Equatorial and Axial Positions: Videos & Practice Problems

In molecular geometry, the arrangement of electron groups around a central atom affects the compound's stability and energy. For molecules with five or six electron groups, the positions of these groups are categorized as equatorial or axial. Equatorial positions are around the central atom's equator, while axial positions are above or below this plane. Generally, the more electronegative elements prefer the axial positions due to energy considerations, leading to greater stability. Lone pairs in six electron group systems are most stable in the axial positions, whereas in five electron group systems, they are most stable in the equatorial positions. A mnemonic to remember this is associating six o'clock with axial positions for six electron groups and five o'clock with equatorial positions for five electron groups, aiding in the visualization of the most stable molecular structure.

Equatorial vs. Axial Positions

Covalent compounds with five or six electron groups exhibit distinct spatial arrangements known as equatorial and axial positions for their surrounding elements. The equatorial position refers to the locations of surrounding elements situated around the "equator" of the molecular structure. For instance, in a compound with five electron groups, three of these groups are positioned along the equator, while in a compound with six electron groups, four are aligned along the equator.

In contrast, the axial or apical positions are those located above or below the equatorial plane. This spatial arrangement is crucial as it influences the repulsion between the surrounding elements. Increased repulsion can lead to a decrease in the overall energy of the compound, which is favorable in chemistry since lower energy states correspond to greater stability.

A key principle to remember is that more electronegative elements tend to occupy the axial positions rather than the equatorial ones. This preference is linked to the concepts of energy minimization and stability enhancement within the molecular structure. Understanding these arrangements and their implications is essential for grasping the stability of covalent compounds with multiple electron groups.

Molecular Shapes Example

To understand the molecular structure of phosphorus trichloride difluoride (PCl₃F₂), we start by identifying the central atom, which is phosphorus (P). As a group 5A element, phosphorus has five valence electrons and can form five bonding groups. In this case, it is bonded to two fluorine (F) atoms and three chlorine (Cl) atoms.

When drawing the Lewis dot structure, we place the least electronegative atom in the center, which is phosphorus. The arrangement of the surrounding atoms is influenced by the VSEPR (Valence Shell Electron Pair Repulsion) theory, which helps predict the geometry based on electron pair repulsion. Since phosphorus has five bonding groups, the molecular geometry is trigonal bipyramidal.

In this geometry, three of the bonding pairs will occupy the equatorial positions, while the remaining two will occupy the axial positions. To maximize stability and minimize repulsion, the more electronegative fluorine atoms should be placed in the axial positions. This is because axial positions experience greater repulsion due to their alignment with other axial bonds, making it preferable to place the more electronegative atoms there.

Fluorine, being in group 7A, has seven valence electrons and forms single bonds when bonded to phosphorus. Chlorine, also a group 7A element, similarly has seven valence electrons and will occupy the equatorial positions. Thus, the final structure of PCl₃F₂ will have two fluorine atoms in the axial positions and three chlorine atoms in the equatorial positions, resulting in a stable configuration.

This arrangement not only adheres to the principles of electronegativity but also optimizes the spatial distribution of the electron pairs around the phosphorus atom, leading to the most stable structure for phosphorus trichloride difluoride.

Lone Pair Positions

Lone pair positions play a crucial role in accurately representing molecular structures. When considering the arrangement of electron groups around a central atom, it is essential to minimize interactions between surrounding elements. In systems with six electron groups, lone pairs are most stable in the axial position, while in systems with five electron groups, they are best placed in the equatorial position.

To aid in remembering these orientations, a mnemonic involving a clock can be helpful. Visualize a clock where the hands indicate different times. For six electron groups, think of the time as 6 o'clock. In this scenario, the hands of the clock point straight up and down, mirroring the axial positions that are oriented above and below the equator of the molecular structure. Thus, 6 o'clock corresponds to the stability of lone pairs in axial positions.

Conversely, for five electron groups, the time is represented as 5 o'clock. Here, the clock hands do not point straight up and down, indicating that these positions are equatorial rather than axial. Therefore, in five electron group systems, lone pairs should be oriented in the equatorial positions to achieve the most accurate molecular representation.

In summary, remember that for six electron groups, lone pairs are stable in axial positions (6 o'clock), while for five electron groups, they are stable in equatorial positions (5 o'clock). This understanding is vital for drawing the best possible molecular structures.

Drawing Molecular Shape Example

To determine the molecular geometry of the ion SCl₃^-, we start by identifying sulfur as the central atom. Sulfur, located in group 6A of the periodic table, has 6 valence electrons. In this ion, sulfur forms single bonds with three chlorine atoms, which are in group 7A and each possess 7 valence electrons. Each chlorine uses one of its electrons to bond with sulfur, resulting in sulfur utilizing 3 of its 6 electrons for bonding.

After forming these bonds, sulfur has 3 remaining valence electrons. Since the ion carries a negative charge (-1), an additional electron is added to sulfur, bringing its total to 4 non-bonding electrons. This results in a total of 5 electron groups around the sulfur atom: 3 bonding pairs (from the S-Cl bonds) and 2 lone pairs.

To visualize the molecular geometry, we recognize that with 5 electron groups, we are dealing with a trigonal bipyramidal arrangement. However, the presence of lone pairs affects the final shape. The lone pairs will occupy the equatorial positions to minimize repulsion, while the chlorine atoms will occupy the axial and equatorial positions. Thus, one chlorine will be in an equatorial position, while the other two chlorines will be in axial positions, oriented vertically.

