B. Matter, Chemical Trends, and Chemical Bonding · Matter, Chemical Trends, and Chemical Bonding

Properties of Ionic and Molecular Compounds

B3.5 B3.4 A1.5 A1.6 A1.10About 70 minutes

Learning goals

I can compare the melting points, hardness, conductivity and solubility of ionic and molecular compounds.
I can explain these properties using ionic bonds and intermolecular forces.
I can explain why ionic compounds are brittle and when they conduct electricity.
I can use properties to classify an unknown compound as ionic or molecular.

Words to know

intermolecular forces
Attractions between separate molecules; much weaker than covalent or ionic bonds.

melting point
The temperature at which a solid changes to a liquid.

brittle
Breaking or shattering, rather than bending, when struck.

electrolyte
A substance that conducts electricity when melted or dissolved in water, because it forms mobile ions.

non-electrolyte
A substance whose solution does not conduct electricity.

dissociation
The separation of an ionic compound into its ions when it dissolves.

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

Salt versus sugar

Table salt and white sugar look almost the same: small, white, crystalline grains. Both dissolve in water. But heat them in a pan and they behave completely differently. Sugar melts at about 186 °C and then turns brown (caramel). Salt does not melt at all on a kitchen stove: it needs 801 °C. Dissolved in water, salt lets electricity flow; sugar does not.

Salt (sodium chloride) is an ionic compound. Sugar (sucrose, C₁₂H₂₂O₁₁) is a molecular compound. In this lesson you will compare the physical properties of these two kinds of compounds and explain the differences using what you know about bonding.

What holds the solid together?

The key to this whole lesson is asking what must be overcome to melt or boil a substance.

In an ionic solid, every ion is held by strong ionic bonds to all its neighbours throughout the lattice. To melt it, you must break apart a huge network of strong attractions. That takes a lot of energy, so the melting point is high.
In a molecular solid, the covalent bonds inside each molecule are strong, but they are not broken when the substance melts. Melting only separates molecules from each other, and molecules are held to each other by much weaker intermolecular forces. So melting and boiling points are low.

When water boils, the steam is still made of H₂O molecules; the O–H bonds are not broken. You will study the different kinds of intermolecular forces (and why water's are unusually strong) in the Gases and Atmospheric Chemistry unit.

Property	Ionic compounds	Molecular compounds
State at room temperature	solids (crystals)	solids, liquids or gases
Melting point	high (often several hundred °C or more)	low (usually below about 300 °C)
Hardness	hard but brittle: shatter when struck	usually soft solids
Conducts electricity as a solid?	no	no
Conducts when melted or dissolved in water?	yes (ions are free to move)	no (no ions), except acids, which form ions in water
Solubility in water	many are soluble	varies: polar molecules (sugar) often dissolve; non-polar ones (wax, oil) do not

Typical properties of ionic and molecular compounds.

Hard but brittle

Ionic crystals are hard because strong attractions hold every ion in place. But they are also brittle: hit a salt crystal and it shatters instead of bending. A sharp blow shifts one layer of ions slightly. Suddenly positive ions sit next to positive ions and negative next to negative. The repulsion between like charges pushes the layers apart and the crystal cracks along a flat surface.

Why ionic crystals are brittle: shifting a layer puts ions with the same charge next to each other, and their repulsion splits the crystal.

Electrical conductivity

Electricity can flow through a substance only if it contains charged particles that are free to move.

Solid ionic compounds do not conduct: the ions are charged but locked in the lattice.
Molten (melted) ionic compounds and aqueous solutions of ionic compounds do conduct: the lattice breaks apart and the ions can move. Solutions that conduct are called electrolytes. (Breaking an ionic compound into separate ions when it dissolves is called dissociation.)
Molecular compounds do not conduct in any state, because their particles are neutral molecules. Their solutions are non-electrolytes. The important exception is acids (such as HCl), which are molecular but react with water to form ions.

A conductivity tester lights only when charged particles can move. Dissolved salt releases mobile ions; solid salt and sugar solution do not conduct.

Substance	Type	Melting point (°C)	Conducts when dissolved or melted?
sodium chloride, NaCl	ionic	801	yes
magnesium oxide, MgO	ionic	2852	yes (molten)
calcium chloride, CaCl₂	ionic	772	yes
water, H₂O	molecular	0	no (pure water barely conducts)
sucrose, C₁₂H₂₂O₁₁	molecular	about 186 (decomposes)	no
iodine, I₂	molecular	114	no
methane, CH₄	molecular	−182	no

Melting points of some ionic and molecular compounds.

Watch out: network solids

A few substances with covalent bonds have very high melting points: diamond (carbon) and quartz (silicon dioxide, SiO₂, about 1700 °C). They are not made of small molecules. Instead, every atom is covalently bonded to its neighbours in a giant network, so melting means breaking strong covalent bonds. You will study these network solids in SCH4U. For now, use the properties in this lesson for compounds made of small molecules.

Worked example

Example 1: Explaining a melting point difference

Potassium bromide (KBr) melts at 734 °C. Bromine (Br₂) melts at −7 °C. Explain the difference.

KBr is ionic (K is a metal, Br a non-metal; ΔEN = 3.0 − 0.8 = 2.2).
Melting KBr means overcoming strong attractions between K⁺ and Br⁻ ions throughout the lattice: high melting point.
Br₂ is molecular. Melting it only separates Br₂ molecules from each other, held by weak intermolecular forces. The Br–Br covalent bond is not broken.

