Properties of Ionic Compounds

students, imagine trying to walk across a floor made of tiny magnets all locked together in a repeating pattern ⚡. That is a useful way to start thinking about ionic compounds. In this lesson, you will learn why ionic compounds have high melting points, why they conduct electricity only in certain situations, and why many of them dissolve in water. These ideas connect directly to the structure of matter and to how bonding explains observable properties in IB Chemistry SL.

What are ionic compounds?

Ionic compounds are substances made from positive ions and negative ions held together by strong electrostatic attraction. A positive ion is called a cation, and a negative ion is called an anion. Ionic compounds usually form between metals and non-metals. For example, sodium and chlorine form sodium chloride, written as $\mathrm{NaCl}$.

The important idea is that ionic compounds do not exist as separate molecules. Instead, they form a giant ionic lattice, which is a regular three-dimensional structure of alternating ions. Every ion is surrounded by ions of the opposite charge. This arrangement is very stable because opposite charges attract strongly.

In the lattice, the attraction acts in all directions, not just between one pair of ions. That is why ionic compounds have very different properties from simple molecular substances such as $\mathrm{CO_2}$ or $\mathrm{CH_4}$.

Key terminology

Ion: a charged particle formed by losing or gaining electrons.
Cation: a positively charged ion.
Anion: a negatively charged ion.
Ionic bond: the strong electrostatic attraction between oppositely charged ions.
Lattice: a repeating three-dimensional arrangement of particles.
Lattice energy: the energy released when gaseous ions form an ionic solid, or the energy needed to separate an ionic solid into gaseous ions.

Understanding this structure helps explain the main properties of ionic compounds.

Why ionic compounds have high melting and boiling points

One of the most famous properties of ionic compounds is that many of them melt only at very high temperatures 🔥. For example, sodium chloride has a melting point of about $801\,\mathrm{^b0C}$.

Why? Because the ions in the lattice are held together by strong electrostatic attractions. To melt an ionic compound, enough energy must be supplied to overcome these attractions. Since the attractions act throughout the entire lattice, a large amount of thermal energy is needed.

This does not mean the bonds are being broken in the same way as breaking a molecule into atoms. Instead, the lattice structure is being disrupted so the ions can move more freely. The stronger the ionic attractions, the more energy needed.

Factors affecting melting point

The melting point of an ionic compound depends on the strength of attraction between its ions. This is affected by:

Ion charge: higher charges create stronger attraction.
Ion size: smaller ions can get closer together, increasing attraction.

For example, $\mathrm{MgO}$ has a much higher melting point than $\mathrm{NaCl}$ because $\mathrm{Mg^{2+}}$ and $\mathrm{O^{2-}}$ have larger charges than $\mathrm{Na^+}$ and $\mathrm{Cl^-}$.

This is a major IB Chemistry idea: structure explains property. Stronger ionic attraction means a stronger lattice, and a stronger lattice means a higher melting point.

Why ionic compounds are brittle

Ionic compounds are hard but brittle. That means they resist scratching or denting, but they shatter when struck.

At first, this may seem strange because the ions are strongly held together. The key is what happens when a force shifts the layers of the lattice. If one layer moves, ions with the same charge can become aligned next to each other. Since like charges repel, the lattice breaks apart.

A real-world example is table salt crystals. They are solid and firm, but if crushed, they break into smaller pieces rather than bending like metal. This is because the lattice does not allow layers to slide smoothly.

This brittle behavior is a direct result of the ordered arrangement of ions in the crystal lattice.

Why ionic compounds conduct electricity only when molten or dissolved

A very important property of ionic compounds is electrical conductivity ⚡. Solid ionic compounds do not conduct electricity, but molten ionic compounds and aqueous solutions of ionic compounds do.

Why is this?

Electrical conduction requires moving charged particles. In a solid ionic compound, the ions are fixed in place in the lattice, so they cannot move. Without mobile charge carriers, no current flows.

When the ionic compound is melted, the lattice breaks down and ions become free to move. Similarly, when it dissolves in water, the ions separate and can move through the solution. These mobile ions carry charge, so the liquid conducts electricity.

Example: sodium chloride

In solid $\mathrm{NaCl}$, the ions are locked in place, so the solid does not conduct. In molten $\mathrm{NaCl}$, $\mathrm{Na^+}$ and $\mathrm{Cl^-}$ can move, so it conducts. In salt water, dissolved ions also conduct electricity.

