Chemistry·Explained

Voids in Close Packed Structures — Explained

Q: What is the significance of the radius ratio rule in relation to voids?

The radius ratio rule ($r_+/r_-$) is crucial for predicting the stability and coordination number of ionic compounds where smaller cations occupy voids formed by larger anions. It defines the minimum and maximum size ratio for a cation to fit stably into a particular void (e.g., tetrahedral or octahedral) without causing significant lattice distortion. If the cation is too small, it 'rattles'; if too large, it forces the anions apart, leading to instability or a different coordination geometry.

NEET UG

Updated 22 Mar 2026

Detailed Explanation

The concept of voids is central to understanding the structure and properties of crystalline solids, particularly those with close-packed arrangements. When identical spheres (atoms or ions) pack together as efficiently as possible, they form either a hexagonal close-packed (HCP) or a cubic close-packed (CCP), also known as face-centered cubic (FCC), structure.

In both these arrangements, the packing efficiency is approximately 74%, meaning 26% of the total volume is empty space. These empty spaces are not random but are specific interstitial sites with defined geometries, known as voids.

Conceptual Foundation: Close Packing and Void Formation

Close packing arises from the tendency of atoms or ions to maximize their coordination number and minimize empty space. Imagine a layer of spheres arranged in a 2D hexagonal pattern. When a second layer is placed over the first, the spheres of the second layer can occupy the depressions (hollows) of the first layer.

There are two types of depressions in the first layer, let's call them 'a' and 'b'. If the second layer spheres occupy 'a' type depressions, then when a third layer is added, it can either occupy the depressions directly above the first layer spheres (forming an 'aba' pattern, which is HCP) or occupy the depressions that are not directly above the first layer spheres (forming an 'abc' pattern, which is CCP/FCC).

Regardless of the stacking sequence (HCP or CCP), specific types of voids are generated.

Key Principles: Types of Voids

There are two primary types of voids in close-packed structures:

Tetrahedral Voids (TVs): — These voids are formed when four spheres surround a central empty space, with their centers forming a regular tetrahedron. The coordination number of a particle occupying a tetrahedral void is 4. This means an atom or ion sitting in a tetrahedral void is in contact with four surrounding spheres. These voids are relatively small.

Octahedral Voids (OVs): — These voids are formed when six spheres surround a central empty space, with their centers forming a regular octahedron. The coordination number of a particle occupying an octahedral void is 6. An atom or ion in an octahedral void is in contact with six surrounding spheres. These voids are larger than tetrahedral voids.

Derivations: Number and Location of Voids

Let's consider a close-packed structure formed by 'N' spheres. In any close-packed arrangement (HCP or CCP/FCC):

The number of octahedral voids is equal to the number of spheres, i.e., N.

The number of tetrahedral voids is twice the number of spheres, i.e., 2N.

This relationship is crucial for determining the stoichiometry of ionic compounds.

Location of Voids in FCC (Cubic Close Packing):

In an FCC unit cell, the effective number of spheres (atoms) is 4 (8 corners $\times$ 1/8 + 6 faces $\times$ 1/2 = 1 + 3 = 4).

Octahedral Voids: — According to the 1:1 relationship, there should be 4 octahedral voids per FCC unit cell. Their locations are:

One at the body center of the unit cell (contributes 1). Twelve at the edge centers (each shared by 4 unit cells, so $12 \times 1/4 = 3$ ). * Total OVs = $1 + 3 = 4$ . * An octahedral void at the body center is surrounded by the six face-centered atoms. An octahedral void at an edge center is surrounded by two corner atoms and four face-centered atoms from adjacent faces.

Tetrahedral Voids: — According to the 2:1 relationship, there should be

2 \times 4 = 8

tetrahedral voids per FCC unit cell. Their locations are:

* Eight tetrahedral voids are located at the body diagonals, one-fourth of the way from each corner. Each body diagonal has two tetrahedral voids. Since there are four body diagonals in a cube, there are $4 \times 2 = 8$ tetrahedral voids. Each of these voids is entirely within the unit cell and is surrounded by one corner atom and three face-centered atoms.

