IB DP Chemistry Structure 2.3: The metallic model – Revision Notes

Use these notes for fast recall of Structure 2.3, The metallic model, as it appears in the IB’s DP Chemistry guide (February 2023 edition, examined from 2025). They cover understandings 2.3.1 and 2.3.2, which are SL and HL, and 2.3.3, which is additional higher level and so HL only. If an idea here feels thin, go back to the metallic model study guide, which explains each point at length.

Full-length questions with marked answers are in the Structure 2.3 practice set. Other units are listed on the course page for DP chemistry, you can print a revision checklist for DP chemistry, and our diagnostics reveal any gaps.

Key definitions

Term	Meaning
Metallic bond	Electrostatic attraction holding a lattice of cations to the delocalized electrons around them
Delocalized electrons	Valence electrons not tied to one atom; spread through the whole lattice and free to move
Cation lattice	Regular, closely packed arrangement of the positive ions left after atoms release their valence electrons
Non-directional bonding	Attraction acts equally in all directions; no bond links one specific pair of particles
Malleable	Can be hammered or rolled into a new shape without breaking
Ductile	Can be drawn out into wire
Electron density	How much delocalized negative charge is packed into a given volume of metal

Structure 2.3.1: bonding and properties

The model in one line: cations fixed in a lattice, held together by a mobile cloud of delocalized electrons.

Property	Explanation in terms of the model
Conducts electricity (solid and liquid)	Under a potential difference the delocalized electrons drift towards the positive terminal; the cations stay put
Conducts heat	Mobile electrons gain kinetic energy at the hot end and carry it quickly to cooler regions; cation vibrations also pass energy on
Malleable and ductile	Layers of cations slide; the electron cloud moves with them, so attraction is kept and the metal reshapes instead of breaking
Usually high melting point	Strong attraction between cations and electrons throughout the lattice needs much energy to disrupt

Method in steps: “relate a property to a use”

1. Name the use (for example, copper water pipes).
2. Name the property it needs (malleable enough to be bent; conducts heat well in a heating system).
3. Explain that property from the bonding (layers slide with the electron cloud; mobile electrons carry energy).
4. Add any other factor the question hints at, such as low density or resistance to corrosion. These are real reasons for a choice, but the bond model alone does not explain them.

Structure 2.3.2: strength of the metallic bond

Bond strength rises with:

• higher cation charge (more delocalized electrons per atom, so greater electron density and stronger attraction)
• smaller ion radius (electrons closer to the nucleus of each cation)

Stronger metallic bond → more energy to disrupt the lattice → higher melting point.

Change	Charge	Radius	Melting point
Down a group (for example group 1, K → Rb → Cs)	same	increases	decreases
Across a period, group 1 → group 2 (for example Rb → Sr)	increases	decreases	increases
Across period 3 into the p-block (Na → Mg → Al)	1+ → 2+ → 3+	decreases	increases (371 K → 923 K → 933 K)

Method in steps: “compare the melting points of two metals”

1. Write the charge of each cation (from the group number).
2. State how many electrons each atom adds to the delocalized cloud.
3. Compare the ion sizes (more shells means larger; more protons in the same shell means smaller).
4. Say which has the stronger attraction between cations and delocalized electrons.
5. Conclude that more energy is needed to overcome it, so that metal has the higher melting point.

Worked reminder

Which melts higher, strontium or rubidium? Both are in period 5.

Rb → Rb⁺ + 1 delocalized electron per atom
Sr → Sr²⁺ + 2 delocalized electrons per atom
Sr²⁺ is smaller (more protons pulling on the same shells)
Higher charge + higher electron density + smaller ion -> stronger metallic bond
Sr has the higher melting point (about 1050 K, against about 312 K for Rb)

For p-block metals such as aluminium, apply the same two factors: Al³⁺ has the highest charge and smallest radius in period 3, so aluminium melts highest, though only just above magnesium. The simple model works well for s-block metals. Some p-block metals break the pattern (gallium melts at 303 K), so trust data over prediction when the two disagree.

Structure 2.3.3: transition elements (HL only)

• In transition elements, 3d electrons as well as 4s electrons can be delocalized.
• More delocalized electrons per atom → greater electron density and higher cation charge, with small ions → very strong metallic bonding → high melting points.
• The extra delocalized electrons are mobile charge carriers → good electrical conductivity.
• Across period 4 the melting point rises sharply from K (337 K) to Ca (1115 K), and keeps rising into the d-block, reaching 2183 K at V.
• The trend is less regular across the d-block than the simple model suggests. Zinc, with a full 3d sublevel, melts at only 693 K.
• Chemical properties of transition elements are in Reactivity 3.4, not here.

