Assessed in Paper 1A (multiple choice), Paper 1B (data-based) and Paper 2 (short and extended response). IB Chemistry guide (first assessment 2025).
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In preparation: 0 of 3 sections compiled and reviewed. The rest show key terms and common misconceptions from our question bank until they are.
S2.3.1 Metallic bond
Metallic bond
The electrostatic attraction between a lattice of cations and delocalized electrons. Each metal atom releases its valence (outer) electrons into a delocalized system shared by the whole lattice, leaving a cation (Na⁺ in sodium, Mg²⁺ in magnesium, Al³⁺ in aluminium). The bonding is non-directional: each cation is attracted to the delocalized electrons around it in all directions, not to one particular neighbour.
Delocalized electrons (in a metal)
Valence electrons that are not attached to any one atom but are free to move throughout the metal lattice. Their total negative charge equals the total positive charge of the cations, so a piece of metal is electrically neutral: solid sodium has one delocalized electron per Na⁺ ion, solid magnesium two per Mg²⁺ ion.
Lattice of cations
The regular, repeating three-dimensional arrangement of metal cations in a solid metal. The cations vibrate about fixed positions and do not migrate through the solid; the delocalized electrons occupy the spaces around them and hold the lattice together.
Electrical conductivity of metals
Metals conduct electricity in the solid and the liquid state because they contain mobile charge carriers: the delocalized electrons. When a potential difference is applied across a metal, the delocalized electrons drift through the lattice towards the positive terminal. The cations do not move through the lattice, so a metal is not chemically changed by conducting, unlike a molten or aqueous ionic compound, in which ions carry the current.
Thermal conductivity of metals
Metals conduct thermal energy far better than non-metallic solids. Energy is passed on by vibrations of the cations in the lattice, but in a metal the delocalized electrons carry most of it: electrons in the hot region gain kinetic energy and move rapidly through the lattice, transferring energy by collisions to other parts of the metal.
Malleability of metals
The ability to be hammered, pressed or rolled into a new shape without breaking. When a force is applied to a metal, layers of cations slide over one another. Because the metallic bond is non-directional, the delocalized electrons continue to attract the cations in their new positions, so the bonding is not broken and the metal does not shatter. (Ductility, being drawn into wires, has the same explanation.)
Relating properties of metals to their uses
Metals are chosen for a use by the property that the use requires. High electrical conductivity: copper for the wires in electrical cables. High thermal conductivity: aluminium for heat sinks that carry energy away from computer processors, and copper or aluminium for the bases of cooking pans. Malleability: aluminium rolled into foil or pressed into drinks cans, and steel sheet pressed into car body panels. Each of these properties is explained by the lattice of cations and delocalized electrons.
Students often think A metal is made of neutral metal atoms; ions exist only in compounds such as salts. In fact Ions. In the metallic model each metal atom has released its valence electrons into a delocalized system, so the lattice consists of cations surrounded by delocalized electrons.
Students often think The electrons in the 'sea' have been lost by the metal atoms, so a piece of metal is a collection of positive ions with a positive overall charge. In fact No. The delocalized electrons stay within the metal. Their total charge balances the total charge on the cations, so a piece of metal is electrically neutral.
S2.3.2 Strength of metallic bonding
Strength of metallic bonding
The strength of the attraction between the cations and the delocalized electrons. It increases as the charge on the cations increases (each cation contributes more delocalized electrons and attracts them more strongly) and as the radius of the cations decreases (the delocalized electrons are closer to the centre of positive charge). Stronger metallic bonding gives a higher melting point.
Electron density (of delocalized electrons)
The number of delocalized electrons per unit volume of the metal. It is higher when each cation contributes more electrons (higher cation charge) and when the cations are smaller and so more closely packed. A higher electron density, together with a higher cation charge, means a stronger attraction between the cations and the delocalized electrons.
Melting-point trends in s- and p-block metals
Across period 3, melting point increases from sodium (97.8 °C) to magnesium (650 °C) to aluminium (660 °C): the cation charge increases from +1 to +3 and the cation radius decreases, so the electron density and the attraction increase. Down group 1 the melting point decreases, from lithium (180.5 °C) to sodium (97.8 °C) to potassium (63.5 °C): the charge stays +1 but the cation radius increases, so the attraction to the delocalized electrons decreases. The simple model does not account for every irregularity (calcium, 842 °C, melts higher than magnesium, 650 °C, although Ca²⁺ is larger than Mg²⁺), so it is applied to comparisons in which the charge or the radius of the cation changes clearly.
Students often think The heavier the atoms, the more energy is needed to melt the metal, so melting point rises with relative atomic mass. In fact No, not as a rule. Melting point depends on the strength of the metallic bonding, which depends on the charge and radius of the cations, not on their mass. Potassium (A_r 39.10) melts at 63.5 °C, far below lithium (A_r 6.94, 180.5 °C).
