Metallic Bonding – Chapter 9

Metallic bonding refers to the strong electrostatic forces of attraction between positively charged metal ions and delocalized electrons.

• Metals form a giant structure where positive ions are surrounded by a sea of delocalized electrons.
• This bonding is responsible for many physical properties of metals, such as conductivity and malleability.
• When sodium atoms bond to form sodium metal, each outer electron becomes free to move, creating a metallic lattice.

Metallic bonding significantly affects the electrical conductivity of metals.

• Delocalized electrons in metallic structures are free to move, allowing metals to conduct electricity efficiently.
• This property is crucial for applications in electrical wiring and electronic components.
• In contrast to insulators, metals can easily transfer electric charge due to their metallic bonds.

The physical properties of metals are largely attributed to metallic bonding.

• Metals are typically hard and have high melting points due to the strong metallic bonds that require substantial energy to break.
• They exhibit malleability and ductility, allowing them to be shaped without breaking.
• This is because the layers of positive ions can slide over each other while maintaining their attraction to delocalized electrons.

Metallic bonds contribute to the malleability of metals by allowing layers of atoms to slide over one another.

• When a force is applied, the positive ions can shift positions without breaking the metallic bond.
• This flexibility is a result of the delocalized electrons that maintain the attraction between the ions.
• Consequently, metals can be hammered or bent into various shapes without fracturing.

The structure of metals in metallic bonding consists of a giant lattice of positive ions surrounded by delocalized electrons.

• There are no individual molecules; all positive ions attract delocalized electrons throughout the structure.
• The ion formed by a metal depends on the number of electrons in its outer shell.
• For example, group 1 elements form +1 ions, while group 2 elements form +2 ions.

Sodium atoms bond in metallic bonding by losing their outer electrons to form positive ions.

• When sodium atoms bond to create sodium metal, each atom loses one electron, resulting in Na+ ions.
• This process creates a lattice structure where these ions are surrounded by delocalized electrons.
• The movement of these electrons contributes to the metallic properties of sodium metal.

Metallic bonding and ionic bonding are two distinct types of chemical bonding.

Examples of metals that exhibit metallic bonding include sodium, magnesium, and aluminum.

• Sodium (Na) forms a +1 ion and has a simple metallic structure.
• Magnesium (Mg) forms a +2 ion and has a more complex lattice.
• Aluminum (Al) is known for its lightweight and strong metallic bonds.

Metallic bonding explains the high melting points of metals due to the strength of the bonds between ions and delocalized electrons.

• Breaking these bonds requires significant energy, resulting in high melting points.
• For example, metals like tungsten and iron have very high melting points due to their strong metallic bonds.
• This property makes metals suitable for high-temperature applications.

Delocalized electrons play a crucial role in metallic bonding by allowing electrical conductivity and structural flexibility.

• These electrons are not bound to any specific atom, enabling them to move freely throughout the metallic structure.
• This movement facilitates the conduction of electricity and heat.
• Additionally, the presence of delocalized electrons contributes to the malleability and ductility of metals.