Ionic, Covalent & Metallic Bonding Explained (MYP Chemistry)
Why does salt dissolve and conduct when melted, while sugar does not? Why is diamond hard but graphite slippery? The answer to nearly every "properties" question in chemistry is bonding: what the particles are and how they are held together. Learn to identify the type of bonding and the properties follow.
Try it yourself: work through ionic, covalent, giant structures and metallic bonding, earning badges as you go.
Ionic bonding
Ionic bonding happens between a metal and a non-metal. The metal atom gives away electrons to become a positive ion and the non-metal atom gains them to become a negative ion. The strong electrostatic attraction between opposite charges holds a giant lattice together, which is why ionic compounds have high melting points and are hard but brittle. They conduct electricity only when molten or dissolved, because only then can the ions move.
Covalent bonding
Covalent bonding happens between non-metal atoms, which share pairs of electrons. Simple molecules such as water or carbon dioxide have strong bonds inside each molecule but only weak forces between molecules, so they have low melting and boiling points and do not conduct electricity — there are no free charged particles.
Giant covalent structures
Some covalent substances are one enormous network of bonded atoms. In diamond each carbon atom bonds to four others, making it extremely hard with a very high melting point. In graphite each carbon bonds to three others in layers, leaving spare electrons that move, so graphite conducts electricity, and the weak forces between layers let them slide.
Metallic bonding
A metal is a lattice of positive ions in a "sea" of delocalised electrons. The electrons move freely, so metals conduct electricity and heat, and because layers of ions can slide over each other without breaking the bonding, metals are malleable and ductile.
Linking structure to properties
In an exam, a full explanation names the particles, the bonding, and then the property: "graphite conducts because it has delocalised electrons that can move through the layers." Naming the property without the particle-level reason will not reach the top band.
Frequently asked questions
The ions are locked in a lattice in the solid, but can move freely when molten or dissolved, carrying charge.
Each carbon in graphite uses only three of its four outer electrons for bonding, leaving delocalised electrons that can move. In diamond all four are used in bonds, so none are free.
Layers of positive ions can slide over each other while the delocalised electrons keep holding the structure together.