location of electrons.

    The reason for this is that particles are always in motion. What this means, is that the electrons do not sit still. Instead, they move around throughout the electron cloud within the various orbitals. This can occur because of the negative charge the electrons carry. Like charges repel one another, and opposite charges attract. You can think of this like magnets. If you bring two north ends of a magnet near one another, they will push each other away (repel), however if you bring a north and a south end of a magnet together they will attract and stick to each other. Electrons will repel one another because of their like charges. Since they are always in motion, they continually will change paths to keep from interacting too closely with other electrons. At the same time, the electrons stay within the electron cloud of the atom because the positive charge of the protons in the nucleus attract the electrons.

    This isn’t to say electrons cannot leave an electron cloud. We have already discussed previously that ions form when an atom does not have equal numbers of electrons and protons. We will discuss this idea more later. For now, the focus is on understanding basic electron behavior.

    At this point you might be wondering about how orbitals are shaped if not in rings around the nucleus? The answer to this is that the shape of an orbital is dependent upon how many electrons can be held within the orbital. We already know that s-orbitals hold two electrons. The shape of this orbital is a sphere. The p-orbital (which holds a maximum of 6 electrons) is a peanut or dumbbell shape, and the d-orbital (holding a maximum of 10 electrons) is a cross peanut or cross dumbbell shape. Figure \(\PageIndex{1}\) shows the breakdown for basic orbital shapes, and the number of suborbitals within each orbital type. Each suborbital can hold a maximum of two electrons. You are only required to know s, p, and d orbital shapes for this class, but it is important to know other orbital shapes exist. Figure \(\PageIndex{1}\) below shows up through the f-orbital.

    Figure \(\PageIndex{1}\): Orbital shapes and their orientation along various axes.

    What you will notice is that for any orbital that holds more than two electrons there are various points in space in which those electrons may be held. Electrons arrange themselves into orbitals based on the amount of energy exists within the atom. The more electrons that exist within an atom, the more orbitals that exist. To fit multiple of each type of orbitals, the orbital space expands to overlap other orbitals. Since the electrons cannot interact, any electron in higher energy orbitals is pushed farther out from the nucleus. In the end, electrons closer to the nucleus are held more tightly by the positive charge of the protons. Meanwhile, electrons further from the nucleus have less attraction to the positively charged protons at the center of the atom. It is the electrons farthest from the nucleus that we call valence electrons. It is the valence electrons that interact with other atoms causing the formation of ionic and covalent compounds.


    You may find the video below from Khan Academy helpful.

     Video 4.8.1 walks through using the periodic table to write out ground state electron configurations for orbitals, instead of using the chart shown in section 4.7.


    4.8: Orbital Shape is shared under a Public Domain license and was authored, remixed, and/or curated by LibreTexts.

    • Was this article helpful?