Not all metallic bonds are equally strong. Sodium melts at 371 K but aluminium needs 933 K. The difference comes down to three things about the metal ion — its charge, its size, and how many electrons it released.
📚 What you need to know
Metallic bond strength depends on the charge on the ion, the radius of the ion, and the number of delocalised electrons.
Higher charge → stronger attraction to the electron sea → stronger bond.
Smaller radius → the electrons sit closer to the nucleus → stronger bond.
More delocalised electrons per atom → stronger bond.
Across Period 3 (Na → Mg → Al) all three factors improve, so melting point rises.
Down Group 1 the charge stays at 1+ but the ion gets bigger, so melting point falls.
What makes a metallic bond strong
Remember what the bond actually is: an attraction between positive ions and the negative electron sea. Anything that makes that attraction stronger makes the metal harder to melt.
The charge on the ion. An Al3+ ion pulls on the surrounding electrons three times as hard as a Na+ ion does. A bigger charge difference means a stronger electrostatic attraction.
The radius of the ion. A smaller ion lets the delocalised electrons get closer to the positive nucleus, and closer means stronger.
The number of delocalised electrons. Each aluminium atom donates three electrons to the sea, sodium only one. A denser electron sea means more attraction to go round.
From sodium to aluminium the ion shrinks, its charge grows and it releases more electrons into the sea — all three strengthen the bond.
These are the same two factors — charge and size — that decided ionic lattice strength back in Topic 2. Electrostatic attraction always works the same way, whatever the bonding. If you can explain why MgO melts higher than NaCl, you can already explain why aluminium melts higher than sodium.
Across a period: Na → Mg → Al
Look at the electron configurations and you can see all three factors improving at once:
Metal
Electron configuration
Ion formed
Delocalised electrons per atom
Melting point / K
Sodium
1s22s22p63s1
Na+
1
371
Magnesium
1s22s22p63s2
Mg2+
2
923
Aluminium
1s22s22p63s23p1
Al3+
3
933
Going from sodium to aluminium, the charge climbs from 1+ to 3+, the number of delocalised electrons climbs from 1 to 3, and the ion actually gets smaller (more protons pulling on the same shell). Every factor points the same way, so the metallic bonding gets much stronger.
Plotting the whole of Period 3 makes the effect obvious — but watch which bars you are allowed to use:
The blue bars are the metals. Silicon towers over them, but that is giant covalent bonding and belongs to a different explanation.
💡 Only the first three are metals
Na, Mg and Al are the only metallically bonded elements in Period 3.
Si has the highest melting point of all, but that is giant covalent bonding, not metallic — different topic, different explanation.
P, S, Cl and Ar are simple molecular or monatomic, held by weak intermolecular forces, so they melt very low.
If a question asks you to explain the trend “for the metals”, stop at aluminium.
Down a group: Li → Cs
Down Group 1 the picture reverses. Every one of these metals forms a 1+ ion and donates exactly one delocalised electron, so charge and electron count are fixed. The only thing that changes is size.
Each element down the group has an extra electron shell, so the ion gets steadily bigger. The delocalised electrons are held further from the nucleus and are more shielded, so the attraction weakens — and the melting point falls.
Charge and electron count are identical all the way down, so the steady fall is caused by the ions getting bigger.
WORKED EXAMPLE
Explain why magnesium has a much higher melting point than sodium.
Mg forms Mg²⁺ but Na forms only Na⁺Higher charge on the ion.Each Mg atom donates 2 electrons, each Na atom only 1More delocalised electrons in the sea.The Mg²⁺ ion is also smaller than Na⁺stronger electrostatic attraction → more energy needed → higher melting point
WORKED EXAMPLE
Potassium melts at 337 K but lithium melts at 454 K. Both are in Group 1. Explain the difference.
Both form 1+ ions and donate 1 electron eachSo charge and electron count cannot be the reason.K⁺ has an extra electron shell, so it is bigger than Li⁺The delocalised electrons sit further from the nucleus and are more shielded.weaker attraction in potassium → lower melting point
A word of caution about aluminium. Look closely at the Period 3 chart and aluminium is only just above magnesium — 933 K against 923 K — even though it has one more charge and one more delocalised electron. These are trends, not rules. Other factors, such as how the ions pack together in the lattice, also affect melting point. That detail is beyond the IB syllabus; you just need to explain the broad trend and not be thrown when the numbers are closer than expected.
💡 Exam tip
A full-mark answer names all three factors where they apply: charge, size, number of delocalised electrons.
Always finish with the consequence: stronger attraction → more energy needed → higher melting point.
For a group trend, say clearly that charge is unchanged so size is the deciding factor. That comparison is worth a mark by itself.
Use the term delocalised electrons, not “free electrons”.
⚠️ Common mix-up
Silicon is not a metal. Its huge melting point is giant covalent bonding — don’t use it as evidence about metallic bonds.
Don’t say “bigger atom, stronger bond”. It is the opposite: bigger ion means the electrons are further away and the bond is weaker.
Atomic radius and ionic radius are not the same. When discussing metallic bonding, talk about the size of the ion.
Melting a metal breaks the metallic bonding — the attraction between ions and electrons. There are no covalent bonds to break.
That completes Metallic Bonding, and with it all three bonding models — ionic, covalent and metallic. You now have everything you need to look at any substance, work out its structure, and predict how it will behave.
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