Master AP Chemistry Topic 3.2 with this free practice quiz. Test your skills on covalent network solids, molecular solids, ionic solids, and metallic solids to prepare for your AP exam.
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Introduction to the Properties of Solids
Solids have a definite shape and volume because their particles are held in relatively fixed positions by attractive forces. However, the properties of solids vary considerably depending on the types of particles present and the forces or bonds holding them together.
The AP Chemistry curriculum distinguishes four major types of solids: ionic solids, covalent network solids, molecular solids, and metallic solids. Their structures explain differences in hardness, electrical conductivity, melting point, and other physical properties.
Solids can also be classified according to their structure as crystalline or amorphous. Alloys, which contain mixtures of elements with metallic properties, demonstrate how changes in particle arrangement can modify the properties of a solid.
Understanding the relationship between microscopic structure and macroscopic properties is central to predicting how a solid behaves.
1. Ionic Solids
Ionic solids consist of positively charged cations and negatively charged anions arranged in an organized three-dimensional lattice. The particles are held together by strong electrostatic attractions between opposite charges.
Sodium chloride, NaCl, is a typical ionic solid. Its crystal structure contains alternating Na⁺ and Cl⁻ ions. The formula NaCl represents the simplest whole-number ratio of ions in the lattice, not a collection of separate NaCl molecules.
Physical properties of ionic solids
Ionic solids generally have high melting points because substantial energy is required to overcome the electrostatic attractions holding the ions together. The strength of these attractions depends on the magnitude of the ionic charges and the distance between ions.
Ionic solids are usually hard but brittle. When a force shifts the layers of ions, ions with the same charge can become adjacent. Their strong electrostatic repulsion causes the crystal to fracture.
Solid ionic compounds do not conduct electricity effectively because their ions are fixed in the lattice and cannot move freely. However, molten ionic compounds and aqueous solutions containing mobile ions can conduct electricity.
For example, solid sodium chloride does not conduct electricity effectively, whereas molten NaCl conducts because its ions are mobile.
The dissolution of sodium chloride in water can be represented by:
NaCl(s) → Na⁺(aq) + Cl⁻(aq)
The resulting solution conducts electricity because the dissolved ions can move and carry charge.
2. Covalent Network Solids
Covalent network solids consist of atoms connected by an extensive network of covalent bonds. Unlike molecular solids, they do not contain separate molecules held together primarily by intermolecular forces.
Diamond is a classic example. Each carbon atom forms four covalent bonds with neighboring carbon atoms, creating a rigid three-dimensional network.
Silicon dioxide, SiO₂, also forms a covalent network solid. Its atoms are connected through an extended arrangement of Si–O bonds rather than existing as separate SiO₂ molecules in the solid structure.
Physical properties of covalent network solids
Covalent network solids generally have very high melting points because melting or breaking apart the structure requires disrupting many strong covalent bonds.
They are often hard and rigid. Diamond, for example, is extremely hard because its carbon atoms are connected by a strong three-dimensional network.
Electrical conductivity varies depending on the structure and electronic properties of the solid. Diamond does not conduct electricity effectively because its electrons are localized in covalent bonds. Graphite, however, conducts electricity along its layers because it contains delocalized electrons that can move through the structure.
Graphite consists of carbon atoms bonded in sheets of hexagonal rings. Each carbon forms three covalent bonds, leaving electrons that are delocalized within the layers. The layers are held together by relatively weak attractions, allowing them to slide over one another. This explains why graphite is soft and useful in pencils, despite the strong covalent bonds within each layer.
The contrast between diamond and graphite demonstrates that the properties of a solid depend not only on the elements present but also on how their atoms are arranged and bonded.
3. Molecular Solids
Molecular solids consist of discrete molecules held together by intermolecular forces. These forces may include London dispersion forces, dipole-dipole interactions, and hydrogen bonding.
Examples include ice, solid carbon dioxide (dry ice), iodine (I₂), and solid sugar.
Within each molecule, atoms are connected by covalent bonds. When a molecular solid melts, the molecules generally remain intact; the process primarily overcomes attractions between neighboring molecules rather than breaking the covalent bonds within them.
Physical properties of molecular solids
Molecular solids generally have lower melting points than ionic and covalent network solids because intermolecular forces are usually weaker than the strong ionic attractions or extended covalent bonds found in those solids.
However, melting points vary considerably. Molecular solids containing hydrogen bonds can have stronger attractions than similar-sized molecules with only dispersion forces. Larger molecules can also have substantial dispersion forces.
Molecular solids are generally poor electrical conductors because they do not contain freely moving charged particles or a mobile population of electrons.
For example, solid iodine consists of nonpolar I₂ molecules. The molecules are attracted mainly by London dispersion forces. When iodine sublimes, it changes directly from solid to gas without requiring the I–I covalent bonds within each molecule to break.
Dry ice behaves similarly:
CO₂(s) → CO₂(g)
This phase change is sublimation. The CO₂ molecules remain intact as they separate from the solid.
