IB DP Chemistry Structure 2.3: The metallic model

Here you will learn Structure 2.3, The metallic model, as described by the IB Diploma Programme Chemistry guide (February 2023 edition; first examined in 2025). The page works through Structure 2.3.1, 2.3.2 and 2.3.3 in order. The first two are studied at both SL and HL. The third, on transition elements, sits in the additional higher level (AHL) part of the sub-topic, so it is HL only. The guiding question asks what decides whether an element behaves as a metal, and what gives metals their properties.

Two companion pages sit alongside this one: condensed metallic model revision notes and original metallic model practice questions. To see this unit in context, open the IB DP chemistry course page or print the chemistry topic list. A short diagnostic quiz is a quick way to check what you already know.

What this unit covers

Guide reference	What you should be able to do	Level
Structure 2.3.1	Describe the metallic bond; explain electrical conductivity, thermal conductivity and malleability; link typical metal properties to the jobs metals do	SL and HL
Structure 2.3.2	Explain why metallic bond strength depends on ion charge and ion radius; use this to explain melting point trends in s-block and p-block metals	SL and HL
Structure 2.3.3	Explain, using delocalized d-electrons, why transition elements have high melting points and conduct electricity	HL only

The metallic bond (2.3.1)

Metal atoms hold their outer electrons loosely. In a solid metal, each atom releases its valence electrons. These electrons no longer belong to any one atom. They spread through the whole structure and are described as delocalized.

What is left behind is a positive ion, a cation. The cations are packed closely in a regular, repeating lattice.

Metallic bonding is the electrostatic attraction that holds a lattice of cations to the delocalized electrons spread around them.

Three features of this model explain almost everything else in the unit.

• The attraction acts in all directions. No electron pair links one particular atom to another, so the bonding is non-directional.
• The delocalized electrons are mobile. They can drift through the lattice.
• The bond is a collective attraction. Every cation is attracted to the electron cloud around it, not to a single partner.

A sample of metal contains no molecules and no fixed ratio of ions to electrons that you write as a formula. You write it with its element symbol, such as Ca or Rb.

Worked example 1: describing the bonding in calcium

Describe the structure and bonding in solid calcium. Calcium is in group 2.

Electron configuration: [Ar] 4s²
Each atom releases its two 4s electrons into the lattice
Particles left in the lattice: Ca²⁺ cations, closely packed in a regular arrangement
Delocalized electrons: two per atom, free to move through the whole structure
Bond: electrostatic attraction between the Ca²⁺ lattice and the delocalized electrons

Notice the two parts of a full answer: what the particles are, and what attracts what.

Explaining metal properties (2.3.1)

Electrical conductivity

An electric current is a flow of charge. A metal conducts because its delocalized electrons can move. When a potential difference is applied across a piece of metal, the electrons drift towards the positive terminal, and electrons enter from the negative terminal to replace them. The cations do not move. They stay at their lattice sites.

This works in the solid and in the liquid. Mercury, a metal that is liquid at 298 K, still conducts, because melting does not remove the delocalized electrons.

Thermal conductivity

When one end of a metal is heated, the delocalized electrons there gain kinetic energy. Because they are mobile, they carry that energy quickly to cooler parts of the metal and pass it on through collisions. Vibrations of the closely packed cations also pass energy along. The mobile electrons are the reason metals conduct heat far better than most non-metals.

Malleability

A malleable material can be hammered or pressed into a new shape without breaking. A related property, ductility, is the ability to be drawn into wire.

When a force is applied to a metal, layers of cations slide over one another. The delocalized electrons move with them, so after the shift each cation is still surrounded by the electron cloud and still attracted to it. Because the bonding is non-directional, no particular bond has to break, and the metal changes shape instead of shattering.

Compare this with an ionic lattice. If one layer of ions slides, ions of the same charge end up side by side. They repel, and the crystal splits. That is why ionic solids are brittle and metals are not.

Relating properties to uses

Use	Property that matters	Link to the bonding
Copper wiring in buildings	High electrical conductivity; ductility	Mobile delocalized electrons; layers slide so copper can be drawn into wire
Aluminium cooking foil	Malleability; thermal conductivity; low density	Non-directional bonding lets it be rolled thin; mobile electrons carry heat
Copper base on a saucepan	High thermal conductivity	Delocalized electrons transfer energy across the base quickly
Thin gold layer on electrical contacts	Conducts electricity; does not corrode	Delocalized electrons carry charge; gold reacts very little with air

When you link a property to a use, give three things: the use, the property, and the feature of metallic bonding that explains it. Density and resistance to corrosion often matter as well, even though they are not explained by the bond model alone.

Strength of the metallic bond (2.3.2)

The strength of the attraction depends on two factors.

• Charge of the cation. A cation with a higher charge attracts the electron cloud more strongly. A metal that forms 2+ ions also releases two electrons per atom, so the cloud has more negative charge in the same space. This is a higher electron density, the term the guide uses in its guidance.
• Radius of the metal ion. A smaller cation lets the delocalized electrons come closer to its nucleus, so the attraction is stronger. A larger cation gives weaker bonding.

To melt a metal, you must loosen the lattice so the cations can move past each other. A stronger metallic bond needs more energy, so it gives a higher melting point.

