Name: molecular and ionic crystal properties
Uploaded: 2020-09-24T15:06:49.000Z
Duration: 2 min 54 s
Channel: JoVE

K12 Standards

Virginia  Standards of Learning

HIgh School

Physics

PH.9

PH.9.e

Molecular and Ionic 
Solids

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&#8211;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&#8211;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&#8217;s properties. Although they are both nonpolar solids, iodine&#8217;s melting point i

molecular and ionic 
solids

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. 
 
&#8226; 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 &#8722;200 &#176;C). Substances consisting of larger, nonpolar molecules have larger attractive forces and melt at higher temperatures. 
&#8226; 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&#8217;s law: 
<img alt='&#60;p lang="en"&#62;Static equilibrium; &#931;Fx=0, MA=0; mechanical diagram with forces and moments illustrated.&#60;/p&#62;' src="/files/ftp_upload/11354/11354_Equation_1.png" style="text-align: center; height: 75px;" data-altupdatedat="1754904545000" data-alt="Static equilibrium, &#931;Fx=0 diagram of forces acting on an object, showcasing equilibrium principles." />
Here, K is a constant of proportionality, r is the distance between the charges, and qa and qc 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&#8211;cation and anion&#8211;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.
<table style="margin-left: auto; margin-right: auto;">
	<tbody>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Type of crystalline solid</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Type of constituent particle</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Type of attractions</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Properties</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Examples</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Molecular Solids</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Molecules</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Intermolecular forces (IMFs): Dispersion forces, dipole-dipole forces, hydrogen bonds</td>
			<td style="color: black; border: 1px solid black; text-align: center;">variable hardness, variable brittleness, low melting points, a poor conductor of heat and electricity</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Ar, H2O (ice), CO2 (dry ice), I2, C12H22O11 (sucrose)
			
			
			</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Ionic Solids</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Ions</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Electrostatic</td>
			<td style="color: black; border: 1px solid black; text-align: center;">hard, brittle, high to very high melting points, conductor of electricity in molten and dissolved state</td>
			<td style="color: black; border: 1px solid black; text-align: center;">NaCl (table salt), 
			MgO(Magnesium Oxide), 
			Al2O3 (alumina)</td>
		</tr>
	</tbody>
</table>
Part of this text has been adapted from <a href="https://openstax.org/books/chemistry-2e/pages/10-5-the-solid-state-of-matter">Openstax, Chemistry 2e, Section 10.5: The Solid State of Matter.</a>

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

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Chemistry

Liquids, Solids, and Intermolecular Forces

KEY TERMS AND CONCEPTS

molecular and ionic crystal properties

Crystalline solids are grouped into four types: molecular, ionic, metallic, and covalent network solids. This page focuses on molecular and ionic crystals and how their particle type and attractions shape their properties.

Molecular crystalline solids are made of neutral molecules. Common examples include ice, sucrose, and iodine. The molecules are held together by weak intermolecular forces, including London dispersion forces, dipole-dipole interactions, and hydrogen bonds. These forces control the melting points and other properties of the crystal.

Small, symmetrical, nonpolar molecules such as H2, N2, O2, and F2 have weak dispersion forces. They form molecular solids with very low melting points, often below -200 °C. Larger nonpolar molecules have stronger attractive forces and melt at higher temperatures. Molecular solids made of polar molecules melt at still higher temperatures. Solid SO2 and table sugar are examples. Hydrogen bonding is the main force that holds together the crystal lattice of ice.

How well the molecules pack also matters. Efficient packing in three dimensions gives stronger overall attraction and a higher melting point. Higher symmetry helps molecules pack closely in the crystal. Lower symmetry makes packing less efficient, so the intermolecular forces are less effective and the melting point is lower.

Ionic crystalline solids, such as sodium chloride, are made of positive and negative ions. These ions are held together by strong electrostatic attractions. Ionic solids usually have high melting points because these attractions are strong. The strength of the ionic force can be estimated with Coulomb's law, where the force depends on the charges and the distance between them. Higher ion charges give stronger attraction, and closer packing in the crystal lattice also increases the force.

Ionic solids are hard and tend to be brittle. They often shatter instead of bending because the lattice contains both attractive ion-pair interactions and repulsive like-charge interactions. They do not conduct electricity when solid because the ions cannot move freely. When melted or dissolved in water, the ions can move and the material can conduct electricity. Examples include NaCl, MgO, and Al2O3.

