S2.3
The metallic model
2 SL · +1 HL (SL 2.3.1-2.3.2) (HL 2.3.3)
Unverified
← View on the mapParent topic: Models of bonding and structure
Guiding question
Guiding questionWhat determines the metallic nature and properties of an element
The one big idea
A metallic bond is the electrostatic attraction between a lattice of cations and a sea of delocalised electrons. It explains conductivity, malleability and melting points — and at HL, the special properties of transition elements.
What this rests on
Our editorial judgement.The dependencies in this section are this site’s own assessment of what a sub-topic rests on. They have notbeen verified by a chemistry teacher and are not part of the official IB guide. By contrast, section 8 (Linking questions) reproduces the IB’s own wording verbatim.
- EditorialS3.1 The periodic table: Classification of elementsMetallic character and cation charge/radius trends come from periodic position.
What rests on this
- EditorialS2.4 From models to materials
- EditorialR3.2 Electron transfer reactions
Core concepts that must be mastered
- Metallic bonding: cations in a sea of electrons(2.3.1)A metallic bond is electrostatic attraction between a lattice of positive cations and delocalised electrons that are free to move throughout the structure. Like ionic bonding (S2.1), metallic bonding is non-directional — the electrostatic attraction acts in all directions with no privileged partner. The difference is structural: ionic bonding operates between discrete ions fixed at lattice sites, whereas metallic bonding uses a delocalised electron sea shared across the entire lattice. This is in contrast to covalent bonding (S2.2), which is directional between specific atoms. The delocalised, mobile nature of the electron sea explains the characteristic properties of metals: electrical conductivity (delocalised electrons can carry charge in any direction — metals conduct in both solid and molten states, unlike ionic compounds which conduct only when molten or dissolved), thermal conductivity (the mobile electrons transfer kinetic energy efficiently — this is why metals feel cold to touch, as they conduct heat away from your hand), malleability (the non-directional bonding means layers of cations can slide past one another without breaking specific bonds — the electron sea simply adjusts, so the metal deforms rather than shatters), and ductility (the same property allows metals to be drawn into wires). In contrast, ionic solids are brittle: if the lattice layers are displaced, ions of like charge are brought next to each other and repel, causing the crystal to shatter. Relate these properties to uses: copper for wiring (conductivity), aluminium for cookware (thermal conductivity), gold for jewellery (malleability and ductility — gold is the most malleable metal).
- Metallic bond strength(2.3.2)The strength of a metallic bond depends on the ionic charge of the cation and the radius of the metal ion. Higher charge means more delocalised electrons per cation and stronger electrostatic attraction; smaller radius means the cations are closer together and the attraction is stronger. This gives trends in melting points of s- and p-block metals: group 1 metals (Na, K, Rb) have low melting points because each atom contributes only one delocalised electron — sodium melts at 98 °C, potassium at 64 °C; group 2 metals (Mg, Ca) have higher melting points because each contributes two — magnesium melts at 650 °C. Across a period, melting points of metals generally increase as charge rises and radius falls — but the trend is complicated by the transition elements. The treatment here is simple: more charge and smaller radius mean stronger bonding and higher melting point. The melting point trend down group 1 (Li → Na → K → Rb → Cs) actually decreases because the ionic radius increases faster than any other factor — the larger cations are farther apart and the electron sea is more diffuse.
- HL: Transition elements and delocalised d-electrons(2.3.3)At HL, transition elements have delocalised d-electrons in addition to their s-electrons. These d-electrons contribute to the metallic bond and explain why transition elements have high melting points — often much higher than s-block metals — and excellent electrical conductivity. The extra bonding from the d-electrons adds significantly to the lattice strength. The chemical properties of transition elements (variable oxidation states, coloured complexes, catalytic activity) are covered in Structure 3.1 and Reactivity 3.4; here, the focus is on how the d-electrons strengthen the metallic bond. The trend in melting points across the d-block is less smooth than across the s-block because the number of d-electrons varies and the contribution to bonding is not as simple as "one electron per atom." Elements near the middle of the d-block (Cr, Mo, W) have the highest melting points because they contribute the most d-electrons to the metallic bond without filling the d-sublevel so much that the electrons become localised.
- How this sub-topic connectsThe metallic model completes the three bonding models (ionic S2.1, covalent S2.2, metallic S2.3) that Structure 2.4 unifies into a continuum. The non-directional nature of metallic bonding explains why metals can form alloys (S2.4): different atoms can sit in the lattice without breaking specific bonds, and the electron sea adjusts to accommodate them. At HL, the d-electron explanation connects to transition element chemistry (S3.1: variable oxidation states, coloured complexes, catalytic activity) and to catalysis (R3.4: transition elements catalyse by changing oxidation state, providing an alternative pathway with lower activation energy). The periodic trend in metallic character (S3.1: metallic character increases down a group and decreases across a period) builds directly on the bond-strength relationship established here — the balance between charge, radius and number of delocalised electrons determines the properties.
Quantitative non-negotiables
No quantitative non-negotiables specified.
Common failure modes
No misconceptions recorded for this sub-topic.
What "HL standard" actually looks like
HLHL adds delocalised d-electrons in transition elements, explaining their high melting points and electrical conductivity. Chemical properties of transition elements are covered in Reactivity 3.4.
Linking questions
Official IB.The questions in this section are the IB’s own linking questions, reproduced verbatim from the guide. They are not this site’s editorial judgement — see section 3 for that distinction.
- Official IB
Why are alloys generally considered to be mixtures, even though they often contain metallic bonding?
- Official IB
What experimental data demonstrate the physical properties of metals, and trends in these properties, in the periodic table?
- Official IB
What trends in reactivity of metals can be predicted from the periodic table?
- Official IB
What are the features of metallic bonding that make it possible for metals to form alloys?
- Official IBwording unverified against the guide
Why is the trend in melting points of metals across a period less evident across the d-block?
- Official IB
What are the arguments for and against including scandium as a transition element?
- Official IB
The surface oxidation of metals is often known as corrosion. What are some of the consequences of this process?
Bridge: GCSE → IB HL
Derived. These are the bridge items tagged to this sub-topic— places where the GCSE model gets redefined, genuinely new territory, or carry-over strengths. The tagging is this site’s editorial judgement.