4. Metallic Solids
Metallic solids consist of metal atoms arranged in a closely packed structure, with valence electrons delocalized throughout the solid. These electrons are not restricted to individual bonds between specific pairs of atoms.
Metallic bonding is the electrostatic attraction between the positively charged metal-ion cores and the delocalized electrons.
Physical properties of metallic solids
Metals conduct electricity because their delocalized electrons can move through the solid in response to an electric field. They also conduct thermal energy effectively because mobile electrons and particle vibrations transfer energy through the structure.
Metals are generally malleable, meaning they can be hammered into sheets, and ductile, meaning they can be drawn into wires. When layers of metal atoms shift, the delocalized electrons continue to attract the metal-ion cores. The structure can therefore deform without the abrupt fracture typical of ionic crystals.
Metallic solids also tend to have a shiny appearance because their mobile electrons interact with incoming light.
Their hardness and melting points vary widely. These properties depend on factors such as the number of delocalized electrons, atomic size, and the structure of the solid.
For example, copper is an effective electrical conductor and is used in electrical wiring because its mobile electrons carry charge efficiently.
5. Interstitial and Substitutional Alloys
An alloy is a material containing a metal combined with one or more other elements. Alloys are often designed to improve properties such as strength, hardness, corrosion resistance, or durability.
The arrangement of atoms in an alloy affects how easily the layers of atoms can slide past one another.
Substitutional alloys
In a substitutional alloy, some atoms of the original metal are replaced by atoms of another element. This is most feasible when the atoms have reasonably similar sizes and compatible crystal structures.
Brass is a substitutional alloy composed mainly of copper and zinc. Some zinc atoms occupy positions normally occupied by copper atoms in the metallic structure.
Because the substituted atoms may differ in size from the original atoms, they distort the regular arrangement of the lattice. This can make the movement of atomic layers more difficult, increasing the hardness or strength of the material.
Interstitial alloys
In an interstitial alloy, smaller atoms occupy spaces between the larger metal atoms in the crystal structure.
Steel is an important example. It contains iron and carbon, with some carbon atoms occupying interstitial positions in the iron structure.
The small carbon atoms distort the metallic lattice and hinder the movement of layers of atoms. This can increase the strength and hardness of the alloy compared with pure iron.
The properties of an alloy depend on its composition, atomic arrangement, and processing. Increasing the concentration of an added element does not necessarily improve every property, because strength, ductility, and other characteristics may change in different ways.
6. Crystalline and Amorphous Solids
Solids can also be classified by the degree of order in their particle arrangements.
Crystalline solids
Crystalline solids have a regular, repeating arrangement of particles extending over long distances. This repeating structure is called a crystal lattice.
Examples include sodium chloride, diamond, quartz, and many metals. Their organized structures can produce characteristic geometric shapes and distinct melting points.
A crystalline solid generally melts over a relatively narrow temperature range at a given pressure because its particles occupy an ordered structure that changes into a liquid at a characteristic temperature.
The physical properties of crystalline materials can depend on direction. For example, electrical conductivity, strength, or the ease of breaking a crystal may vary with the direction in which the property is measured. This behavior is called anisotropy.
Amorphous solids
Amorphous solids lack long-range repeating order. Their particles do not form the same regular, extended arrangement found in crystalline solids, although local organization may still exist.
Glass is a common example of an amorphous solid. Many plastics and some forms of solid silica are also amorphous.
Amorphous solids generally soften over a range of temperatures rather than having a single, sharply defined melting point. As the temperature rises, the material gradually becomes softer and more able to flow.
Crystalline and amorphous solids can have the same or similar chemical compositions but different physical properties because their particles are arranged differently.
For example, crystalline quartz and amorphous silica glass both consist primarily of SiO₂, but their atomic arrangements differ. Quartz has a repeating crystal structure, whereas glass lacks long-range crystalline order.
7. Connecting Solid Structure to Observable Properties
To predict the properties of a solid, identify the particles present, the attractions or bonds holding them together, and whether charge carriers can move through the structure.
Ionic solids have strong electrostatic attractions between ions, generally high melting points, and electrical conductivity when molten or dissolved if mobile ions are present.
Covalent network solids contain extended covalent bonds, generally have very high melting points, and vary in conductivity according to their electronic structure.
Molecular solids contain discrete molecules held together by intermolecular forces. They generally have lower melting points and poor electrical conductivity.
Metallic solids contain delocalized electrons, explaining their electrical conductivity, thermal conductivity, malleability, and ductility.
Alloys modify the regular arrangement of metal atoms through substitutional or interstitial incorporation, often improving strength and hardness. Crystalline and amorphous structures further demonstrate that particle arrangement influences physical behavior.
In AP Chemistry, these relationships are used to interpret experimental observations and compare unknown solids. A high melting point alone does not uniquely identify a solid type, and conductivity must be considered alongside the state of the substance and the mobility of its charged particles. The most reliable predictions come from combining evidence about bonding, structure, and particle mobility.