Molecular and Ionic Crystal Properties

molecular comparison of gases, liquids, and solids

molecular-comparison-of-gases-liquids-and-solids

How Intermolecular Forces Shape States

Intermolecular forces and particle motion shape the states of matter: solids, liquids, and gases. In a solid, particles are tightly packed and often arranged in a regular pattern. A solid has a fixed shape, and its particles vibrate in place rather than flowing past one another. In a liquid, particles stay close together but do not have a regular arrangement. Liquids have no fixed shape, yet their particles move past each other while remaining in contact. That is why liquids can flow and take...

intermolecular vs intramolecular forces

How Forces Shape Liquids and Solids

Intermolecular forces are electrostatic attractions between molecules. Their strength depends on how far apart the molecules are. These forces matter most in solids and liquids, where particles are close together. In gases, they become important mainly under high pressure because the molecules are then closer together. Intermolecular forces help control key physical properties. They affect melting point, boiling point, density, and the enthalpies of fusion and vaporization. When a liquid is...

intermolecular forces

How Molecules Stick Together

Molecules and ions stick together through intermolecular forces. These are electrostatic attractions and repulsions caused by permanent, partial, or temporary charges. They act between ions, polar molecules, nonpolar molecules, and neutral atoms, and they vary in strength. Some intermolecular forces act between neutral particles and are called van der Waals forces. These include dipole-dipole forces, hydrogen bonding, and dispersion forces. Another important interaction is ion-dipole force,...

comparing intermolecular forces: melting point, boiling point, and miscibility

comparing-intermolecular-forces-melting-point-boiling-point

How Molecule Shape Affects Intermolecular Forces

Intermolecular forces are the attractions between molecules, and they help explain melting point, boiling point, and miscibility. Their strength depends on molar mass, molecular shape, and polarity. These forces shape the bulk behavior of substances that students see in chemistry labs. Dispersion forces are temporary attractions that occur in every substance, whether it is polar or nonpolar. They form when electrons are spread unevenly for a moment, creating a temporary dipole. Stronger...

surface tension, capillary action, and viscosity

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How Liquids Behave on Surfaces

Liquids behave in different ways on surfaces because of surface tension, adhesive forces, capillary action, and viscosity. These properties help explain why some liquids bead up, some spread out, and some flow more easily than others. Surface tension comes from cohesive forces, which are the attractions between identical molecules. Molecules inside a liquid are pulled equally in all directions, but molecules at the surface are pulled by fewer neighbors. That imbalance makes the liquid contract...

phase transitions

Melting, Boiling, and State Changes

Melting, boiling, and other state changes depend on how particles are arranged in a substance. In solids, particles pack closely in an ordered pattern. They vibrate in fixed positions but do not move past nearby particles. In liquids, particles are still close together, but they are randomly arranged and can move past one another. Solids and liquids are called condensed phases because their particles remain close and their intermolecular forces are relatively strong. Gases behave differently...

phase transitions: vaporization and condensation

phase-transitions-vaporization-and-condensation

Heat and Mass Transfer in Phase Changes

Heat and mass transfer drive phase changes in liquids. When a liquid is heated, its molecules move faster. Some surface molecules can then overcome intermolecular forces and escape into the gas phase. This process is called vaporization, and it happens at the surface of the liquid. Vaporization can also happen in a closed container. In that case, gas molecules cannot leave the container. The gas molecules move randomly and may collide with the surface of the condensed phase. Some of those...

vapor pressure

Vapor Pressure and Boiling Point

Vapor pressure and boiling point describe how a liquid behaves in a closed container and when it starts to boil. As a liquid vaporizes, some gas molecules stay above the liquid and can collide with the surface again. When those collisions turn gas back into liquid, the process is called condensation. If vaporization and condensation happen at the same rate, the liquid and vapor amounts stay constant. The vapor is then in equilibrium with the liquid. This is a dynamic equilibrium, which means...

clausius-clapeyron equation

Predicting Vapor Pressure with Temperature

Predicting vapor pressure with temperature is a key use of the Clausius-Clapeyron equation. It describes how the equilibrium between a liquid and its vapor changes as temperature changes. When temperature rises, the vapor pressure of the liquid also rises. The equation gives a quantitative link between vapor pressure, P, and temperature, T. It shows the rate at which vapor pressure increases for each unit increase in temperature. In this relationship, ΔHvap is the enthalpy of vaporization, R...

phase transitions: melting and freezing

Melting Point and Freezing Equilibrium

Melting point and freezing equilibrium describe how a crystalline solid changes between solid and liquid states. As a solid is heated, the average energy of its atoms, molecules, or ions increases, and the solid becomes hotter. At a certain point, the added energy becomes large enough to partly overcome the forces that hold the particles in fixed positions. The solid then starts melting, or changing into a liquid. During melting, the temperature stops rising even though heat is still being...

phase transitions: sublimation and deposition

phase-transitions-sublimation-and-deposition

Sublimation, Frost, and Enthalpy Changes

Sublimation is a phase change in which a solid turns directly into a gas. It skips the liquid state. Dry ice is a clear example at room temperature and standard pressure, where solid carbon dioxide seems to disappear as it sublimes without forming a liquid first. Some other solids also sublime under the right conditions. Snow and ice can sublimate below water’s melting point, and wind or the lower air pressure at high altitudes can speed up that slow process. Solid iodine does the same when...

