Void Meaning
The word "void" refers to gaps between constituent particles. In a densely packed structure, voids refer to the space between constituent particles (voids in chemistry). Solids can be packaged in one of three ways: one-dimensional (1D), two-dimensional (2D), or three-dimensional (3D).
When atoms are arranged in square close packing of hexagonal close packing, we see empty spaces between them in 2-dimensional structures.
These empty spaces are known as voids, and in hexagonal packing, these voids have triangular shapes and are referred to as triangular voids. Thus, the vacant spaces in a closely packed arrangement are called voids.
Tetrahedral and Octahedral Voids
In hexagonal packing, these triangular voids are seen in two different orientations. The apex of the triangle in one row points upward, while the apex of the triangle in the other row points downward.
In the three-dimensional structure, about 26% of total space is empty and not occupied by spheres in both CCP and HCP near packing in solids. Interstitial voids, interstices, or gaps are the names given to these empty spaces. The above voids in solids are proportional to the number of spheres present.
In a three-dimensional structure, there are two types of interstitial voids:
Tetrahedral Voids: In a cubic close-packed structure, the second layer's spheres are above the first layer's triangular voids. Each sphere touches the first layer's three spheres. It forms a tetrahedron by joining the centers of these four spheres, and the space created by joining the centers of these spheres forms a tetrahedral void. In a closed packed structure, the number of tetrahedral voids is two times the number of spheres. Let the number of spheres be n. Then the number of tetrahedral voids will be 2n.
Octahedral Voids: Adjacent to tetrahedral voids you can find octahedral voids. Octahedral voids are located next to tetrahedral voids. So now, what are Octahedral Voids? When the triangular voids of the first layer coincide with the triangular voids of the layer above or below it, we get a void that is formed by enclosing six spheres. This vacant space formed by combining the triangular voids of the first layer and that of the second layer is called Octahedral Voids. Octahedral Voids refer to the space created by combining the triangular voids of the first and second layers. If the number of spheres in a close-packed structure is n, then the number of octahedral voids will be n.
Number of Voids
The number of these two types of voids depends on the number of closed-packed spheres.
If the number of closed packed spheres is N, then
The octahedral void be N
The tetrahedral void be 2N
What is the Primary Difference between Tetrahedral and Octahedral Voids?
Tetrahedral voids are unoccupied empty spaces present in substances having a tetrahedral crystal system. Octahedral voids are unoccupied empty spaces present in substances having an octahedral crystal system. It can be found in substances having a tetrahedral arrangement in their crystal system. A tetrahedral void is a simple triangular void in a crystal and is surrounded by four spheres arranged tetrahedrally around it. On the other hand, an octahedral void is a double triangular void with one triangle vertex upwards and the other triangle vertex downwards and is surrounded by six spheres.
Difference between Tetrahedral and Octahedral Voids
In-depth Concept of Void
Voids mean gaps between the constituent particles. Voids in solid states mean the vacant space between the constituent particles in a closed-packed structure. Close packing in solids can be generally done in three ways: 1D close packing, 2D close packing, and 3D close packing.
In 2 dimensional structures when the atoms are arranged in square close packing and hexagonal close packing, we see empty spaces left over between the atoms. These empty spaces are called voids and in the case of hexagonal packing, these voids are in triangular shapes and are known as the triangular voids.
Did You Know?
The unit cell, or building block of a crystal, is the smallest repeating unit of the crystal lattice.
The identical unit cells are described in such a way that they fill the available space without overlapping. A crystal lattice is a three-dimensional arrangement of atoms, molecules, or ions within a crystal. It comprises a large number of unit cells. Per lattice point is occupied by one of the three constituent particles.
Numerous unit cells together make a crystal lattice. Constituent particles like atom, molecules are also present. Each lattice point is occupied by one of these particles.
Primitive Cubic Unit Cell
Body-Centered Cubic Unit Cell
Face Centered Cubic Unit Cell
FAQs on Void
1. What are voids in the context of solid-state chemistry?
In solid-state chemistry, voids (also known as interstitial sites or holes) are the empty spaces left between the constituent particles (atoms, molecules, or ions) in a close-packed crystal structure. Even in the most efficient packing arrangements, only about 74% of the space is occupied, with the remaining 26% consisting of these voids. They are crucial for understanding the properties and structures of crystalline solids.
