IB Chemistry HLTopic 2 — Models of Bonding & StructurePaper 1 & 2Trends~9 min read
s-Block and p-Block Elements
Sodium melts at 371 K. Aluminium, two places along the same row, melts at 933 K. Both are metals with the same kind of bonding, so something must be making aluminium’s grip on its electrons far tighter. Two factors explain it, and between them they handle every metallic trend on the syllabus.
📘 What you need to know
Not all metallic bonds are equally strong. Two things control the strength:
The charge on the metal ion — a bigger charge means more delocalised electrons and a bigger charge difference, so stronger attraction.
The radius of the metal ion — a smaller ion holds the electron sea more tightly, so stronger attraction.
Across a period the charge increases and the radius decreases, so metallic bonding gets stronger and melting points rise.
Down a group the charge stays the same but the radius increases, so metallic bonding gets weaker and melting points fall.
Only the first three elements of period 3 (Na, Mg, Al) are metallic — silicon is a covalent network and the rest are simple molecular.
These are trends, not laws. Packing structure causes small deviations that the IB does not require you to explain.
The two factors
Metallic bonding is an electrostatic attraction, so the same logic applies as anywhere else with charges: bigger charges attract more strongly, and charges closer together attract more strongly.
⚡
MORE CHARGE
Al3+ releases three electrons into the sea, so the charge difference either side is large.
📏
BIGGER RADIUS
Cs+ is huge and gives up only one electron, so its hold on the sea is feeble.
The rule in one line
higher charge and smaller radius → stronger electrostatic attraction → stronger metallic bond → higher melting point
Across period 3
Sodium, magnesium and aluminium form Na+, Mg2+ and Al3+. Look at their electron configurations and you can see where the extra electrons come from:
Na = 1s2 2s2 2p6 3s1 — one outer electron to give
Mg = 1s2 2s2 2p6 3s2 — two
Al = 1s2 2s2 2p6 3s2 3p1 — three
So going across, the ionic charge climbs from 1+ to 3+, the number of delocalised electrons per atom climbs from 1 to 3, and the ionic radius gets smaller because the growing nuclear charge pulls the remaining electrons in. All three changes push in the same direction: stronger bonding.
Melting points across period 3
Do not try to explain the whole chart with metallic bonding. From silicon onwards a different structure takes over, and that is what the huge silicon bar is really telling you.
Notice that aluminium is barely above magnesium — 933 against 923 K. If you had only the theory you would expect a much bigger jump. This is a good reminder that these are broad trends, and that other factors such as how the atoms pack can shift things slightly. The IB does not ask you to explain those; it asks you to explain the trend.
Down group 1
Going down a group the charge on the ion stays fixed at 1+, so only one thing changes: the ions get bigger as each element adds another shell.
A bigger ion means the delocalised electrons sit further from the nucleus, with more inner shells shielding them. The attraction weakens, less energy is needed to break the lattice apart, and the melting point falls steadily from lithium to caesium.
Melting points down group 1
Caesium melts at 302 K, which is 29°C — below body temperature. Hold a sealed ampoule of it and it turns to liquid, which is a strikingly direct demonstration of weak metallic bonding.
Direction
Ionic charge
Ionic radius
Delocalised electrons
Metallic bonding
Melting point
Across a period
increases
decreases
increases
stronger
rises
Down a group
stays the same
increases
stays the same
weaker
falls
🧠 How to structure the answer every time
Three sentences, in this order. One: state what happens to the charge and the radius. Two: say what that does to the electrostatic attraction between the cations and the delocalised electrons. Three: say what that means for the energy needed, and therefore the melting point. Get those three in and you have the marks.
Worked examples
WE 1
Explain why magnesium has a much higher melting point than sodium [3]
Mark 1: the ions and the electrons
Mg forms Mg²⁺ and releases 2 delocalised electrons per atom; Na forms Na+ and releases only 1.
Mark 2: the size
Mg2+ is also smaller than Na+, because of the greater nuclear charge.
Mark 3: link to the attraction and the energy
Both factors give a stronger electrostatic attraction between the cations and the electron sea, so more energy is needed to break the lattice.
371 K → 923 Kall three of charge, radius and electron count are creditable — give as many as you can
WE 2
Explain the trend in melting points from Li to Cs [3]
Mark 1: state the trend and what stays constant
Melting point decreases down the group. Every ion has the same 1+ charge and gives one delocalised electron.
Mark 2: what does change
The ionic radius increases as each element adds a shell, and shielding increases.
Mark 3: consequence
The delocalised electrons are further from the nucleus, so the attraction is weaker and less energy is needed.
454 K down to 302 Ksaying what stays the same is what proves you understand which factor is doing the work
WE 3
Silicon has by far the highest melting point in period 3. Explain why this is not evidence of very strong metallic bonding.
Step 1: identify silicon’s structure
Silicon is not a metal. It forms a giant covalent network, with each atom bonded to four others.
Step 2: say what must be broken to melt it
Melting requires breaking many strong covalent bonds throughout the lattice, not overcoming metallic attraction.
Step 3: draw the conclusion
The high value comes from a different bonding model entirely, so it cannot be compared with Na, Mg and Al.
Giant covalent, not metallic — a different explanation appliesonly the first three period 3 elements are metallic; check this before you explain any trend
💡 Exam tips
Name both factors where they apply: charge on the ion and radius of the ion.
Always finish by linking to the energy needed to overcome the attraction. That is the melting point sentence.
Down a group, say explicitly that the charge is unchanged so radius is the only variable.
Check the structure before explaining any melting point — only Na, Mg and Al in period 3 are metallic.
Use the phrase delocalised electrons, and mention how many each atom contributes.
If asked about a small anomaly like Al versus Mg, say these are trends rather than rules.
⚠ Common mix-ups
Explaining silicon’s melting point with metallic bonding. Silicon is a covalent network.
Saying “the metallic bond gets stronger” without saying why. Name the charge and the radius.
Confusing atomic radius with ionic radius. It is the cation in the lattice that matters.
Claiming the charge increases down a group. It does not — only the size does.
Using “molecules” for metals. A metallic lattice has no molecules in it.
Expecting a perfectly smooth trend. Al is only just above Mg, and that is fine.
Up next: Physical Properties of Transition Elements — the metals that break the pattern entirely, melting hundreds of degrees higher than anything in the s-block.
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