This explains why seawater can carry electrical current and why electrolysis can be used with ionic solutions.

Why many ionic compounds dissolve in water

Many ionic compounds dissolve in water because water is a polar solvent. Water molecules have a partial negative charge near oxygen and partial positive charges near hydrogen atoms.

When an ionic compound is placed in water, water molecules surround the ions. The oxygen end of water attracts cations, and the hydrogen end attracts anions. These attractions help pull ions away from the lattice.

If the attraction between water molecules and the ions is strong enough to overcome the lattice energy, the ionic compound dissolves.

However, not all ionic compounds are equally soluble. Solubility depends on the balance between lattice energy and hydration energy. If the lattice is very strong, the compound may be only slightly soluble or insoluble.

Real-world example

$\mathrm{NaCl}$ dissolves easily in water, which is why salt can be added to soups and used in food preservation. But some ionic compounds, such as $\mathrm{AgCl}$, are very insoluble because their lattice is too strong for water to pull apart effectively.

So, students, when you see an ionic compound, do not automatically assume it dissolves completely. Solubility must be checked using data or solubility rules.

Structure-property relationships in ionic compounds

This lesson fits into the broader IB topic because chemistry often asks: how does microscopic structure explain macroscopic properties?

For ionic compounds:

the giant lattice explains high melting points and boiling points;
the fixed ions in solids explain poor electrical conductivity;
the mobile ions in molten or aqueous states explain conductivity in those conditions;
the rigid lattice explains brittleness;
the interaction with water explains why many ionic compounds dissolve.

This is structure-property reasoning. The properties are not random. They follow from the arrangement of ions and the strength of electrostatic attraction.

Comparing ionic compounds with other bonding models

It helps to compare ionic compounds with covalent substances and metals.

Simple molecular covalent substances often have low melting points because the forces between molecules are weak.
Giant covalent structures such as diamond have very high melting points because many strong covalent bonds must be broken.
Metals conduct electricity as solids because they contain delocalized electrons.
Ionic compounds conduct only when the ions are free to move, such as when molten or dissolved.

These comparisons are a major part of Structure 2 — Models of Bonding and Structure. You are not just memorizing facts; you are linking model and evidence.

How to answer IB-style questions about ionic compounds

When solving exam questions, students, use a clear cause-and-effect structure.

If asked why an ionic compound has a high melting point, write:

It consists of a giant ionic lattice.
Strong electrostatic attractions exist between oppositely charged ions.
A large amount of energy is needed to overcome these attractions.
Therefore, the melting point is high.

If asked why solid ionic compounds do not conduct electricity, write:

Conductivity requires mobile charge carriers.
In the solid, ions are fixed in the lattice.
The ions cannot move.
Therefore, the solid does not conduct.

If asked why molten or aqueous ionic compounds conduct, write:

Melting or dissolving frees the ions.
The ions become mobile.
Mobile ions carry electric current.
Therefore, the substance conducts.

This step-by-step reasoning is exactly the kind of explanation IB Chemistry rewards.

Conclusion

Ionic compounds are a clear example of how bonding and structure determine properties. Their giant ionic lattices cause high melting and boiling points, brittleness, and specific conductivity behavior. Their interaction with water helps explain solubility and the importance of mobile ions in solution. By connecting microscopic ion arrangement to macroscopic observations, you build the core reasoning skill needed in IB Chemistry SL. If you can explain ionic compound properties using the language of lattice structure, electrostatic attraction, and ion mobility, you have mastered the central ideas of this lesson ✅

Study Notes

Ionic compounds are made of cations and anions arranged in a giant ionic lattice.
Ionic bonding is the electrostatic attraction between oppositely charged ions.
Strong attractions in the lattice give ionic compounds high melting and boiling points.
Ionic compounds are brittle because shifting layers can bring like charges together, causing repulsion.
Solid ionic compounds do not conduct electricity because ions are fixed in place.
Molten ionic compounds conduct because the ions are free to move.
Aqueous ionic solutions conduct because dissolved ions are mobile.
Many ionic compounds dissolve in water because polar water molecules attract and separate ions.
Solubility depends on the competition between lattice energy and hydration energy.
Structure-property relationships are central to IB Chemistry SL and help explain observable behavior from particle models.