2. What is the main difference between a tetrahedral and an octahedral void?
The primary differences between tetrahedral and octahedral voids lie in their formation and geometry:
- Formation: A tetrahedral void is formed by and surrounded by four spheres. An octahedral void is formed by and surrounded by six spheres.
- Coordination Number: The coordination number of a particle in a tetrahedral void is 4, while for an octahedral void, it is 6.
- Shape: A tetrahedral void is the space enclosed by four spheres arranged in a tetrahedral geometry. An octahedral void is the space created between two opposing sets of triangular voids from adjacent layers.
3. How is a tetrahedral void formed in a crystal lattice?
A tetrahedral void is formed when a sphere from one layer of a close-packed structure is placed directly on top of a triangular depression formed by three touching spheres in the layer below it. The empty space enclosed by these four spheres is the tetrahedral void. Connecting the centres of these four spheres creates a geometric shape called a tetrahedron.
4. How is an octahedral void formed in a crystal lattice?
An octahedral void is formed when the triangular void of one layer aligns perfectly with an opposing triangular void from the layer directly above or below it. This space is enclosed by a total of six spheres—three in the bottom layer and three in the top layer. Connecting the centres of these six spheres forms a geometric shape known as an octahedron.
5. What is the relationship between the number of close-packed particles and the number of voids?
There is a fixed mathematical relationship between the number of voids and the number of particles in a close-packed structure. If the number of close-packed spheres is represented by N, then:
- The number of octahedral voids is equal to N.
- The number of tetrahedral voids is equal to 2N.
- Therefore, the total number of voids in the lattice is 3N.
6. Why are there always twice as many tetrahedral voids as close-packed spheres (N)?
This 2:1 ratio arises from the geometry of close-packing. Each sphere in a close-packed structure has one triangular void directly above it and one directly below it. When the next layers are added, a sphere covering each of these triangular voids generates a tetrahedral site. Since each sphere contributes to the formation of two potential tetrahedral voids (one above, one below), the total number of tetrahedral voids is twice the number of spheres (2N).
7. Where are the octahedral and tetrahedral voids located within a face-centred cubic (FCC) unit cell?
In a face-centred cubic (FCC) or cubic close-packed (CCP) lattice, which has 4 effective atoms (N=4), the voids are at specific locations:
- Octahedral Voids: There are 4 octahedral voids. One is at the absolute body centre of the cube, and there is one at the centre of each of the 12 edges (each edge-centred void is shared by 4 unit cells, contributing 12 x 1/4 = 3 to the count). Total = 1 + 3 = 4.
- Tetrahedral Voids: There are 8 tetrahedral voids, all located completely inside the unit cell. They are positioned on the body diagonals, with two voids on each of the four body diagonals.
8. What is the practical importance of understanding voids in crystal structures?
Understanding voids is essential for several reasons in chemistry and materials science:
- Ionic Compounds: The structure of many ionic solids (e.g., NaCl) depends on which type of void the smaller cations occupy within the lattice formed by the larger anions.
- Interstitial Compounds: Small non-metal atoms like hydrogen, carbon, or nitrogen can fit into these voids in a metal's crystal lattice to form interstitial alloys like steel, which have significantly altered properties like hardness and strength.
- Density and Stability: The type and extent of void filling directly influence the density, stoichiometry, and overall stability of a crystalline solid.
9. What determines whether a smaller atom or ion will occupy a tetrahedral or an octahedral void?
The primary factor is the radius ratio rule, which relates the size of the smaller particle (cation, r+) to the size of the larger particle forming the lattice (anion, r-). For maximum stability, the smaller particle should fit snugly into the void. A specific range of the radius ratio (r+/r-) is ideal for a particle to occupy a tetrahedral void, while a different range is suitable for the larger octahedral void. If the radius ratio falls outside the stable range for a particular void, the crystal structure will be less stable or adopt a different arrangement.