Students often think Metals are held together by intermolecular forces such as London forces, so their melting points rise with the number of electrons in each atom, as for molecular substances. In fact No. A metal is held together by metallic bonding, the attraction between the cations and the delocalized electrons throughout the lattice. Its melting point reflects the strength of this attraction, not London forces.
S2.3.3 Transition element HL
Transition element
An element whose atom, or one of whose common ions, has an incomplete d sublevel, for example titanium ([Ar] 3d² 4s²) and iron ([Ar] 3d⁶ 4s²). The first-row transition elements are all metals.
Delocalized d-electrons in transition elements
In transition elements the 3d electrons, as well as the 4s electrons, are delocalized in the metallic lattice. This gives a higher density of delocalized electrons and stronger metallic bonding than in the s-block metals of the same period, which explains the high melting points of transition elements (titanium 1668 °C and iron 1538 °C, against calcium 842 °C and potassium 63.5 °C). The delocalized s and d electrons are also mobile charge carriers, which explains the electrical conductivity of transition elements.
Students often think Only the outer 4s electrons of a transition element are delocalized; the 3d electrons are inner electrons that stay on the cations. In fact Yes. In transition elements the 3d electrons, as well as the 4s electrons, are delocalized in the metallic lattice. This raises the density of delocalized electrons, giving strong metallic bonding and high melting points, and supplies mobile charge carriers for electrical conduction.
Diagnostic a bearings check, not a test
6 questions, one per part of the topic where we can. Answer them, then see which statements you own and which to read.
1 Which statement describes a metallic bond?
Answer and reasoning
The sharing of electron pairs between each metal atom and its nearest neighbours — A student who pictures metallic bonds as localized bonds between neighbouring atoms, like covalent bonds, picks this. In a metal the valence electrons are delocalized over the whole lattice, not shared in pairs between two particular atoms.
The electrostatic attraction between a lattice of cations and delocalized electrons — This is the metallic bond as the IB defines it. Each atom releases its valence electrons into a delocalized system shared by the whole lattice, and the resulting cations are attracted to these electrons in all directions.
The attraction between cations and anions formed by electron transfer between metal atoms — A student who models metallic bonding on ionic bonding picks this. A pure metal contains no anions: the negative charge is carried by delocalized electrons that belong to no particular atom.
The London forces acting between the closely packed atoms of a metal lattice — A student who treats a metal like a molecular substance, held by intermolecular forces, picks this. A metal is a giant structure held by metallic bonding, the attraction between cations and delocalized electrons.
2 The melting points of the group 1 metals decrease down the group: lithium 180.5 °C, sodium 97.8 °C, potassium 63.5 °C. Which statement explains this trend?
Answer and reasoning
Reactivity increases down the group, so the bonds break more easily when the metal is heated. — A student who treats one group 1 trend as the cause of another picks this. Reactivity and melting point are separate properties; both trends follow from the increasing radius, but reactivity does not cause the fall in melting point.
The cation radius increases, so the delocalized electrons are attracted less strongly. — Every group 1 cation has a charge of +1, but the cation radius increases from Li⁺ to Na⁺ to K⁺. The delocalized electrons are further from the centre of positive charge, so the metallic bonding weakens and less energy is needed to melt the metal.
Larger atoms form longer and weaker covalent bonds with each of their neighbouring atoms. — A student who pictures metallic bonds as localized bonds between pairs of atoms picks this. There are no covalent bonds in a metal; the attraction is between the cations and the delocalized electrons in all directions.
The cations and anions in the lattice get larger, so their ionic attraction weakens. — A student who models metallic bonding on ionic bonding picks this. A group 1 metal contains no anions: the lattice is of cations only, held by the delocalized electrons, and it is the increasing radius of these cations that weakens the attraction.
3 Iron ([Ar] 3d⁶ 4s²) is a transition element. Which statement explains the electrical conductivity of solid iron? HL
Answer and reasoning
Its 4s electrons are delocalized, but its 3d electrons are core electrons fixed on the cations. — A student who treats the 3d electrons as inner electrons picks this. In transition elements the 3d electrons are delocalized along with the 4s electrons; the guide's model of transition elements rests on these delocalized d-electrons.
Its cations move through the lattice towards the negative terminal, carrying the charge. — A student who carries over conduction by moving ions from electrolytes picks this. The iron cations stay in the lattice; solid iron conducts because its delocalized electrons are mobile.
Electrons from the power supply pass through it while its own electrons stay in place. — A student who thinks the cell supplies the moving charge picks this. The charge carriers are iron's own delocalized electrons; the power supply provides the potential difference that makes them drift.
