C. Structure and Properties of Matter · Structure and Properties of Matter

Ionic, Molecular, Covalent Network and Metallic Solids

C2.5 C3.4About 75 minutes

Learning goals

I can describe the particles and forces in ionic, molecular, covalent network and metallic solids.
I can explain the hardness, melting point, conductivity and malleability of each type.
I can predict the type of solid a substance forms from the elements it contains, including products of reactions.
I can explain the different properties of diamond and graphite.

Words to know

ionic solid
A solid made of cations and anions held by ionic bonds in a lattice.

molecular solid
A solid made of molecules held by intermolecular forces.

covalent network solid
A solid in which atoms are joined by covalent bonds in a continuous network, such as diamond or quartz.

metallic bonding
The attraction between metal cations and a sea of delocalized valence electrons.

allotropes
Different structural forms of the same element, such as diamond, graphite and graphene.

alloy
A mixture of a metal with one or more other elements, such as steel or brass.

Four kinds of solid

A diamond ring, a grain of salt, a chunk of ice and a copper wire are all solids, but they could hardly be more different. Diamond is the hardest natural material; salt shatters; ice melts at 0 °C; copper bends and carries electricity. Chemists explain the differences by asking two questions about any solid:

What particles are at the lattice points: ions, molecules, atoms or metal cations?
What forces hold them together?

Remember from Grade 11: in SCH3U you compared ionic and molecular compounds and explained their melting points, brittleness and conductivity. Grade 12 adds two more types, covalent network and metallic solids, and asks you to predict which type a substance forms.

1. Ionic solids

Particles: positive and negative ions. Forces: ionic bonds (electrostatic attraction) in all directions through a lattice.
Properties: high melting points (NaCl 801 °C), hard but brittle, do not conduct as solids but do conduct when molten or dissolved (the ions can then move), many dissolve in water.
Strength depends on charge and size: by Coulomb's law, the attraction is stronger for larger charges and smaller ions (closer together). MgO (Mg²⁺, O²⁻) melts at about 2800 °C, far above NaCl (Na⁺, Cl⁻), which is why MgO is used to line furnaces.

2. Molecular solids

Particles: molecules (or single atoms, for the noble gases). Forces: intermolecular forces only (London, dipole-dipole, hydrogen bonding).
Properties: low melting points (ice 0 °C; iodine 114 °C; sucrose melts and starts to decompose at about 186 °C), soft, non-conductors in every state, solubility depends on polarity.

3. Covalent network solids

Particles: atoms. Forces: a continuous 3-D (or 2-D) network of covalent bonds. The whole crystal is effectively one giant molecule.
Properties: very high melting points (quartz, SiO₂, about 1700 °C; diamond does not melt at normal pressure but turns to vapour above about 3500 °C), very hard, insoluble, usually non-conductors.
Examples: diamond (C), quartz and sand (SiO₂), silicon carbide (SiC, used in sandpaper and cutting tools), and silicon.

Carbon's allotropes show how structure decides properties:

Diamond: every carbon is sp³ and bonded to four others in a tetrahedral network: extremely hard, electrically insulating.
Graphite: each carbon is sp² and bonded to three others in flat hexagonal sheets. The fourth electron of each carbon is in a delocalized π system, so graphite conducts electricity along the sheets. The sheets are held to each other only by London forces, so they slide apart easily: graphite is soft and slippery, which is why it is used in pencils and as a dry lubricant.
Graphene is a single sheet of graphite, one atom thick (see the materials lesson).

4. Metallic solids

Particles: metal cations in a "sea" of delocalized valence electrons that belong to the whole crystal. Forces: metallic bonding, the attraction between the cations and the electron sea.
Properties: conduct electricity and heat as solids and liquids (mobile electrons); lustrous (shiny: the free electrons absorb and re-emit light); malleable and ductile (layers of cations can slide without breaking the bonding, because the electron sea moves with them). Melting points vary enormously: mercury −39 °C, sodium 98 °C, copper 1085 °C, tungsten 3422 °C (more valence electrons and smaller cations usually give stronger metallic bonding).
Alloys such as steel (iron with a little carbon), brass (copper and zinc) and bronze (copper and tin) are usually harder than the pure metal, because atoms of a different size disrupt the neat layers and make sliding harder.

The four types of crystalline solids, classified by the particles in them and the forces holding those particles together.

	Ionic	Molecular	Covalent network	Metallic
Particles	cations and anions	molecules	atoms	cations + electron sea
Forces	ionic bonds	intermolecular forces	covalent bonds	metallic bonds
Melting point	high	low	very high	low to very high
Hardness	hard, brittle	soft	very hard	variable; malleable, ductile
Conducts as solid?	no	no	no (except graphite)	yes
Conducts molten or dissolved?	yes	no	(does not dissolve)	yes (molten)
Examples	NaCl, MgO, CaF₂	H₂O(s), I₂, CO₂(s), sucrose	diamond, SiO₂, SiC	Cu, Fe, Al, Na

Summary of the four types of crystalline solids.

Why metals bend but ionic crystals shatter. In the ionic crystal the bottom row has shifted by one ion, so each + now sits under a + and each − under a −.

Predicting the type of solid

Use the elements in the substance:

Metal only (or an alloy of metals) → metallic.
Metal + non-metal (or a polyatomic ion) → ionic (large ΔEN). Example: magnesium burning in oxygen forms MgO, an ionic solid.
Non-metals only → usually molecular. A few non-metal substances form covalent networks: carbon (diamond, graphite), silicon, SiO₂, SiC, boron and boron nitride. Clues: extremely high melting point, very hard, insoluble.

Then check the prediction against properties (the next lesson is an investigation that does exactly this).

