11.15

Molecular and Ionic Solids

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Molecular solids are a type of crystalline solids which have molecules or atoms as their constituent particles and are held together by non-ionic intermolecular forces like hydrogen bonds, dispersion forces, or dipole–dipole interactions.

The strength of these intermolecular forces dictates the properties of molecular solids. Overall, these solids tend to be soft, have low melting points, and have low thermal and electrical conductivity.

Nonpolar or net nonpolar molecular solids like solid nitrogen or dry ice are primarily held together by weak dispersion forces. Such solids have very low melting points and sublime easily.

Polar molecular solids like ice and solid sulfur dioxide exhibit hydrogen bonds and dipole–dipole interactions. Such solids have comparatively higher melting points and exist as soft solids or volatile liquids at standard temperature and pressure.

Another example of how the strength of the intermolecular forces influences properties of molecular solids is illustrated by solid iodine.

The increased strength of certain intermolecular forces between larger molecules is reflected in iodine’s properties. Although they are both nonpolar solids, iodine’s melting point is substantially higher than that of solid nitrogen.

Ionic solids are crystalline solids with electrically charged species or ions as constituent particles held together by strong electrostatic forces. For example, sodium chloride is an ionic solid composed of sodium cations and chloride anions.

The packing of ionic solids maximizes the interaction between oppositely charged ions and minimizes the interaction between ions of the same charge. This is often visualized as one set of ions on lattice points and the opposing ions occupying some or all of the spaces between them, or the interstitial sites.

Due to strong coulombic interactions, ionic solids have high melting temperatures. The ionic interactions typically become stronger with an increase in charge or a decrease in ion size.

For example, caesium chloride melts at 645 °C and sodium chloride melts at 801 °C, which can be attributed to the smaller sodium cation enabling closer packing. Calcium oxide, which has higher ionic charges, melts at 2572 °C.

Overview

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Crystalline solids are divided into four types: molecular, ionic, metallic, and covalent network based on the type of constituent units and their interparticle interactions.

Molecular Solids

Molecular crystalline solids, such as ice, sucrose (table sugar), and iodine, are solids that are composed of neutral molecules as their constituent units. These molecules are held together by weak intermolecular forces such as London dispersion forces, dipole-dipole interactions, or hydrogen bonds, which dictate their properties (Table 1).

The strengths of the attractive forces between the units present in different crystals vary widely, which is reflected in the melting points of such crystals.

Small symmetrical nonpolar molecules, such as H₂, N₂, O₂, and F₂, have weak dispersion forces and form molecular solids with very low melting points (below −200 °C). Substances consisting of larger, nonpolar molecules have larger attractive forces and melt at higher temperatures.
Molecular solids composed of polar molecules with permanent dipole moments melt at still higher temperatures. Examples include solid SO₂ and table sugar. Intermolecular hydrogen bonding is mainly responsible for maintaining the three-dimensional lattice of such molecular solids, as seen in frozen water or ice.

The properties of molecular solids depend on the efficient packing of their constituent units, the molecules, in three dimensions. As the intermolecular forces are contact-dependent, higher symmetry of constituent molecules ensures close and compact packing within the crystal structure with high intermolecular attractions. This increases the melting point. The lower symmetry of the molecules prevents their efficient packing. The intermolecular forces, thus, are not as effective, and the melting point is lower.

Ionic Solids

Ionic crystalline solids, such as sodium chloride, are composed of positive and negative ions that are held together by strong electrostatic attractions.

Ionic solids have high melting points due to the strong ionic attractions. The strength of ionic interaction between the cations and the anions in an ionic solid can be approximated by the electrostatic force, given by Coulomb’s law:

<p lang="en">Static equilibrium; ΣFx=0, MA=0; mechanical diagram with forces and moments illustrated.</p>

Here, K is a constant of proportionality, r is the distance between the charges, and q_a and q_c represent the charges on the anions and cations, respectively. Higher the charge on the cations and anions, the stronger the force of ionic attraction. Similarly, close packing of anions and cations in the crystal lattice reduces the distance between the charges, resulting in stronger forces of ionic attraction.

Ionic solids are hard, they also tend to be brittle, and they shatter rather than bend. Their brittleness is attributed to the presence of both attractive (cation-anion) and repulsive (cation–cation and anion–anion) interactions in the crystal lattice. As the ions are unable to move freely due to the strong coulombic forces, Ionic solids do not conduct electricity. However, in the molten state or when dissolved in water, the ions become free to move and conduct electricity.

Table 1. Characteristics of Molecular and Ionic Solids.

Type of crystalline solid	Type of constituent particle	Type of attractions	Properties	Examples
Molecular Solids	Molecules	Intermolecular forces (IMFs): Dispersion forces, dipole-dipole forces, hydrogen bonds	variable hardness, variable brittleness, low melting points, a poor conductor of heat and electricity	Ar, H₂O (ice), CO₂ (dry ice), I₂, C₁₂H₂₂O₁₁ (sucrose)
Ionic Solids	Ions	Electrostatic	hard, brittle, high to very high melting points, conductor of electricity in molten and dissolved state	NaCl (table salt), MgO(Magnesium Oxide), Al₂O₃(alumina)

Part of this text has been adapted from Openstax, Chemistry 2e, Section 10.5: The Solid State of Matter.

Transcript

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$$\rightleftharpoonup{xx}$$ $$\longleftharp{xx}$$, $$\longrightharp{xx}$$,

Molecular solids are a type of crystalline solids which have molecules or atoms as their constituent particles and are held together by non-ionic intermolecular forces like hydrogen bonds, dispersion forces, or dipole–dipole interactions.

The strength of these intermolecular forces dictates the properties of molecular solids. Overall, these solids tend to be soft, have low melting points, and have low thermal and electrical conductivity.

Nonpolar or net nonpolar molecular solids like solid nitrogen or dry ice are primarily held together by weak dispersion forces. Such solids have very low melting points and sublime easily.

Polar molecular solids like ice and solid sulfur dioxide exhibit hydrogen bonds and dipole–dipole interactions. Such solids have comparatively higher melting points and exist as soft solids or volatile liquids at standard temperature and pressure.

Another example of how the strength of the intermolecular forces influences properties of molecular solids is illustrated by solid iodine.

The increased strength of certain intermolecular forces between larger molecules is reflected in iodine’s properties. Although they are both nonpolar solids, iodine’s melting point is substantially higher than that of solid nitrogen.

Ionic solids are crystalline solids with electrically charged species or ions as constituent particles held together by strong electrostatic forces. For example, sodium chloride is an ionic solid composed of sodium cations and chloride anions.

The packing of ionic solids maximizes the interaction between oppositely charged ions and minimizes the interaction between ions of the same charge. This is often visualized as one set of ions on lattice points and the opposing ions occupying some or all of the spaces between them, or the interstitial sites.

Due to strong coulombic interactions, ionic solids have high melting temperatures. The ionic interactions typically become stronger with an increase in charge or a decrease in ion size.

For example, caesium chloride melts at 645 °C and sodium chloride melts at 801 °C, which can be attributed to the smaller sodium cation enabling closer packing. Calcium oxide, which has higher ionic charges, melts at 2572 °C.

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