Radius Ratio Rule and Void Occupancy:

The stability of an ionic compound where smaller cations occupy voids formed by larger anions depends on the radius ratio ( $r_+/r_-$ ), where $r_+$ is the radius of the cation and $r_-$ is the radius of the anion. For a cation to fit perfectly into a void without distorting the lattice, its radius must be within a specific range relative to the anion's radius.

For Tetrahedral Voids: — The ideal radius ratio for a cation to fit into a tetrahedral void is

r_+/r_- = 0.225

. If the cation is too small, it rattles; if it's too large, it pushes the anions apart, distorting the structure. The range for stable tetrahedral coordination is

0.225 \le r_+/r_- < 0.414

For Octahedral Voids: — The ideal radius ratio for a cation to fit into an octahedral void is

r_+/r_- = 0.414

. The range for stable octahedral coordination is

0.414 \le r_+/r_- < 0.732

Derivation of Radius Ratio for Octahedral Void:

Consider an octahedral void at the body center of an FCC unit cell. The anions (spheres) are at the face centers. Let 'R' be the radius of the anion and 'r' be the radius of the cation occupying the void.

The anions touch along the face diagonal. The length of the face diagonal is $4R$ . Also, the face diagonal is $\sqrt{2}a$ , where 'a' is the edge length of the unit cell. So, $4R = \sqrt{2}a \implies a = 4R/\sqrt{2} = 2\sqrt{2}R$ .

Now, consider the edge of the unit cell. The cation in the octahedral void at the body center is in contact with the anions at the face centers. The distance from the body center to a face center is $a/2$ .

This distance is also $R+r$ . So, $R+r = a/2$ . Substitute 'a': $R+r = (2\sqrt{2}R)/2 = \sqrt{2}R$ . $r = \sqrt{2}R - R = R(\sqrt{2}-1) = R(1.414-1) = 0.414R$ . Thus, $r/R = 0.414$ .

Derivation of Radius Ratio for Tetrahedral Void:

Consider a tetrahedral void located at $1/4$ th of the body diagonal from a corner in an FCC unit cell. The four spheres forming the tetrahedron are one corner atom and three face-centered atoms. The distance from the corner to the center of the tetrahedral void is $\sqrt{3}a/4$ .

This distance is also $R+r$ . The distance between the centers of two touching spheres (anions) along the face diagonal is $2R$ . The face diagonal is $\sqrt{2}a$ . So, $2R = \sqrt{2}a/2$ (distance from corner to face center) is incorrect.

Let's use a simpler approach. Consider a small cube of side length $a/2$ within the FCC unit cell. A corner of this small cube is occupied by a corner atom of the FCC cell, and the body center of this small cube is a tetrahedral void.

The distance from the corner of the small cube to its body center is $\sqrt{3}(a/2)/2 = \sqrt{3}a/4$ . This distance is $R+r$ . We know $a = 2\sqrt{2}R$ from the FCC structure. So, $R+r = \sqrt{3}(2\sqrt{2}R)/4 = \sqrt{6}R/2 \approx 1.

225R $. Therefore,$ r = 1.225R - R = 0.225R $. Thus,$ r/R = 0.225$.

Real-World Applications:

Voids are fundamental to the structure of many ionic compounds. For example:

In NaCl, the larger Cl

^-

ions form an FCC lattice, and the smaller Na

^+

ions occupy all the octahedral voids. This explains its 1:1 stoichiometry and rock-salt structure.

In ZnS (zinc blende), the S

^{2-}

ions form an FCC lattice, and the Zn

^{2+}

ions occupy half of the tetrahedral voids. This also gives a 1:1 stoichiometry.

In CaF

_2

(fluorite structure), Ca

^{2+}

ions form an FCC lattice, and F

^-

ions occupy all the tetrahedral voids. This leads to a 1:2 stoichiometry.

Common Misconceptions:

Confusing number of voids with number of atoms: — Students often forget the 2N and N relationship. It's crucial to remember that for 'N' spheres, there are 2N tetrahedral and N octahedral voids.

Incorrectly locating voids: — Visualizing the exact positions of voids in a 3D unit cell (especially FCC) can be challenging. Practice with diagrams is essential.

Mixing up radius ratio ranges: — Memorizing the specific radius ratio ranges for different void types is important for predicting coordination numbers and structures.

Assuming all voids are occupied: — In many ionic compounds, only a fraction of the available voids are occupied, which determines the compound's stoichiometry.

