Metallic bonds are the Coulombic attractions between positively charged metal cations and a "sea" of delocalized valence electrons, which explains why metals conduct electricity, are malleable, and tend to have high melting points (AP Chem Unit 3, Topic 3.2).
A metallic bond is the attraction between metal cations and the valence electrons those atoms have given up to the whole structure. Instead of belonging to one atom or one bond, the valence electrons are delocalized, meaning they're free to move throughout the entire solid. Picture metal cations sitting in a fixed lattice while a sea of mobile electrons flows around them and glues everything together. That's the electron-sea model, and it's the picture AP Chem wants you to draw and describe.
This structure is the whole story behind metallic properties. Mobile electrons carry charge, so metals conduct electricity as solids (no melting or dissolving required). Cations can slide past each other without breaking the bond, because the electron sea just flows around the new arrangement, so metals are malleable and ductile instead of shattering. The strength of the Coulombic attraction between cations and the electron sea also gives many metals high melting points. On the exam, every metallic property you cite should trace back to this particulate-level structure.
Metallic bonding lives in Unit 3 (Properties of Substances and Mixtures), mainly Topic 3.2 (Properties of Solids), with roots in Topic 3.1's Coulombic logic. It directly supports AP Chem 3.2.A, which asks you to explain the relationship among a substance's macroscopic properties, its particulate-level structure, and the interactions between particles. Metallic solids are one of the four solid types you have to classify (alongside ionic, covalent network, and molecular), and the exam loves giving you a mystery solid's properties and asking which type it is. Knowing that "conducts as a solid" is the metallic fingerprint, while "conducts only when melted or dissolved" screams ionic, is one of the highest-yield distinctions in the unit. Metallic bonding also connects forward to alloys (interstitial and substitutional), where adding a second element changes malleability and hardness.
Keep studying AP Chemistry Unit 3
Delocalized Electrons (Unit 3)
Delocalized electrons ARE the metallic bond. Every metallic property you'll ever explain on the exam comes back to electrons that aren't stuck to one atom. If your answer doesn't mention mobile or delocalized electrons, it's probably incomplete.
Conductivity (Unit 3)
Metals conduct electricity in the solid state because the electron sea moves freely when a voltage is applied. Compare this to ionic solids, which only conduct when melted or dissolved, because there the charge carriers are ions that need to be freed from the lattice first.
Malleability (Unit 3)
When you hammer a metal, layers of cations slide past each other and the electron sea simply re-glues them in their new positions. An ionic solid shatters under the same stress because shifting the lattice puts like charges next to each other. Alloys change this, since interstitial atoms like carbon in copper jam the layers and reduce malleability.
Coulombic Interactions (Units 1 & 3)
Metallic bonding is just Coulomb's law at work. Positive cations attract negative delocalized electrons, the same attraction logic behind ionization energy in Unit 1 and intermolecular forces in Topic 3.1. AP Chem rewards explanations built on charge attraction, not memorized labels.
Covalent Network Solids (Unit 3)
Both metallic and covalent network solids can have very high melting points, but for different reasons. Network solids like diamond hold every electron in localized covalent bonds, so they're hard, brittle, and (usually) nonconductive. Metals stay strong while still bending and conducting because their electrons roam.
Metallic bonding shows up most often in classify-the-solid multiple choice questions. A typical stem describes an unknown solid's melting point, conductivity (as a solid, melted, and dissolved), and hardness, then asks you to identify the bonding type or explain a property. For example, a solid that conducts only when melted or dissolved is ionic, not metallic, and you need to articulate why. Alloy questions are also fair game, like explaining why a copper-carbon interstitial alloy is harder and less malleable than pure copper (the carbon atoms sit between copper layers and block sliding). No released FRQ has centered on metallic bonds verbatim, but free-response questions regularly ask you to connect particulate structure to macroscopic properties under LO 3.2.A, and a particulate drawing of cations in an electron sea is exactly the kind of representation the CED expects you to interpret or sketch.
Both are strong Coulombic attractions in a lattice, but the players differ. Ionic bonds attract cations to anions (two oppositely charged ions), while metallic bonds attract cations to delocalized electrons. The conductivity test separates them instantly. Metals conduct as solids because electrons are already mobile; ionic compounds only conduct when melted or dissolved, since the ions have to be freed first. Mechanically, metals bend and ionic solids shatter, because sliding an ionic lattice forces like charges together and the lattice repels itself apart.
A metallic bond is the Coulombic attraction between metal cations and a sea of delocalized valence electrons, not an attraction between two specific atoms.
Metals conduct electricity as solids because their delocalized electrons are free to move; ionic solids only conduct when melted or dissolved.
Metals are malleable and ductile because cation layers can slide past each other while the electron sea keeps holding the structure together.
Adding atoms to a pure metal makes an alloy, and interstitial atoms (like carbon in copper) wedge between layers, making the alloy harder but less malleable.
On the exam, always explain metallic properties by linking them to particulate-level structure, which is exactly what learning objective 3.2.A asks for.
It's the attraction between positively charged metal cations and delocalized valence electrons that move freely throughout the solid. This electron-sea model explains why metals conduct electricity, bend without breaking, and often have high melting points.
No. Metallic bonding is an interparticle force within a single continuous solid, not a force between separate molecules. Metals don't have discrete molecules at all, which is why metallic solids behave so differently from molecular solids held together by dispersion or dipole-dipole forces.
An ionic bond attracts cations to anions, while a metallic bond attracts cations to free-moving electrons. The quickest tell on an exam question is conductivity. Metals conduct as solids, while ionic compounds conduct only when melted or dissolved in water.
Conduction requires mobile charged particles. In metals, delocalized electrons are already free to move through the solid. In ionic solids, the ions are locked in a rigid lattice and can only carry charge after melting or dissolving frees them.
When a metal is deformed, layers of cations slide past each other and the electron sea immediately re-bonds them in their new positions. Ionic solids shatter under the same stress because shifting the lattice lines up like charges, which repel and crack the crystal.