What is typically true about the atomic structure of metals?

Metals are characterized by their unique atomic structure, which includes the presence of a "sea of free electrons." This model describes how electrons in metallic bonds are not bound to individual atoms but are instead delocalized and move freely throughout the metal lattice. This electron mobility is responsible for many of the distinctive properties of metals, such as electrical conductivity and malleability.

The ability to conduct electricity, for instance, arises because these free electrons can carry charge through the metal when an electric field is applied. Additionally, the malleability of metals allows them to be shaped without breaking, a property that also stems from the ability of the layers of atoms to slide over one another while maintaining a cohesive bond through the surrounding delocalized electrons.

In contrast, tightly bound electrons, as suggested in one of the other options, are more characteristic of non-metals or insulating materials, where electrons are localized and do not contribute to conductivity. The formation of covalent bonds is generally not a common behavior among metals, which typically exhibit metallic bonding rather than covalent bonding. Lastly, the notion that metals have no mobile charge carriers contradicts their fundamental properties, as it is the presence of these free-moving electrons that enables their conductive qualities.