Trend down a group

Going down group 1, the cation charge stays at 1+, but the ion radius increases because each element has an extra occupied shell. The attraction between cations and delocalized electrons weakens, so melting points fall.

Metal	Melting point / K
K	337
Rb	312
Cs	302

Caesium melts just above room temperature. Its large ions hold the electron cloud weakly.

Trend across a period

Across a period, from group 1 to group 2, the cation charge rises from 1+ to 2+, the number of delocalized electrons per atom doubles, and the ion becomes smaller because the nuclear charge is greater. All three changes strengthen the metallic bond. Potassium melts at 337 K, while calcium, next to it in period 4, melts at 1115 K.

Worked example 2: comparing electron density

Use these values to compare the number of delocalized electrons in 1.00 cm³ of potassium and of calcium. Take the Avogadro constant as 6.02 × 10²³ mol⁻¹.

Metal	Density / g cm⁻³	Molar mass / g mol⁻¹
K	0.89	39.10
Ca	1.55	40.08

Mass of 1.00 cm³: K = 0.89 g        Ca = 1.55 g

Amount: n = m / M
  K:  0.89 / 39.10 = 0.0228 mol
  Ca: 1.55 / 40.08 = 0.0387 mol

Atoms: n × 6.02 × 10²³
  K:  1.37 × 10²² atoms
  Ca: 2.33 × 10²² atoms

Delocalized electrons: atoms × electrons released per atom
  K:  1.37 × 10²² × 1 = 1.37 × 10²²
  Ca: 2.33 × 10²² × 2 = 4.66 × 10²²

Ratio Ca : K = 4.66 / 1.37 ≈ 3.4

Calcium has more than three times the electron density of potassium, and each of its cations carries double the charge. This is the simple picture behind its much higher melting point.

A note on p-block metals

The same ideas apply to metals in the p-block, such as aluminium (Al³⁺) and tin. Treat the trend as a guide, not a law. Real metals pack their atoms in different arrangements, and some p-block metals do not fit the pattern well: gallium, in group 13, melts at only 303 K. If a question gives you data that do not fit, describe what the data show and say the simple model does not fully account for it.

Transition elements (2.3.3, HL only)

Transition elements have electrons in a partly filled d sublevel as well as in the outer s sublevel. In the metallic lattice, both the 4s and the 3d electrons can become delocalized. A period 4 transition metal can therefore release more electrons per atom than potassium or calcium.

This has two consequences.

• High melting points. With more delocalized electrons, the electron density is higher and each cation left behind carries a larger positive charge. The cations are also small, because the nuclear charge is high. The metallic bond is strong, so a lot of energy is needed to disrupt the lattice.
• Electrical conductivity. The delocalized d-electrons add to the mobile charge carriers, so transition elements are good conductors of electricity.

Worked example 3: period 4 melting points (HL only)

Use these melting points to explain the trend from potassium to vanadium.

Element	K	Ca	Sc	Ti	V
Melting point / K	337	1115	1814	1941	2183
Valence configuration	4s¹	4s²	3d¹ 4s²	3d² 4s²	3d³ 4s²

K to Ca: 4s electrons only; delocalized electrons per atom rise from 1 to 2
         Cation charge rises, ion radius falls -> big rise in melting point (778 K)

Ca to V: 3d electrons can now join the 4s electrons in the delocalized cloud
         More delocalized electrons per atom -> higher electron density
         Stronger attraction to the cations -> melting points keep rising

Answer: the number of electrons that can be delocalized rises across
the series, which strengthens the metallic bond and raises the melting point.

The rise is not smooth across the whole d-block. Zinc, with a full 3d sublevel, melts at only 693 K, so its d-electrons take little part in the bonding. The chemical properties of transition elements, such as variable oxidation states and coloured complexes, belong to Reactivity 3.4.

Common errors

• Calling the electrons “free” without saying they are delocalized through the whole lattice, or describing the lattice as made of atoms rather than cations.
• Saying the cations move when a metal conducts. Only the delocalized electrons carry the current.
• Explaining malleability by “bonds breaking and reforming” between particular atoms. The point is that the bonding is non-directional and the electron cloud moves with the layers.
• Mixing up metals and ionic compounds: a solid ionic compound does not conduct, but a solid metal does.
• Using only ion size when comparing metals from different groups. The charge and the number of delocalized electrons change too.
• Writing melting points in degrees Celsius with a “K” label. If data are in kelvin, subtract 273 to get degrees Celsius.
• (HL only) Saying only the 4s electrons are delocalized in a transition metal.

Where to go next

• Condensed recall: Structure 2.3 revision notes
• Exam-style work: Structure 2.3 practice questions
• Related unit: Structure 2.1, The ionic model, for the contrast with brittle ionic lattices
• Electron configurations, which you need for cation charges: Structure 1 topic page
• Course overview: all IB DP chemistry resources
• Tick off what you have revised: printable DP chemistry checklist
• Check yourself in ten minutes: free diagnostics

Official syllabus

This study guide is based on the Chemistry guide for the IB Diploma Programme, published by the International Baccalaureate in February 2023 (first assessment 2025). It follows sub-topic Structure 2.3, The metallic model. Understandings 2.3.1 and 2.3.2 apply to SL and HL; understanding 2.3.3 is additional higher level, so HL only.