heating and cooling curves

Phase Changes on Heating and Cooling Curves

Phase changes on heating and cooling curves show how a substance responds when it absorbs or releases heat while isolated from its surroundings. The graph tracks changes in temperature and state as the substance is heated or cooled. When a solid is heated, its temperature rises first. The amount of heat absorbed depends on the solid’s heat capacity, which is written as q = mcΔT. In this equation, m is the mass and c is the specific heat. This relationship applies when matter changes...

phase diagrams

Water and CO2 Phase Behavior

Water and carbon dioxide phase behavior is shown with phase diagrams that map pressure against temperature. These plots show where a substance exists as a solid, liquid, or gas. They also show how melting points, sublimation points, and boiling points change with pressure. In a phase diagram, the labeled regions for solid, liquid, and gas represent single phases. The lines or curves between them show two phases in equilibrium. The triple point is the one condition of pressure and temperature...

structures of solids

Crystalline and Amorphous Solid Patterns

Crystalline and amorphous solids differ in how their particles are arranged. In crystalline solids, atoms, ions, or molecules follow a definite repeating pattern. Metals and ionic compounds typically form these ordered solids. Because the same particles are held in place by the same forces or energy, crystalline solids have a precise melting temperature. Amorphous solids, also called non-crystalline solids or glasses, do not have an ordered internal structure. Their particles are arranged...

Crystalline solids are grouped into four types: molecular, ionic, metallic, and covalent network solids. This page focuses on molecular and ionic crystals and how their particle type and attractions shape their properties. Molecular crystalline solids are made of neutral molecules. Common examples include ice, sucrose, and iodine. The molecules are held together by weak intermolecular forces, including London dispersion forces, dipole-dipole interactions, and hydrogen bonds. These forces...

ionic crystal structures

Ionic Crystal Packing and Ion Holes

Ionic crystal packing depends on ion size, charge, and the ratio of cations to anions. In these solids, the ions usually behave like charged spheres. Stable structures form when each ion is surrounded by as many oppositely charged ions as possible and the ions remain in contact. In simple ionic crystals, the larger anions usually form a close-packed array. The smaller cations fit into the spaces between them. Two common spaces are tetrahedral holes, which are surrounded by four anions, and...

metallic solids

Metallic Bonding and Crystal Packing

Metallic solids are made of metal atoms arranged in crystals. Copper, aluminum, and iron are common examples. These solids are often described as metal nuclei spread through a sea of delocalized electrons. The atoms are held together by metallic bonding, and that bonding gives metals many of their useful bulk properties. Metallic solids are shiny, conduct heat and electricity well, and can be shaped without breaking. This shaping ability is called malleability, which means a metal can deform...

band theory

Band Gaps in Conductors and Insulators

Band gaps in solids help explain why some materials conduct electricity and others do not. When many atoms come together, their atomic orbitals combine into molecular orbitals with distinct energies. In a solid, so many orbitals are involved that these energies merge into continuous regions called bands. The energy difference between bands is called the band gap. Electrons in a solid must move across orbitals of different energies to travel through the material. In conductors, the valence...

network covalent solids

Why Network Covalent Solids Are So Hard

Network covalent solids are made of atoms joined in a three-dimensional network of covalent bonds. They appear in the crystal structures of nonmetals such as diamond, graphite, and silicon. Some covalent compounds also form this kind of structure, including silicon dioxide, which is found in sand, and silicon carbide, the abrasive used on sandpaper. Many minerals also contain networks of covalent bonds. These solids are hard to melt or break because their covalent bonds must be broken first.

x-ray crystallography

X-ray Diffraction in Crystal Structure

X-ray diffraction can reveal the unit cell size and atom arrangement in a crystal. This method is called X-ray crystallography. It uses diffraction, which is the bending of an electromagnetic wave when it meets a barrier about the same size as its wavelength. X-rays work well for this method because their wavelengths are about the same as the spacing between neighboring atoms in crystals. They are measured on the scale of a few angstroms. When monochromatic X-rays, or X-rays with one...

lattice centering and coordination number

lattice-centering-and-coordination-number

Crystal Packing in Cubic Unit Cells

Crystal packing in cubic unit cells shows how atoms are arranged in a repeating solid structure. A crystal is best described by its unit cell, which is the smallest repeating part of the lattice. Lattice points mark where atoms or ions are located. The full crystal is built when this unit cell repeats in three dimensions. The cubic lattice has three common unit cells. A simple cubic structure, also called a primitive cubic structure, can be pictured as identical spheres stacked in layers with...

key terms and concepts

Explore how molecular and ionic crystal solids differ in bonding, packing, melting point, hardness, brittleness, and electrical conductivity.

Molecular and Ionic Crystal Properties | JoVE K12

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Molecular and Ionic Crystal Properties

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