In this tutorial, you will learn about metallic bonding, its characteristics and properties, and even learn some examples!
Topics Covered in Other Articles
- Ionic vs. Covalent Bonds
- Polar vs. Non-polar Bonds
- Metals, Nonmetals, Metalloids
- Bond Enthalpy & Bond Energy
- Polar-Covalent Bonds
Definition of Metallic Bonding
The short answer: metallic bonding is a type of chemical bonding between two or more metal atoms, which arises from the attraction between positively charged metal nuclei and their delocalized valence electrons.
In the rest of this article, we will take a look at the different parts of this definition and break down what it means, and explore the consequences of metallic bonding interactions!
What is a Metallic Bond?
Metallic bonds are not discrete directional bonds between specific atoms, so it often makes sense to talk about metallic “bonding” rather than individual bonds. But what are the characteristics of this kind of bonding? First, as the name implies, this type of bonding is found in metals and metal alloys. Bonding between metals and nonmetals is usually ionic, while bonding among nonmetals is usually covalent. In substances containing only metal atoms, however, the interaction involved has a different character, and we call it metallic bonding.
Characteristics of metallic bonding
Metallic elements have low electronegativity, which means that they hold onto their valence electrons loosely. This is the reason why they give up valence electrons around electronegative atoms to form ionic compounds. As a result, when there are only metal atoms around, the valence electrons of one atom are hard to distinguish from those of its neighbors. All the atoms in a piece of metal share these electrons, and they can flow freely through the whole material. Many people use the term “sea of electrons” to refer to this shared pool of valence electrons. In the figure below, the yellow contour lines depict this sea of electrons approximately in terms of electron density.
There are some cases where electrons can flow through covalent compounds, such as in conjugated pi-systems. However, this comes from delocalization of electrons along specific bonds, and not from a general sharing of electrons through the whole material.
Metallic Crystal Structure
Another characteristic of metallic bonding is that all nearest neighbor atoms bond identically. This means that metal crystals usually organize in close-packed patterns, like hexagonal close-packing or cubic close-packing.
Many of the properties of metals come from the nature of metallic bonding. Among these are electrical conductivity, thermal conductivity, and sheen.
Electrical conduction is the transfer of electrons through a material in response to a potential difference. Electrical conductivity arises primarily from electrons in the conduction band of a metal. Because electrons always populate the conduction band in metals, they can always transfer charge efficiently.
The high thermal conductivity of metals also comes from electron mobility. The atoms themselves contribute to heat transfer, but the electrons also absorb and transfer kinetic energy.
The sheen of metals comes from the high density of electrons at the surface of the solid. Incoming photons are more likely to reflect off these electrons than transmit or absorb.
Examples of Metallic Bonding
- Alkali metals, like sodium and potassium
- Alkaline earth metals, like calcium and magnesium
- Transition metals, like iron, titanium, gold, tungsten, copper and zinc
- Alloys, like bronze, pewter, stainless steel, electrum, and duralumin