KBr's ions are held by strong ionic bonds; Br₂ molecules are held to each other only by weak intermolecular forces.

Worked example

Example 2: Identifying an unknown

An unknown white solid melts at 776 °C. It does not conduct electricity as a solid, but its water solution does. Is it ionic or molecular? Explain.

High melting point → strong forces throughout the solid, typical of an ionic lattice.
Not conducting as a solid but conducting in solution → it contains ions that become free to move when it dissolves.
Both clues agree.

The solid is ionic. (Potassium chloride, KCl, which melts at about 770 °C, fits this description.)

Worked example

Example 3: Predicting properties

Predict the state at room temperature, relative melting point and conductivity of (a) calcium fluoride, CaF₂, and (b) ammonia, NH₃.

(a) CaF₂: metal + non-metal, ΔEN = 3.0 → ionic. Predict: solid, high melting point, conducts when molten (it is not very soluble in water). Actual melting point: 1418 °C.
(b) NH₃: two non-metals → molecular. Predict: low melting and boiling points; does not conduct as a pure substance. Actual: a gas at room temperature (boils at −33 °C).

(a) Ionic: hard solid, high melting point, conducts when molten. (b) Molecular: gas, low melting point, no conductivity as a pure substance.

Just for fun

Why don't ionic compounds make good comedians?

They crack under pressure.

Hands-on activity

Salt versus sugar showdown

With a grown-up. You need: table salt, white sugar, two clear glasses of warm water, two spoons, a magnifying glass (optional), two small squares of aluminium foil folded up at the edges, a frying pan and a stove. The pan, foil and melted sugar get very hot: a grown-up handles the stove and pan, and nobody touches or tastes the melted sugar.

Look: examine a few grains of each with the magnifying glass. Salt grains are cubes; sugar crystals have a different, slanted shape.
Hardness: press a few grains of each with the back of a spoon on a cutting board. Both crunch. Describe any difference.
Solubility: stir 1 teaspoon of each into its own glass of warm water. Both dissolve. (Both are soluble: salt because water attracts its ions; sugar because its molecules are polar.)
Melting (grown-up does this): put a small pinch of each on its own foil square in the pan. Heat on medium-low for a few minutes. Watch from a safe distance. Turn off the heat and let everything cool completely before cleaning up.
Conductivity: you cannot safely test this at home. Your teacher may use a conductivity tester in class; use the diagram above to predict the result.

What to look for: the sugar melts into a clear liquid and then turns brown; the salt does not change at all. Write a conclusion that explains the result using the words ionic bonds, intermolecular forces and melting point.

Your turn

Why do ionic compounds have high melting points?

When ice melts, are the covalent bonds in the water molecules broken? Explain.

Why does solid NaCl not conduct electricity, while molten NaCl does?

Explain why ionic crystals shatter when hit.

Classify as ionic or molecular from the properties: (a) melts at −95 °C; (b) melts at 1074 °C and conducts when molten.

Sugar dissolves in water but the solution does not conduct. Explain.

More practice: warm-up

Do molecular compounds generally have high or low melting points?

What is an electrolyte?

Does solid KCl conduct electricity?

Name one molecular compound that is a gas at room temperature.

What word describes a solid that shatters instead of bending?

More practice: on your own

Arrange in order of increasing melting point and explain: MgO, CH₄, NaCl.

Hydrogen chloride is a molecular gas, yet hydrochloric acid conducts electricity. Explain.

A student says, "Paraffin wax melts at a low temperature because its covalent bonds are weak." Correct the statement.

Predict whether CaO or KCl has the higher melting point, and explain.

Why can a conductivity test of the water solution be used to tell sugar from salt, but not tell salt from another soluble ionic compound?

Challenge

Design an investigation to classify three unknown white solids as ionic or molecular, using a hot plate, a conductivity tester, distilled water and beakers. Include one safety precaution.

Silicon dioxide (SiO₂) contains only non-metals but melts at about 1700 °C. Explain why it does not fit the molecular-compound pattern.

Ionic compounds are brittle, but metals (also held by electrostatic attractions) bend. Using the idea of a "sea of electrons", suggest why.

Check your understanding

Quiz

Which property is typical of ionic compounds?

When a molecular solid melts, what is overcome?

A substance melts at 45 °C and does not conduct electricity in any state. It is most likely

Why does molten NaCl conduct electricity?

Ionic crystals are brittle because

Which compound has the highest melting point?

A solution of glucose (C₆H₁₂O₆) in water

For parents and teachers

Big idea: Expectation B3.5 asks students to compare and contrast the physical properties of ionic and molecular compounds. The explanation always comes back to what must be overcome: strong ionic bonds throughout a lattice, or weak forces between molecules.

Safety: in the activity, a grown-up does all heating. Melted sugar is extremely hot and sticks to skin. Let everything cool before cleaning up.

How to help: Ask your teen to explain the salt-and-sugar results in their own words, using "ionic bonds" and "intermolecular forces". Kitchen examples work well: salt does not melt in a pan, butter and chocolate (mixtures of molecular compounds) melt easily.

Common difficulties:

Saying covalent bonds break when a molecular substance melts or boils. They do not.
Thinking ionic solids conduct because they contain ions. The ions must be free to move.
Assuming everything with covalent bonds melts easily. Network solids such as quartz are the exception.