One column, three personalities
You met boron in the previous guide as the oddball at the top of Group 13 — too small and too electron-poor to behave like a metal, forced into electron-deficient clusters and behaving as a Lewis acid. Now we descend the rest of the column: boron, aluminum, gallium, indium, thallium. Every one of them has the same outer arrangement, ns2 np1, so the textbook expectation is a tidy +3 oxidation state across the board. The reality is far more interesting, and it is the perfect place to watch a single periodic trend — [[metallic-character|metallic character]] — strengthen as you slide down a group.
Boron is a hard, brittle metalloid; it forms no simple B3+ ion at all, because tearing three electrons off so small an atom costs far more energy than any lattice or hydration can ever repay. Aluminum just below it is a true light metal — shiny, conducting, structural — yet its bonds still carry strong covalent character. Gallium and indium are unmistakably metallic (gallium famously melts in your warm hand at about 30 degrees Celsius), and thallium at the bottom is a soft, dense, heavy metal. So a single group spans the full metal-to-nonmetal gradient: nonmetal-ish at the top, solidly metallic at the bottom. The driving force is simple — going down, atoms get bigger and their outer electrons are held more loosely, so they let go of electrons more readily, which is exactly what "metallic" means at the chemical level.
Aluminum: the two-faced oxide
Aluminum is the most abundant metal in the Earth's crust, yet you almost never see it corrode away the way iron rusts. The secret is its oxide. When fresh aluminum meets air it grows a thin, tough, transparent skin of Al2O3 that clings to the metal and seals it — passivation. That same oxide sits right on the dividing line between basic and acidic, so it answers to both: it neutralizes acids like a base (Al2O3 + 6 H+ gives 2 Al3+ + 3 H2O) and dissolves in strong alkali like an acid (Al2O3 + 2 OH- + 3 H2O gives 2 [Al(OH)4]-). This double behaviour is [[amphoteric-aluminum-oxide|amphoterism]], and you saw it earlier in the acid-base rung; here it is the chemical fingerprint of an element straddling the metal-nonmetal border.
Why is the oxide on the fence rather than firmly basic like Na2O or MgO? Because the Al3+ ion packs a +3 charge into a small radius, giving it a fierce electric pull — high charge density. By charge and size bookkeeping this small, highly charged cation polarizes nearby oxide and hydroxide ions strongly, dragging electron density toward itself and giving the bonds real covalent character. A bond that is partly covalent can be pulled apart from either side: a base can pry the metal loose, an acid can pry the oxide loose. Sodium's big, gently charged Na+ does no such polarizing, so Na2O stays purely basic. The amphoterism of aluminum is therefore not a quirk — it is charge density made visible.
Al2Cl6: a molecule holding hands
Now look at aluminum's most famous covalent quirk. A bare AlCl3 molecule would have aluminum surrounded by only three chlorines — six electrons, two short of an octet. Aluminum hates that electron shortfall just as boron does, so it is a strong [[inorg-lewis-acid-base|Lewis acid]], hungry for an electron pair. With nothing better around, two AlCl3 units solve the problem by pairing up: each aluminum borrows a lone pair from a chlorine on the other molecule. The result is the dimer [[aluminum-chloride-dimer|Al2Cl6]], in which two chlorines sit in bridging positions, each one bonded to both aluminums at once. Those bridges donate the missing electron density, and each aluminum finally reaches a full octet in a roughly tetrahedral environment.
Cl Cl Cl
\ / \ /
Al Al terminal Cl: ordinary 2-electron Al-Cl bonds
/ \ / \ bridging Cl: each lone pair shared with BOTH Al
Cl Cl Cl
^^
the two BRIDGING chlorines
count per Al: 2 terminal Al-Cl + 2 bridge Al-Cl(donated) -> octet, ~tetrahedralTwo honest notes. First, the bridge bonds here are ordinary two-electron dative bonds — a chlorine simply donates a lone pair it already had spare. That is different from boron's bridges in diborane B2H6, where two electrons must be smeared over a three-atom B-H-B span (a genuine three-centre two-electron bond) because hydrogen has no lone pair to give. Aluminum's chlorines do have lone pairs, so Al2Cl6 needs no such exotic bonding. Second, this dimer is what you get in the vapour or in nonpolar solvents; in water or any donor solvent, those donors out-compete the bridges, AlCl3 falls apart, and you are back to the hydrated [Al(H2O)6]3+ aqua ion. The Lewis acidity that builds the dimer is the very same property exploited in industry, where AlCl3 is the classic catalyst for Friedel-Crafts reactions.
Gallium, indium, and the lazy electron pair
Keep descending. Gallium and indium are still mostly +3 elements — GaCl3 and InCl3, Ga2O3 and In2O3 are the everyday compounds — but a second oxidation state, +1, starts to appear and grows more comfortable lower down. By the time you reach thallium at the bottom, the situation has flipped: thallium's +1 state (the thallous ion Tl+) is more stable than its +3 state (thallic, Tl3+). Tl3+ is a strong oxidizer, itching to grab two electrons and drop back to Tl+. The element happily forms Tl+ salts that look almost alkali-metal-like — TlOH is a strong, soluble base. So the group that began with an element too electron-poor to lose even one electron easily, ends with one that would rather keep two of them.
What stabilizes that +1 state is the [[inert-pair-effect|inert-pair effect]]: in the heaviest p-block elements the outer ns2 pair tends to stay put rather than join in bonding, so the element shows an oxidation state two below its group maximum. The name suggests the pair is intrinsically lazy, but the real story is an energy balance. Down a heavy group the ns electrons are held a little tighter (a relativistic tug on the inner s electrons, plus poor shielding by the bloated d and f shells beneath), so it costs more to use them — and at the same time the bonds the element forms get weaker, so there is less energy returned to pay that cost. When the books no longer balance, the element simply leaves the pair alone. Honest caveat: the pair is not literally inert and the relativistic part matters mostly for the very heaviest elements; "inert-pair effect" is a convenient label for that whole tug-of-war, not a force in itself.
Reading the gradient end to end
Step back and the whole column tells one connected story. The oxides soften as you go down: B2O3 is acidic (it gives boric acid), Al2O3 and Ga2O3 are amphoteric, and by thallium Tl2O is basic — exactly tracking the growing metallic character. Notice how oxidation state and oxide character move together: the high +3 state favours acidic-to-amphoteric oxides, while the low +1 state of heavy thallium gives a frankly basic one, just as it would for an alkali metal. The same charge-density logic that made Al2O3 amphoteric, applied to the much smaller, harder Tl3+, is part of why high-valent heavy-element oxides lean acidic.
It is worth naming a misconception while we are here. "Inorganic" chemistry is not the lifeless half of the subject, and Group 13 makes the point: aluminum's compounds run our packaging, transport, and antacids; gallium arsenide and indium phosphide are the semiconductors behind LEDs and fast electronics; gallium nitride drives modern efficient lighting. And remember that an oxidation state like Tl(+1) or Al(+3) is a bookkeeping device — a way of tracking electrons on paper — not a literal count of charges sitting on the atom. The Al-Cl bonds in Al2Cl6 are sharply covalent even though we cheerfully write aluminum as +3.