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8.3: Metallic Bonds

Difficulty Level: At Grade Created by: CK-12
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Lesson Objectives

  • Describe the electron-sea model of metallic bonding.
  • Explain how metallic bonding is responsible for the conductivity and luster of metals.
  • Explain why metals are malleable and ductile, while crystalline ionic compounds are not.
  • Describe how metal atoms are arranged, including the three most common packing systems.
  • Identify some common alloys and explain their importance.

Lesson Vocabulary

  • alloy
  • closest packing
  • metallic bond

Check Your Understanding

Recalling Prior Knowledge

  • How are the cations and anions in an ionic crystal arranged?
  • What happens to an ionic crystal when it is put under a large stress?

The bonding that occurs in a metal is responsible for its distinctive properties, including luster, malleability, ductility, and excellent conductivity.

The Metallic Bond

Pure metals are crystalline solids, but unlike ionic compounds, every point in the crystal lattice is occupied by an identical atom. The electrons in the outer energy levels of a metal are mobile and capable of drifting from one metal atom to another. This means that the metal is more properly viewed as an array of positive ions surrounded by a “sea of mobile valence electrons.” Electrons which are capable of moving freely throughout the empty orbitals of the metallic crystal are called delocalized electrons (Figure below). A metallic bond is the attraction of the stationary metal cations to the surrounding mobile electrons.

In a metal, the stationary metal cations are surrounded by a sea of mobile valence electrons that are not associated with any one cation.

Properties of Metals

The metallic bonding model explains the physical properties of metals. Metals conduct electricity and heat very well because of their free-flowing electrons. As electrons enter one end of a piece of metal, an equal number of electrons flow outward from the other end. When light is shone onto the surface of a metal, its electrons absorb small amounts of energy and become excited into one of its many empty orbitals. The electrons immediately fall back down to lower energy levels and emit light. This process is responsible for the high luster of metals (Figure below).

The American Platinum Eagle is the official platinum bullion coin of the United States and was first minted in 1997. The luster of a metal is due to its metallic bonds.

Recall that ionic compounds are very brittle. Application of a force results in like-charged ions in the crystal coming too close to one another, causing the crystal to shatter. When a force is applied to a metal, the free-flowing electrons can slip in between the stationary cations and prevent them from coming in contact. Imagine ball bearings that have been coated with oil sliding past one another. As a result, metals are very malleable and ductile. They can be hammered into shapes, rolled into thin sheets, or pulled into thin wires.

Crystal Structures of Metals

If you wanted to make a stack of identical spheres, you might come up with an arrangement like the one shown in Figure below.

When identical spheres are stacked, each successive layer fits into the small spaces where different spheres come together. Closest packing minimizes the amount of empty space and is how metal atoms are arranged in a crystal.

This orderly and regular arrangement of the metal balls minimizes the empty space between them. Closest packing is the most efficient way to pack spherical objects. The atoms in a metal crystal are arranged in similar patterns, called close-packed structures. Pure metals adopt one of several related close-packed structures, as shown in Figure below.

Most pure metals naturally adopt one of these three closest packing arrangements.

On the far left is the body-centered cubic (bcc) structure. In that crystal, metal atoms occupy the eight corners of a cube along with one atom in the very center. The coordination number of each atom in the body-centered cubic structure is 8. In the face-centered cubic (fcc) structure, there are eight atoms at each corner of the cube and six atoms in the center of each face. The coordination number of each atom in the face-centered cubic structure is 12. The atoms in a hexagonal close-packed (hcp) structure also have a coordination number of 12, but crystals of this type are hexagonally shaped rather than cubic.


An alloy is a mixture composed of two or more elements, at least one of which is a metal. You are probably familiar with some alloys such as brass and bronze (Figure below). Brass is an alloy of copper and zinc. Bronze is an alloy of copper and tin. Alloys are commonly used in manufactured items because the properties of these metal mixtures are different, and sometimes more useful, than those of a pure metal. For example, bronze is harder than copper and more easily cast. Brass is very malleable, and its acoustic properties make it useful for musical instruments.

Bronze, an alloy of copper and tin, has been in use since ancient times. The Bronze Age saw the increased use of metals rather than stone for weapons, tools, and decorative objects. Brass, an alloy of copper and zinc, is widely used in musical instruments like the trumpet and trombone.

Steels are a very important class of alloys. The many types of steels are primarily composed of iron, with various amounts of the elements carbon, chromium, manganese, nickel, molybdenum, and boron. Steels are widely used in building construction because of their strength, hardness, and resistance to corrosion. Most large modern structures like skyscrapers and stadiums are supported by a steel skeleton (Figure below).

The Willis Tower (formerly called the Sears Tower) in Chicago was once the tallest building in the world and is still the tallest in the Western Hemisphere. The use of steel columns makes it possible to build taller, stronger, and lighter buildings.

Alloys can be one of two general types. In one type, called a substitutional alloy, the various atoms simply replace each other in the crystal structure. In another type, called an interstitial alloy, smaller atoms such as carbon fit in between the larger atoms in the crystal packing arrangement.

Lesson Summary

  • A metallic bond involves the attraction of stationary metal cations to a surrounding sea of mobile valence electrons. Metallic bonding directly contributes to properties such as high electrical and thermal conductivity, malleability, ductility, and luster.
  • Metal atoms are arranged in regular, orderly patterns that often correspond to various closest packing arrangements for spheres.
  • Alloys are mixtures of metals. Many alloys are widely used because their properties are often more useful for certain applications than those of pure metals.

Lesson Review Questions

Reviewing Concepts

  1. Which of the following elements displays metallic bonding?
    1. Se
    2. Ti
    3. Ru
    4. Te
  2. Explain the behavior of electrons in a metal.
  3. How does a metallic bond contribute to the electrical conductivity of metals?
  4. Name the three common closest packing arrangements of metal atoms in a crystal. Give an example of a metallic element that crystallizes in each form.


  1. Why are ionic crystals brittle, while most metals are malleable?
  2. What are some advantages of using steel rather than iron?
  3. Sodium has one valence electron per atom, while magnesium has two. Predict whether sodium or magnesium has stronger metallic bonds. How do you think that you could test your hypothesis?
  4. What are the coordination numbers of a metal atom in a body-centered cubic structure and in a face-centered cubic structure? In which type are the atoms more closely packed? Iridium is one of the densest known elements at 22.6 g/cm3. In which crystal structure do you think iridium is more likely to crystallize?

Further Reading / Supplemental Links

Points to Consider

Molecular compounds are a class of substances that take the form of individual molecules.

  • Which types of elements make up molecular compounds?
  • How is the chemical bonding within a molecular compound different from the bonding that occurs in ionic compounds or metals?
  • How are the physical properties of molecular compounds different from those of ionic compounds or metals?

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Date Created:
Aug 02, 2012
Last Modified:
Aug 30, 2016
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