Structure from properties
Table salt, table sugar, quartz sand, and a copper wire are all solids, yet they behave very differently: salt shatters when struck and conducts when melted; sugar melts soon after heating; quartz doesn’t melt until well over a thousand degrees; copper wire bends and conducts. AP topic 3.2 asks you to sort solids into four types and explain these properties by “which particles sit at the lattice points, and what holds them together,” and, working backward, to identify the type of a solid from its properties.
Four types of crystalline solids
| Type | Particles at lattice points | Forces between particles | Melting point | Conductivity | Mechanical properties | Examples |
|---|---|---|---|---|---|---|
| Ionic solid离子晶体 | cations and anions | ionic bonding | high | none as a solid; conducts when molten or dissolved | hard and brittle | NaCl, MgO |
| Molecular solid分子晶体 | molecules | intermolecular forces | low | none | soft | , ice, dry ice, sucrose |
| Covalent network solid共价网状晶体 | atoms | covalent bonds throughout the crystal | very high | usually none (graphite is the exception) | very hard (graphite is the exception) | diamond, , Si |
| Metallic solid金属晶体 | metal cations (in a sea of electrons) | metallic bonding | varies widely | conducts as a solid and a liquid | malleable and ductile | Cu, Fe, Na |
The key question: what has to be overcome to melt (or boil) the solid? In an ionic solid it is the Coulombic attraction between ions, and in a covalent network solid it is covalent bonds, so both melt at high temperatures. In a molecular solid only the intermolecular forces have to be overcome and the molecules stay intact, so the melting point is low.
Ionic and metallic solids (review)
- Ionic solids (U2.3): the Coulombic attraction between ions is strong, so the melting point is high and the vapor pressure very low. When a layer of ions shifts, like charges face each other and repel, so the solid is brittle. The ions can move, and the substance conducts, only when it is molten or dissolved in water.
- Metallic solids (U2.4): delocalized electrons move freely, so the solid conducts; metallic bonding is nondirectional, so layers can slide and the metal is malleable. Melting points vary widely: cesium melts at 28.4 °C, in the palm of your hand, while lithium melts at 180.5 °C (CRC).
Molecular solids
The lattice points of a molecular solid hold individual molecules. Inside each molecule are strong covalent bonds, but between molecules there are only intermolecular forces (U3.1), so molecular solids melt at low temperatures and are soft. They don’t conduct either: their electrons are localized in the covalent bonds and lone pairs of each molecule, so there are no mobile charged particles. Melting moves the molecules a little farther apart and leaves the molecules themselves intact. The stronger the intermolecular forces, the higher the melting point: melts at −101.5 °C, and , with far more electrons, at 113.7 °C (CRC).
Molecular solids often have fairly high vapor pressures: iodine slowly sublimes at room temperature, and dry ice (solid ) turns directly into gas.
Covalent network solids
In a covalent network solid, atoms are joined by covalent bonds into a network that runs through the whole crystal, with no separate molecules. Melting requires breaking many covalent bonds, so these solids have very high melting points and are very hard; with no mobile charged particles, they usually don’t conduct.
- Diamond: each C atom forms 4 covalent bonds to 4 neighboring C atoms (sp³, tetrahedral). It is the hardest natural substance. Cubic boron nitride (BN) has the same structure and is also extremely hard: covalent bonds with fixed directions and angles lock the atoms in place in all three dimensions.
- Silicon dioxide (quartz): each Si atom is bonded to 4 O atoms, and each O atom to 2 Si atoms. and have similar formulas but entirely different structures: forms a molecular solid, and dry ice sublimes at room temperature, while quartz doesn’t melt until above 1500 °C.
- Graphite: each C atom forms 3 covalent bonds to 3 C atoms in the same layer (sp²), making flat sheets of hexagons. The remaining p electron on each C becomes a delocalized π electron that can move along the layer, so graphite conducts. Only dispersion forces hold the layers together, so they slide apart easily, making graphite soft and slippery, useful as a lubricant and in pencils.
Graphite and diamond are both made of carbon atoms; they are allotropes. Their very different properties come entirely from how the atoms are connected.