NEET-Specific Angle:

NEET questions often test the ability to:

Calculate the number of tetrahedral and octahedral voids in a given unit cell (e.g., FCC, HCP).

Determine the formula of an ionic compound given the arrangement of anions and the fraction of voids occupied by cations.

Identify the location of voids within an FCC unit cell.

Apply the radius ratio rule to predict the type of void occupied or the coordination number.

Relate packing efficiency to the presence of voids.

Often confused with

Side-by-side differences the NEET paper likes to test.

Voids in Close Packed Structures vs Octahedral Voids

Aspect	Voids in Close Packed Structures	Octahedral Voids
Geometry	Tetrahedral (surrounded by 4 spheres)	Octahedral (surrounded by 6 spheres)
Coordination Number	4	6
Relative Size	Smaller	Larger
Number per 'N' spheres (close-packed)	2N	N
Location in FCC unit cell (4 effective atoms)	8 voids, along body diagonals (1/4th and 3/4th from corners)	4 voids, at body center and edge centers
Ideal Radius Ratio ($r_+/r_-$)	$0.225$	$0.414$

Tetrahedral voids are smaller, more numerous (2N per N spheres), and have a coordination number of 4, typically accommodating smaller cations with a radius ratio around 0.225. Octahedral voids are larger, less numerous (N per N spheres), and have a coordination number of 6, suitable for larger cations with a radius ratio around 0.414. Their distinct geometries and locations within the crystal lattice are crucial for determining the structure and stoichiometry of ionic compounds.

Why it is tested: For NEET, understanding the differences between tetrahedral and octahedral voids is fundamental. Questions frequently test their relative numbers, locations in common unit cells (like FCC), and the application of radius ratio rules to predict which void type a cation will occupy. This knowledge is directly applicable to determining the chemical formulas of ionic solids and understanding their crystal structures.

Questions students ask

5 answered on this topic.

What is the fundamental difference between tetrahedral and octahedral voids?

The fundamental difference lies in their geometry and the number of surrounding spheres. A tetrahedral void is surrounded by four spheres, forming a tetrahedron, and thus has a coordination number of 4. An octahedral void, conversely, is surrounded by six spheres, forming an octahedron, and has a coordination number of 6. Consequently, octahedral voids are generally larger than tetrahedral voids, which is reflected in their respective radius ratio ranges for stable occupation by smaller ions.

How many tetrahedral and octahedral voids are present per atom in a close-packed structure?

For every 'N' spheres (atoms or ions) in a close-packed arrangement (either HCP or CCP/FCC), there are always 'N' octahedral voids and '2N' tetrahedral voids. This 1:1:2 ratio (spheres:octahedral voids:tetrahedral voids) is a universal principle for close-packed structures and is critical for determining the stoichiometry of ionic compounds where smaller ions occupy these interstitial sites.

Where are the tetrahedral and octahedral voids located in an FCC unit cell?

In an FCC unit cell, which has 4 effective atoms, there are 4 octahedral voids and 8 tetrahedral voids. The 4 octahedral voids are located at the body center (1 void) and at the center of each of the 12 edges (each contributing 1/4th, totaling 3 voids). The 8 tetrahedral voids are located along the body diagonals, with two voids on each of the four body diagonals, positioned at 1/4th and 3/4th of the diagonal length from the corner.

What is the significance of the radius ratio rule in relation to voids?

The radius ratio rule ( $r_+/r_-$ ) is crucial for predicting the stability and coordination number of ionic compounds where smaller cations occupy voids formed by larger anions. It defines the minimum and maximum size ratio for a cation to fit stably into a particular void (e.g., tetrahedral or octahedral) without causing significant lattice distortion. If the cation is too small, it 'rattles'; if too large, it forces the anions apart, leading to instability or a different coordination geometry.

Can all voids in a close-packed structure be occupied by smaller ions?

No, not necessarily. While voids provide potential sites for smaller ions, the actual occupancy depends on several factors, including the stoichiometry of the compound, the relative sizes of the ions (radius ratio), and charge balance. For instance, in NaCl, all octahedral voids are occupied, but in ZnS, only half of the tetrahedral voids are occupied. The specific fraction of occupied voids dictates the chemical formula of the resulting ionic solid.