Must-know distinctions

	Metal	Ionic compound
Particles in the lattice	Cations only, plus delocalized electrons	Cations and anions
Charge carriers when it conducts	Electrons	Ions
Conducts as a solid?	Yes	No
Effect of a strong force	Layers slide; metal bends	Like charges meet; crystal shatters

• Delocalized (spread over the whole lattice) is not the same as lone or bonding pair (localized between atoms).
• A metal sample is written with its element symbol. There is no molecular or empirical ratio of cations to electrons to show.
• Group number gives the cation charge for s-block metals. It does not give a fixed charge for transition elements.
• Electrical conductivity is about moving charge; thermal conductivity is about moving energy. The same mobile electrons explain both, but the wording of each answer is different.
• Melting a metal does not break the metallic bond completely. The cations can move past each other, but they are still held by the delocalized electrons, which is why liquid metals still conduct.

Quick self-test

1. Define a metallic bond.
2. Which particles carry the current when a copper wire conducts?
3. Why does liquid mercury conduct electricity?
4. Explain in one sentence why a metal does not shatter when it is hammered.
5. Give one reason why aluminium is used for cooking foil, and link it to the bonding.
6. How many electrons does each barium atom add to the delocalized cloud, and what is the charge on its cation?
7. Predict which has the higher melting point, sodium or potassium, and give two reasons.
8. Explain why calcium has a higher melting point than potassium.
9. Calculate the number of delocalized electrons released by 0.500 mol of aluminium atoms. Use 6.02 × 10²³ mol⁻¹ for the Avogadro constant.
10. Convert caesium’s melting point of 302 K to degrees Celsius.
11. (HL only) Why do transition elements have higher melting points than s-block metals in the same period?
12. (HL only) Which electrons in solid titanium are delocalized?

Answers

1. The electrostatic attraction that holds a lattice of cations to the delocalized electrons around them.
2. The delocalized electrons. The Cu cations stay in their lattice sites.
3. Melting does not remove the delocalized electrons, so they can still move through the liquid and carry charge.
4. The layers of cations slide over each other and the delocalized electrons move with them, so the non-directional attraction is kept.
5. It is malleable, so it can be rolled very thin: the layers of Al³⁺ cations slide while staying attracted to the electron cloud. (Thermal conductivity from mobile electrons is also accepted.)
6. Two electrons; Ba²⁺.
7. Sodium. Both form 1+ ions, but Na⁺ is smaller (one fewer shell), so the delocalized electrons are attracted more strongly and more energy is needed to melt it.
8. Calcium forms Ca²⁺ and releases two electrons per atom, so the cation charge and electron density are higher. Ca²⁺ is also smaller than K⁺. The attraction is stronger, so more energy is needed to melt calcium.
9. 0.500 × 3 × 6.02 × 10²³ = 9.03 × 10²³ electrons.
10. 302 − 273 = 29 °C.
11. Their 3d electrons, as well as their 4s electrons, can be delocalized, giving a higher electron density and stronger attraction to small, highly charged cations.
12. Both 4s electrons and the 3d electrons (valence configuration 3d² 4s²).

Where marks are usually lost

• Writing “sea of free electrons” but never saying the electrons are attracted to the cations. The bond is the attraction, not the electrons.
• Describing a metal lattice as “atoms” or as “positive and negative ions”. It is cations with delocalized electrons.
• Saying ions move when a solid metal conducts.
• For malleability, stopping at “layers slide”. Add that the delocalized electrons move with them and the attraction is kept.
• For thermal conductivity, saying only “electrons move”. Say they gain and carry kinetic energy.
• Comparing melting points with one factor only. Give the charge (or number of delocalized electrons) and the ion radius.
• Saying “more energy to break bonds” without naming the bond as the attraction between cations and delocalized electrons.
• Linking a use to a property without explaining the property from the bonding.
• (HL only) Forgetting to say the d-electrons are delocalized when explaining transition metal melting points.

Official syllabus

Written to the International Baccalaureate’s DP Chemistry guide (February 2023; first assessment 2025). The notes track sub-topic Structure 2.3, The metallic model: understandings 2.3.1 and 2.3.2 (SL and HL) and 2.3.3 (additional higher level, HL only).