Metallic bonds are responsible for many properties of metals, such as conductivity. Why is this possible?(1 point)
The reason metallic bonds are responsible for many properties of metals, such as conductivity, can be explained by understanding the nature of metallic bonding.
To understand metallic bonding, we need to consider the structure and properties of metals. Metals consist of closely packed positive metal ions surrounded by a "sea" of delocalized electrons. In metallic bonding, the valence electrons of the metal atoms are not confined to individual atoms, but rather can move freely throughout the crystal lattice.
This "sea" of delocalized electrons is what gives metals their unique properties, including conductivity. When an electric field is applied to a metal, it causes these delocalized electrons to move in a coordinated manner, creating an electric current. Since these electrons are not tied to any specific atom, they can easily flow through the metal lattice, allowing for the efficient transfer of electric charge.
The reason why this is possible is because of the nature of metallic bonding itself. The positive metal ions in the lattice hold onto their loosely held valence electrons fairly weakly, allowing them to move freely and conduct electricity. This is in contrast to other types of chemical bonds, such as covalent or ionic bonds, where the electrons are tightly held and not as mobile.
In summary, metallic bonds allow for the movement of delocalized electrons, which in turn allows metals to exhibit properties like conductivity. The loosely held valence electrons in metallic bonding enable the efficient transfer of electric charge, making metals excellent conductors of electricity.