Look at the particle diagrams and properties of five solids, then switch to “Identify an unknown solid” and classify samples from experimental observations:
Cations and anions alternate, with oppositely charged ions next to each other.
| Particles at lattice points | Cations and anions |
|---|---|
| Forces between particles | Ionic bonding (Coulombic attraction) |
| Melting point | High |
| Electrical conductivity | None as a solid; conducts when molten or dissolved |
| Hardness and malleability | Hard and brittle |
| Examples | NaCl, MgO, CaF₂ |
Individual molecules (diatomic ones here) in an orderly arrangement: covalent bonds inside each molecule, only intermolecular forces between molecules.
| Particles at lattice points | Molecules |
|---|---|
| Forces between particles | Intermolecular forces (covalent bonds within molecules) |
| Melting point | Low |
| Electrical conductivity | None as a solid or liquid |
| Hardness and malleability | Soft |
| Examples | I₂, ice, dry ice (CO₂), sucrose |
Si (yellow) and O (red) are linked by covalent bonds into one huge network, so the whole crystal is like a single giant molecule (a flat sketch of a 3D structure).
| Particles at lattice points | Atoms |
|---|---|
| Forces between particles | Covalent bonds throughout the crystal |
| Melting point | Very high |
| Electrical conductivity | Usually none |
| Hardness and malleability | Very hard |
| Examples | SiO₂ (quartz), diamond, Si, SiC |
Side view: within each layer the C atoms are covalently bonded, and the π electrons (blue dots) can move along the layer; between layers there are only dispersion forces.
| Particles at lattice points | C atoms |
|---|---|
| Forces between particles | Covalent bonds within layers, dispersion forces between layers |
| Melting point | Very high |
| Electrical conductivity | Conducts along the layers (delocalized π electrons) |
| Hardness and malleability | Soft and slippery; layers slide apart easily |
| Examples | Graphite |
Metal cations form a lattice, and the valence electrons (blue dots) move freely through the whole piece of metal.
| Particles at lattice points | Metal cations (in a sea of electrons) |
|---|---|
| Forces between particles | Metallic bonding |
| Melting point | Usually fairly high, varies widely |
| Electrical conductivity | Conducts as a solid and a liquid |
| Hardness and malleability | Malleable and ductile |
| Examples | Cu, Fe, Na, alloys |
Each sample comes with experimental observations. Decide what type of solid it is:
Example 1: SiO₂ and CO₂
C and Si are both in group 14, yet is a gas at room temperature, while (quartz) melts above 1500 °C. Explain in terms of particles.
-
Identify the type of solid. is made of separate linear O=C=O molecules, so solid (dry ice) is a molecular solid. In , each Si is bonded to 4 O and each O to 2 Si, linking the atoms covalently in a three-dimensional network: a covalent network solid with no separate “ molecules.”
-
What must be overcome on melting (or subliming)?
- (molecular solid): only the dispersion forces between molecules, which are weak.
- (covalent network solid): many Si–O covalent bonds must break, which are very strong.
-
Conclusion. Covalent bonds are far stronger than intermolecular forces, and breaking Si–O bonds throughout the network takes a great deal of energy, so melts far higher; needs to overcome only weak dispersion forces and is a gas at room temperature.
-
A common mistake. Don’t argue that “Si–O bonds are stronger than C=O bonds”: no C=O bonds break when becomes a gas, so comparing bond energies is beside the point.
Example 2: Graphite and diamond
Graphite conducts electricity and is soft, while diamond doesn’t conduct and is extremely hard. Explain using their structures.
-
Compare the structures first.
- Diamond: each C forms 4 covalent bonds in a three-dimensional network; there are no mobile electrons.
- Graphite: each C forms 3 covalent bonds in flat layers; each layer has delocalized p electrons; only dispersion forces act between layers.
-
Conductivity. Conduction needs mobile charged particles. In diamond all the valence electrons are in localized σ bonds and can’t move, so it doesn’t conduct; in graphite each C has one leftover p electron delocalized over the whole layer, which can move along the layer, so graphite conducts.
-
Hardness. Diamond is held by covalent bonds in all three directions, so deforming it means breaking covalent bonds, which makes it extremely hard; graphite’s layers are held together only by weaker dispersion forces and slide apart easily, so it is soft and slippery.
-
What they share. Both have very high melting points, because melting either one breaks C–C covalent bonds. Keep the properties apart: melting point depends on what breaks on melting, hardness on what must be disrupted to deform the solid, and conductivity on whether there are mobile charged particles.
- KBr
A reactive metal and a nonmetal: and .
- Dry ice (solid )
Individual molecules.
- Diamond
C atoms covalently bonded into a 3D network.
- Iron
Metal cations and a sea of electrons.
- Quartz,
Si and O covalently bonded into a network, with no separate molecules.
- Sucrose
Made of molecules; low melting point, and its solution doesn't conduct.
- MgO
and .
- Ice
Water molecules held together by hydrogen bonds and other intermolecular forces.
- Silicon carbide, SiC
Si and C each form four covalent bonds in a network; very hard.
Only dispersion forces between molecules.
and .
- Graphite
It conducts, but it is a covalent network solid, not a metal.
- Brass
A copper–zinc alloy, with a sea of electrons.
