📘 Class 12 – Chemistry
Chapter 01: The Solid State ⚪
📌 PART – 4
Welcome to Part 4 of The Solid State chapter from ExamsPoint Notes, where we focus on voids, types of interstitial sites, and the role of radius ratios in determining ionic structure and stability. These concepts are foundational in solid-state chemistry and materials science.
🔍 Topics Covered in This Video (Solid State – Part 4):
✅ Voids / Interstitial Sites
In crystalline solids, atoms or ions are not always packed perfectly; gaps exist between them. These empty spaces are called voids or interstitial sites. They play a significant role in determining the physical properties of solids, including density, ionic conductivity, and stability.
Voids are crucial in the formation of interstitial compounds, where smaller atoms or ions occupy these gaps. Understanding voids is especially important in studying alloys and ionic crystals.
✅ Types of Voids – Trigonal, Tetrahedral, Octahedral
There are mainly three types of voids formed during the packing of spheres:
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Trigonal Voids: Formed when three spheres are in contact, creating a small triangular space.
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Tetrahedral Voids: Formed when a sphere is placed in the depression of three touching spheres. Four spheres form the corners of a tetrahedron around the void.
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Octahedral Voids: Formed when two layers of spheres are placed such that a void lies between six surrounding spheres.
The number of tetrahedral voids is double the number of atoms, while the number of octahedral voids is equal to the number of atoms in a close-packed structure.
✅ Radius Ratio
The radius ratio is the ratio of the radius of the smaller ion (usually a cation) to that of the larger ion (anion). It helps in predicting the type of void a cation can occupy in a crystal structure:
Radius Ratio = r+ / r-
Depending on the ratio, the geometry of the resulting ionic structure varies. For example:
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0.225 to 0.414: Tetrahedral voids
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0.414 to 0.732: Octahedral voids
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0.732 to 1.0: Cubic voids
✅ Conditions for Ionic Stability
For a stable ionic crystal:
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The radius ratio must fall within certain limits so that the cation fits perfectly in the void.
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There must be strong electrostatic attraction between oppositely charged ions.
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Lattice energy should be high to overcome repulsive forces.
Ionic compounds with unsuitable radius ratios tend to distort or form unstable lattices. Hence, radius ratio is a key factor in determining coordination number and geometry.
✅ Calculating Radius Ratio
To calculate the radius ratio:
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Identify the radii of the cation (r+) and anion (r-).
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Use the formula: Radius Ratio = r+ / r-
Example: For Na+ (95 pm) and Cl- (181 pm):
Radius Ratio = 95 / 181 ≈ 0.524 ➔ Octahedral void
This calculation helps predict the most probable structure type and coordination number in ionic crystals.
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