Worked example

Example 1: Classify by properties

A white solid melts at 1710 °C, is extremely hard, does not dissolve in water, and does not conduct electricity even when melted. Classify it.

Very high melting point: ionic, network or metallic, not molecular.
No conductivity when melted rules out ionic (molten ionic compounds conduct) and metallic.
Extremely hard and insoluble fits a covalent network.

A covalent network solid (this is quartz, SiO₂).

Worked example

Example 2: Predict the solid formed in a reaction

Predict the type of solid formed in each reaction: (a) 2 Mg(s) + O₂(g) → 2 MgO(s) (b) Si(s) + O₂(g) → SiO₂(s) (c) 2 H₂(g) + O₂(g) → 2 H₂O, cooled below 0 °C.

(a) Mg is a metal, O a non-metal: ΔEN = 3.4 − 1.3 = 2.1: ionic.
(b) Si and O are both non-metals, but SiO₂ forms a giant network of Si–O bonds (each Si bonded to 4 O, each O to 2 Si): covalent network.
(c) H and O are non-metals that form separate H₂O molecules: ice is molecular (held by hydrogen bonds).

(a) ionic, (b) covalent network, (c) molecular.

Worked example

Example 3: Comparing ionic solids

Rank NaCl, MgO and KCl in order of increasing melting point and explain.

Compare charges: MgO has 2+ and 2− ions; NaCl and KCl have 1+ and 1−. MgO has much the strongest attraction.
Compare sizes for NaCl and KCl: K⁺ is larger than Na⁺, so the ions in KCl are farther apart and attract less strongly.
Measured: KCl 770 °C, NaCl 801 °C, MgO about 2800 °C.

KCl < NaCl < MgO: higher charges and smaller ions give stronger ionic bonding.

Just for fun

Gallium is a metal that melts at about 30 °C, so a piece of it can melt in your hand (body temperature is about 37 °C). Chemists sometimes prank each other with gallium spoons that "melt" when stirred into hot tea. Gallium is not something to handle at home, but the videos are worth a search.

Hands-on activity

A kitchen tour of solids

You need: samples you can find at home: table salt, sugar, a candle (paraffin wax), clean sand or a piece of quartz from a craft store, a pencil, aluminium foil and a copper wire or coin; a magnifying glass; a spoon; a notebook.

Make a table with columns: substance, appearance (crystal shape, lustre), hardness (scratch it with a fingernail; press with the spoon), what happens if you bend it, and your predicted type of solid.
Look at salt and sugar grains under the magnifying glass and sketch their crystal shapes.
Try to crush a salt crystal and a sugar crystal with the back of the spoon on a plate. Try to bend the foil and the wire.
Rub the pencil tip on paper and feel the mark. Scratch the candle with your fingernail. Try to scratch a glass jar with the sand grains, gently, on an old jar you do not need (grown-up note: check the jar for chips first and handle glass carefully).
Classify each substance.

What to look for: salt: cubic crystals, shatters (ionic). Sugar: crushes, soft-ish (molecular). Wax: very soft, scratched by a fingernail (molecular). Sand/quartz: scratches glass, very hard (covalent network). Pencil graphite: slippery, leaves layers on paper (network of sheets held together by weak forces). Foil and copper: lustrous, bend without breaking (metallic). Explain each result using the particles and forces in the solid.

Your turn

Classify each solid: (a) KBr (b) CO₂(s) (c) diamond (d) iron.

Why do metals conduct electricity as solids, but ionic compounds do not?

Explain why graphite conducts electricity but diamond does not.

Why are metals malleable?

A solid melts at 80 °C, is soft, and does not conduct in any state. What type is it?

Predict the type of solid formed when calcium reacts with chlorine.

Why does MgO have a much higher melting point than NaF, although the ions are similar in size?

More practice: warm-up

What particles make up a metallic solid?

Name a covalent network solid other than diamond.

What holds the molecules together in ice?

Does molten NaCl conduct electricity?

Why is graphite used in pencils?

More practice: on your own

Classify and justify: a shiny solid that conducts electricity, melts at 660 °C and can be rolled into thin sheets.

Silicon dioxide (SiO₂) melts at about 1700 °C but carbon dioxide (CO₂) is a gas at room temperature. Explain.

Explain why steel (an alloy) is harder than pure iron.

Predict the type of solid for (a) SiC (b) P₄ (c) Na₂SO₄ (d) brass.

Rank in order of increasing melting point and name the type of each: Cu, CH₄, NaCl, diamond.

Challenge

Sodium (Na) melts at 98 °C, magnesium at 650 °C and aluminium at 660 °C. Use the metallic bonding model to explain why the melting point rises from Na to Mg.

Boron nitride (BN) can be made in a form with a hexagonal sheet structure like graphite. Unlike graphite, it does not conduct electricity. Suggest why, using electronegativity.

A student finds an unknown white solid that does not dissolve in water, does not conduct, and melts at 210 °C. She concludes it is a covalent network solid because it does not dissolve. Evaluate her conclusion.

Check your understanding

Quiz

Which type of solid conducts electricity in the solid state?

What holds the particles together in a covalent network solid?

Which substance is a molecular solid?

Why are ionic solids brittle?

What type of solid forms when magnesium burns in oxygen?

Graphite conducts electricity because:

For parents and teachers

Big idea (C2.5, C3.4): the properties of a solid follow from its particles and the forces between them. Students should classify solids from formulas and from properties, and predict the type of solid formed in a reaction.

How to help: the kitchen tour uses ordinary materials; please supervise any scratching of glass. Ask your teen to justify each classification with a property and a reason.

Common difficulties:

Classifying SiO₂ as molecular because CO₂ is molecular.
Saying ionic solids conduct (they conduct only when molten or dissolved).
Thinking all non-metal compounds have low melting points (network solids are the exception).