Common mistakes
Mistake 1: “Melting a molecular solid breaks covalent bonds.” Melting and boiling only overcome the forces between molecules; the covalent bonds inside each molecule stay intact. Melted ice is still molecules.
Mistake 2: “A solid that conducts must be a metal.” Graphite is a covalent network solid; its delocalized π electrons let it conduct.
Mistake 3: “Covalent bonds are strong, so every substance with covalent bonds has a high melting point.” Only covalent network solids break covalent bonds when they melt. Molecular solids contain covalent bonds inside their molecules, but melting overcomes only intermolecular forces, so their melting points are low.
Quick check
Dry ice sublimes at room temperature, but quartz () melts above 1500 °C. Which is the best explanation?
The key is what must be overcome. B: subliming doesn't break any C=O bonds, so comparing bond energies is pointless. C: mass isn't the reason. D is a wrong fact.
Which of these solids conducts electricity in the solid state?
The π electrons in each graphite layer can move. Diamond's electrons are all in localized σ bonds; the ions in solid NaCl can't move; is a molecular solid.
A solid has a high melting point, is hard and brittle, and does not conduct as a solid but conducts when melted. It is most likely
Conducting when melted shows mobile ions, so it is ionic. Cu conducts as a solid; doesn't conduct even when molten; sucrose has a low melting point.
Molten MgO conducts electricity. Which is the best explanation?
The charge carriers are mobile ions. A: MgO has no delocalized electrons; B: there is no water in molten MgO; C: MgO is ionic and has no MgO molecules.
Which change requires breaking covalent bonds?
Diamond is a covalent network solid, so turning it into gaseous C atoms breaks C–C covalent bonds. Subliming iodine or dry ice only overcomes intermolecular forces, leaving the I–I and C=O bonds intact; melting NaCl overcomes ionic bonding (Coulombic attraction), not covalent bonds.
Which lists , NaCl, and (quartz) in order of increasing melting point?
is a molecular solid (113.7 °C); NaCl is ionic (801 °C); is a covalent network solid (above 1500 °C).
Graphite can be used as a lubricant. Which is the best explanation?
The bonds within a layer are strong covalent bonds (A is wrong); only dispersion forces act between layers. B and C are wrong facts: the mobile particles in graphite are π electrons, and graphite is not a metal.
At room temperature, which solid has the highest vapor pressure (sublimes most easily)?
is a molecular solid held only by dispersion forces, so its molecules leave the solid most easily. In the other three, the particles are held by ionic, covalent, or metallic bonding, and their vapor pressures are tiny.
AP-style free response
Four solids: NaCl, , (quartz), and Cu.
(a) Identify the type of each solid.
(b) Which solids conduct electricity as solids? Which conduct only when molten? Explain.
(c) In terms of particles and forces, explain why melts at a far higher temperature than .
(d) Explain why Cu can be hammered into thin sheets while NaCl shatters.
Model answer and scoring
(a) NaCl: ionic; : molecular; : covalent network; Cu: metallic.
(b) Cu conducts as a solid because it has freely moving delocalized electrons. NaCl conducts only when molten (or dissolved), when its ions can move; in the solid they are fixed in the lattice. and don’t conduct in either state, because they have no mobile charged particles.
(c) Melting only overcomes the dispersion forces between molecules; no I–I bonds break. Melting requires breaking many Si–O covalent bonds, which are far stronger than intermolecular forces, so it takes much more energy.
(d) In Cu, metallic bonding is nondirectional, so layers of atoms can slide and remain held together by the sea of electrons. In NaCl, a blow shifts a layer of ions so that like charges face each other; they repel strongly, and the crystal splits along that layer.
Scoring: (a) all four correct; (b) identifies the delocalized electrons in Cu and the mobile ions in molten NaCl; (c) contrasts intermolecular forces with covalent bonds; (d) contrasts the sea of electrons with repulsion between like charges.
Summary
- Four types of crystalline solids: ionic (ions, ionic bonding), molecular (molecules, intermolecular forces), covalent network (atoms, covalent bonds), and metallic (cations in a sea of electrons, metallic bonding).
- What must be overcome on melting sets the melting point: low for molecular solids, high for ionic and covalent network solids.
- Conduction needs mobile charged particles: metals and graphite conduct as solids; ionic solids conduct when molten or dissolved; molecular solids and other covalent network solids don’t conduct.
- Graphite and diamond are both carbon: graphite has delocalized π electrons within layers and only dispersion forces between them, so it conducts and is soft; diamond is a 3D covalent network, so it doesn’t conduct and is extremely hard.
AP tip: Questions often describe the properties of several solids and ask you to identify their types, or give two substances and ask you to compare their melting points or conductivity with a justification. Base the justification on which particles sit at the lattice points, what must be overcome on melting, and whether there are mobile charged particles.