When drawing the structure, sulfur is placed at the center, with the two lone pairs in the equatorial positions and one chlorine atom also in an equatorial position. The other two chlorine atoms will be positioned axially, pointing straight up and down. It is essential to represent the ion with brackets and indicate the charge outside, resulting in the final structure of SCl₃^-.

Understanding the arrangement of lone pairs and bonding pairs is crucial for determining the best possible molecular structure. This knowledge allows for accurate predictions of molecular geometry based on the number of electron groups surrounding the central atom.

Draw the most likely shape for the following compound:XeF₄

Xenon tetrafluoride (XeF₄) is a molecular compound that exhibits a unique geometry due to the arrangement of its valence electrons. Xenon, a noble gas in group 8A, possesses 8 valence electrons. When forming bonds with four fluorine atoms, each requiring one electron to create a single bond, xenon utilizes 4 of its valence electrons for bonding.

After bonding with the four fluorine atoms, xenon retains 4 electrons, which are arranged as two lone pairs. The molecular geometry of xenon tetrafluoride can be understood through the VSEPR (Valence Shell Electron Pair Repulsion) theory, which predicts the spatial arrangement of electron groups around a central atom to minimize repulsion.

In the case of XeF₄, the molecule is classified as a 6 electron group system, consisting of 4 bonding pairs (from the Xe-F bonds) and 2 lone pairs. The most stable configuration places the lone pairs in the axial positions, while the fluorine atoms occupy the equatorial positions. This arrangement minimizes electron pair repulsion, leading to a square planar geometry.

Thus, the best representation of xenon tetrafluoride shows the four fluorine atoms positioned at the corners of a square, with the lone pairs situated above and below the plane of the square, effectively illustrating the molecule's square planar shape.

Draw and determine the geometry for the following molecule:Br₂CO

To determine the geometry of the molecule bromine dibromide carbon monoxide (Br₂CO), we start by identifying the central atom. Carbon, being the least electronegative element, is placed at the center of the structure. As a member of Group 4A, carbon has four valence electrons available for bonding.

In this molecule, the two bromine atoms, which are halogens from Group 7A, will each form a single bond with the carbon atom. Since carbon aims to achieve a total of four bonds, it will also need to form a double bond with the oxygen atom. Oxygen, found in Group 6A, possesses six valence electrons and will use two of these to create a double bond with carbon, while retaining two lone pairs of electrons.

The resulting structure features carbon at the center, bonded to three surrounding elements: two bromine atoms and one oxygen atom. This configuration does not conform to a typical five or six electron system, but it is essential to practice drawing such molecular structures to enhance understanding of molecular geometry and bonding.

How many lone pairs reside in the equatorial position of the KrCl₅⁺ ion. 

• A. 0
• B. 2
• C. 1
• D. 3
• E. 4

Answer: (A)

Here’s what students ask on this topic:

What are the axial and equatorial positions?

Axial and equatorial positions refer to the locations of atoms or groups in a molecule that has a three-dimensional geometric structure, particularly in cyclic structures like cyclohexane. Imagine a cyclohexane molecule as a chair, which is the most stable conformation for this six-membered ring.

In this "chair" form, there are two types of positions for the hydrogen atoms or substituents attached to the carbon atoms of the ring:

1. Axial positions: These are oriented parallel to the axis of the ring, which you can think of as being perpendicular to the imaginary plane that the ring lies in. If you visualize the ring flat on a table, axial substituents would either point straight up or straight down, like the legs and backrest of a chair.
2. Equatorial positions: These are located around the equator of the ring, almost in the plane of the ring itself, or slightly above and below it. They are angled outward and are more spread out, akin to the seat and armrests of a chair.

The difference in these positions is important in chemistry because they can influence the molecule's reactivity and stability, particularly when larger substituents are involved, due to steric hindrance and strain.

How can one determine equatorial and axial positions?

To determine equatorial and axial positions, it's essential to understand the geometry of the molecule in question. Typically, these terms are used in the context of trigonal bipyramidal molecular geometry, such as in the case of phosphorus pentachloride (PCl₅).

In a trigonal bipyramidal structure, there are five positions around the central atom. Three of these positions lie in the same plane and form an equilateral triangle; these are the equatorial positions. They are spaced at 120-degree angles from each other. The remaining two positions are above and below this plane, aligned along the vertical axis; these are the axial positions.

To determine which is which, look at the molecule's 3D structure:

• Equatorial positions: They are in the horizontal plane and tend to have more space around them.
• Axial positions: They are aligned vertically, along the central atom's axis, and typically have less space due to the linear arrangement with the central atom.

The distinction is important because it can affect the molecule's physical and chemical properties, such as dipole moment and reactivity.

What is the difference between axial and equatorial positions?

Axial and equatorial positions refer to the locations of atoms or groups in a molecule with a trigonal bipyramidal geometry, which is common in molecules with five atoms or groups around a central atom.

Axial positions are aligned parallel to the principal axis of the molecule. Imagine a molecule shaped like a seesaw; the axial positions are like the seats of the seesaw, located at the top and bottom, along the imaginary vertical axis passing through the center of the molecule. These positions are often less stable due to increased electron repulsion, as they are 90 degrees relative to more atoms compared to the equatorial positions.

Equatorial positions, on the other hand, lie along the plane perpendicular to the axial positions. In our seesaw analogy, these would be like the handles on the sides of the seesaw. There are three equatorial positions forming a triangle around the central atom in the middle plane of the molecule. These positions are usually more stable because they experience less electron repulsion, being 120 degrees apart from each other and 90 degrees from only two axial positions.