Ionic, covalent, metallic — three names, three diagrams, and for most students, three disconnected things to memorise. They're actually three different solutions to the exact same problem: every atom is more stable with a full outer electron shell, and bonding is just the different strategies atoms use to get there, depending on what they're bonding with.
GCSE coreIonic bonding: give it away entirely
When a metal atom (with few outer electrons, like sodium's one) meets a non-metal atom that's nearly full (like chlorine's seven), the easiest route to a stable outer shell for both is a straight transfer: sodium gives its one outer electron to chlorine completely. That leaves sodium as a positively-charged ion (it lost a negative electron) and chlorine as a negatively-charged ion (it gained one) — and opposite charges attract, which is the ionic bond itself.
Ionic bonding transfers an electron completely, creating two oppositely-charged ions that attract. Covalent bonding shares electrons between two atoms instead — neither one fully gives or takes.
Covalent bonding: share instead of give
When two non-metal atoms bond — like the hydrogen and oxygen in water — neither one has spare electrons to simply hand over; both need to gain, not lose. The solution is sharing: a pair of electrons sits between the two nuclei, counting toward a full outer shell for both atoms simultaneously. Water forms two of these shared pairs — one between oxygen and each hydrogen atom — which is exactly why its formula is H₂O, not H₃O or HO.
Exam-style question
Explain why sodium chloride (NaCl) has a high melting point, but water has a much lower melting point, even though water molecules also contain strong bonds. [4 marks]
Show model answer and mark breakdown
- Sodium chloride is a giant ionic lattice — every ion is strongly, electrostatically attracted to its neighbours in every direction (1 mark)
- Melting it means overcoming this whole network of attractions throughout the lattice at once, which needs a lot of energy (1 mark)
- Water is made of individual small molecules; the covalent bonds within each molecule are strong, but the forces between separate water molecules are much weaker (1 mark)
- Melting water only requires overcoming those weaker forces between molecules, not breaking the strong covalent bonds inside them (1 mark)
Examiner focus: keep the bonds within a molecule (strong, covalent) clearly separate from the forces between molecules (weak) — this is the exact distinction the question is testing.
Common mistake: saying water has a low melting point because its covalent bonds are weak — the O–H covalent bonds are actually very strong; it's the much weaker forces between separate molecules that are overcome on melting.
Metallic bonding: share with everyone at once
Metal atoms don't bond in pairs at all. Their outer electrons detach and move freely through the entire structure, leaving behind a lattice of positive metal ions held together by that shared, moving "sea" of electrons around them. This is why metals are malleable (the ions can shift position without breaking a specific bond to a specific neighbour) and why they conduct electricity so well — those free electrons can carry a current through the whole piece of metal. If ionic, covalent and metallic bonding need explaining as one connected idea rather than three separate diagrams to memorise, that's exactly what our GCSE chemistry tutoring is for — see the full learning pathway here.
Frequently asked questions
What's the actual difference between ionic and covalent, in one sentence?
Ionic bonding transfers electrons completely from one atom to another, creating two oppositely-charged ions that attract; covalent bonding shares electrons between atoms instead of giving them away, with neither atom fully gaining or losing anything.
Why does metallic bonding need its own category, separate from ionic and covalent?
In a metal, outer electrons aren't transferred to a specific other atom or shared between two specific atoms — they detach from their original atoms entirely and move freely through the whole structure, a 'sea of electrons' shared collectively by all the positive metal ions around them. That free-moving electron sea is also why metals conduct electricity so well: those electrons can carry a current through the whole structure.
Why do ionic compounds have such high melting points?
Ionic bonds are electrostatic attractions between oppositely-charged ions, and that attraction is strong and acts in every direction throughout a rigid 3D lattice of ions — melting the compound means overcoming that whole network of attractions at once, which takes a lot of energy, hence the high melting point.
Reader ratings & reviews
Rate this article across a few different things — your rating helps other readers, and helps us improve.
Was this article helpful?
Tell us what you think — a correction, a question it left unanswered, or a topic you'd like covered next.
Prefer To Talk? Book a free video call at a time that suits you: 15 minutes for a quick question, 30 for anything else. Times show in your own time zone.
About the author
Sudershan Soni
Founder & Lead Tutor at Mostak Services — an MSc-qualified Mathematics, Science, Computer Science & STEM tutor with 20+ years of professional experience, teaching students from 11+ and GCSE to A-Level and beyond, online worldwide.
Read full profile