Its 3d and 4s electrons are delocalized and drift when a potential difference is applied. — In a transition element the 3d electrons, as well as the 4s electrons, are delocalized in the metallic lattice. These mobile electrons drift through the lattice of cations when a potential difference is applied, carrying the current.
4 Which statement about a piece of solid sodium is correct?
Answer and reasoning
It carries a positive charge, because its atoms have lost their outer electrons. — A student who reads the 'sea of electrons' as electrons lost by the metal picks this. The delocalized electrons remain in the metal and balance the charge of the Na⁺ ions, so solid sodium is electrically neutral.
It contains no ions; its lattice is made only of neutral sodium atoms. — A student who thinks ions exist only in compounds picks this. In the metallic model each sodium atom has released its valence electron, so the lattice consists of Na⁺ cations surrounded by delocalized electrons.
It is neutral, with one delocalized electron for each Na⁺ ion in its lattice. — Each sodium atom releases its one valence electron into the delocalized system, leaving Na⁺. The electrons stay in the metal, so their total charge balances the total charge of the cations and the piece of sodium is neutral.
Its atoms are joined in pairs, each pair held together by a shared electron pair. — A student who pictures metallic bonds as covalent-style bonds between neighbouring atoms picks this. Solid sodium has no pairs or molecules: its valence electrons are delocalized over the whole lattice of Na⁺ ions.
5 The melting points of three period 3 metals are sodium 97.8 °C, magnesium 650 °C and aluminium 660 °C. Which statement explains why the melting point increases from sodium to aluminium?
Answer and reasoning
The relative atomic mass rises, so the heavier atoms need more energy to separate. — A student who links melting point to mass picks this. Mass does not set the strength of metallic bonding: down group 1 the mass rises while the melting point falls (lithium 180.5 °C, potassium 63.5 °C).
Each atom forms more covalent bonds to its neighbours: one for Na, two for Mg and three for Al. — A student who pictures metallic bonds as localized, covalent-style bonds counted from the group number picks this. Metallic bonding is non-directional; the group number matters because it sets the cation charge and the number of delocalized electrons per cation.
The cations have a higher charge and a smaller radius, so the metallic bonding is stronger. — From Na⁺ to Mg²⁺ to Al³⁺ each cation contributes more delocalized electrons and is smaller, so the electron density and the attraction between the cations and the delocalized electrons increase. Stronger metallic bonding needs more energy to overcome in melting.
The London forces between the atoms become stronger as the number of electrons in each atom rises. — A student who treats melting a metal like melting a molecular substance picks this. Metals are held by metallic bonding, the attraction between cations and delocalized electrons, not by London forces between atoms.
6 Calcium ([Ar] 4s²) melts at 842 °C. Titanium ([Ar] 3d² 4s²), the next metal but one in period 4, melts at 1668 °C. Which statement explains the higher melting point of titanium? HL
Answer and reasoning
Its atoms are heavier than calcium atoms, so more energy is needed to separate them. — A student who links melting point to mass picks this. Mass does not set the strength of metallic bonding; potassium is heavier than lithium yet melts far lower. Titanium melts high because its delocalized d-electrons strengthen the bonding.
Its extra 3d electrons give stronger London forces between its atoms than calcium's have. — A student who treats a metal's melting point as a matter of intermolecular forces picks this. Titanium is a giant metallic structure; its extra 3d electrons raise the melting point by joining the delocalized system, not through London forces.
Its 3d as well as its 4s electrons are delocalized, so the electron density is higher. — In a transition element the 3d electrons join the 4s electrons in the delocalized system. Titanium therefore has more delocalized electrons per cation than calcium, a higher electron density and stronger attraction between cations and delocalized electrons, so more energy is needed to melt it.
Its 3d electrons are not delocalized, so its smaller cation is what makes the bonding stronger. — A student who treats the 3d electrons as core electrons picks this. In transition elements the 3d electrons are delocalized too; this higher electron density is the main reason transition elements melt far higher than the s-block metals of the same period.
Read the ones marked not yet in Learn, then Verify.
Verify confirm before you go
4 more questions. Every wrong answer here is a real misconception, and you see why it is wrong straight away.
1 Copper is used for the wires in household electrical cables because it conducts electricity very well. What carries the electric current through a copper wire?
Answer and reasoning
The copper's delocalized electrons, drifting through the lattice — The delocalized electrons are mobile charge carriers already present throughout the copper. When a potential difference is applied they drift through the lattice of cations, which stay in place. This high electrical conductivity is why copper is used for wiring.
Copper cations, moving through the lattice towards the negative terminal — A student who carries over conduction by moving ions from molten or aqueous ionic compounds picks this. In solid copper the cations stay in the lattice; a copper wire does not change in composition when it conducts.
Electrons from the power supply; the copper's own electrons stay in place — A student who thinks the cell supplies the moving charge picks this. The charge carriers are the copper's own delocalized electrons; the power supply provides the potential difference that makes them drift.
Vibrating copper ions, each passing the energy on to the next ion along — A student who pictures a current as energy handed on from particle to particle picks this. An electric current is a flow of charge; in copper the charge carriers are the delocalized electrons, and the cations stay in place.
2 Aluminium is used for heat sinks, which carry thermal energy away from computer processors. Which statement explains why aluminium is suitable?
Answer and reasoning
Its high melting point lets it absorb all the heat without melting. — A student who thinks heat-related uses depend on withstanding heat picks this. Aluminium melts at 660 °C, far above a processor's temperature; what the use needs is rapid transfer of energy, which is high thermal conductivity.
Its vibrating ions pass the energy on; its electrons play no part. — A student who uses only the vibrating-particle model of conduction picks this. Ion vibrations do contribute, but in a metal the delocalized electrons carry most of the energy, which is why metals conduct heat far better than non-metallic solids.
Metals stay colder than their surroundings, so it keeps drawing heat away. — A student who thinks metals are naturally colder than other materials picks this. Aluminium left in a room reaches room temperature; it removes energy from the hotter processor because it conducts that energy away quickly.
Its delocalized electrons carry kinetic energy rapidly through the whole metal. — A heat sink must conduct energy away quickly. In aluminium the delocalized electrons gain kinetic energy near the hot processor and move rapidly through the lattice, transferring the energy to the cooler fins. This high thermal conductivity is the reason for the choice.
3 Aluminium is rolled into kitchen foil less than 0.02 mm thick without cracking. Which statement explains this property of aluminium?
Answer and reasoning
Its layers of cations slide, staying attracted to the delocalized electrons around them. — This is malleability. Because the metallic bond is non-directional, the delocalized electrons keep attracting the cations after the layers have slid into new positions, so the bonding is not broken and the metal changes shape without cracking.
Its metallic bonds are weak, so the rollers can easily push its atoms apart. — A student who reads 'easily shaped' as 'weakly bonded' picks this. Aluminium's metallic bonding is strong (it melts at 660 °C); it can be rolled thin because the non-directional bonding is kept as layers of cations slide.
Its bonds between neighbouring atoms bend like hinges instead of breaking. — A student who pictures metallic bonds as localized bonds between pairs of atoms picks this. There are no bonds between particular atoms to bend: the attraction is between the cations and the delocalized electrons in all directions.
Its individual atoms are soft, so each one is flattened as it passes between the rollers. — A student who gives individual particles the properties of the bulk material picks this. The cations keep their size and shape; the metal changes shape because layers of cations move relative to one another.
4 Data for three metals are shown as (A_r; cation; cation radius; melting point): lithium (6.94; Li⁺; 76 pm; 180.5 °C), sodium (22.99; Na⁺; 102 pm; 97.8 °C), calcium (40.08; Ca²⁺; 100 pm; 842 °C). Which statement explains these data using the metallic model?
Answer and reasoning
Melting point rises with A_r: calcium has the heaviest atoms and the highest melting point. — A student who links melting point to mass picks this. The data rule it out: lithium (A_r 6.94) is lighter than sodium (22.99) but melts higher.
Calcium is less reactive than sodium, so its bonds are harder to break on heating. — A student who treats reactivity as the cause of melting point picks this. Reactivity and melting point are separate properties; calcium melts high because Ca²⁺ has twice the charge of Na⁺ at almost the same radius, which gives stronger metallic bonding.
Cation radius decides bond strength: the smallest ion, Li⁺, should give the strongest bonding. — A student who learned "smaller ion, stronger bond" from group 1 as the whole rule picks this. Li⁺ (76 pm) is the smallest ion, yet calcium, whose Ca²⁺ ion (100 pm) is larger, melts about 660 °C higher. The doubled charge of Ca²⁺ outweighs its larger size.
Na⁺ and Ca²⁺ are almost the same size, so the higher charge of Ca²⁺ gives calcium much stronger bonding. — Comparing sodium with calcium isolates the effect of charge, because the radii (102 pm and 100 pm) are almost equal. Each Ca²⁺ contributes two delocalized electrons and attracts them more strongly, so calcium's bonding is much stronger. Comparing lithium with sodium isolates radius: the larger Na⁺ gives the lower melting point.
That was your twenty minutes. Real practice on S2.3 is past-paper questions marked against the mark scheme.
Paper 1A tests it as multiple choice; Paper 1B through data you have not seen before; Paper 2 with short answers and, at HL, extended responses. Look for the command words — outline, explain, compare, evaluate — and give exactly what each asks for.
Compiled from the IB Chemistry guide (first assessment 2025) and our question bank · Specialist review in progress. How these pages are made · Free, no account ·
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