Metallic Crystals


Learning Objective

  • Describe metallic crystals.

Key Points

    • Atoms in metals lose electrons to form cations. Delocalized electrons surround the ions. Metallic bonds (electrostatic interactions between the ions and the electron cloud) hold the metallic solid together. Atoms are arranged like closely packed spheres.
    • Because outer electrons of metal atoms are delocalized and highly mobile, metals have electrical and thermal conductivity. The free electron model can be used to calculate electrical conductivity as well as the electrons’ contribution to the heat capacity and heat conductivity of metals.
    • Metals are ductile, or capable of plastic deformation. Hooke’s law describes reversible elastic deformation in metals, in which the stress is linearly proportional to the strain. Forces larger than the elastic limit, or heat, may cause an irreversible deformation of the object.
    • In general, metals are denser than nonmetals. This is due to the tightly packed crystal lattice of the metallic structure. The larger the amounts of delocalized electrons, the stronger the metallic bonds are.


  • metalAny of a number of chemical elements in the periodic table that form a metallic bond with other metal atoms. It is generally shiny, malleable, and a conductor of heat and electricity.
  • metallic bondA chemical bond in which mobile electrons are shared over many nuclei; this leads to electrical conduction.

Metallic Properties

In a metal, atoms readily lose electrons to form positive ions (cations). These ions are surrounded by delocalized electrons, which are responsible for conductivity. The solid produced is held together by electrostatic interactions between the ions and the electron cloud. These interactions are called metallic bonds. Metallic bonding accounts for many physical properties of metals, such as strength, malleability, ductility, thermal and electrical conductivity, opacity, and luster.

Metallic BondingLoosely bound and mobile electrons surround the positive nuclei of metal atoms.

Understood as the sharing of “free” electrons among a lattice of positively charged ions (cations), metallic bonding is sometimes compared to the bonding of molten salts; however, this simplistic view holds true for very few metals. In a quantum-mechanical view, the conducting electrons spread their density equally over all atoms that function as neutral (non-charged) entities.

Atoms in metals are arranged like closely-packed spheres, and two packing patterns are particularly common: body-centered cubic, wherein each metal is surrounded by eight equivalent metals, and face-centered cubic, in which the metals are surrounded by six neighboring atoms. Several metals adopt both structures, depending on the temperature.

Metals in general have high electrical conductivity, high thermal conductivity, and high density. They typically are deformable (malleable) under stress, without cleaving. Some metals (the alkali and alkaline earth metals) have low density, low hardness, and low melting points. In terms of optical properties, metals are opaque, shiny, and lustrous.

Melting Point and Strength

The strength of a metal derives from the electrostatic attraction between the lattice of positive ions and the “sea” of valence electrons in which they are immersed. The larger the nuclear charge (atomic number) of the atomic nucleus, and the smaller the atom’s size, the greater this attraction. In general, the transition metals with their valence-level d electrons are stronger and have higher melting points:

  • Fe, 1539°C
  • Re, 3180 °C
  • Os, 2727 °C
  • W, 3380°C.

The majority of metals have higher densities than the majority of nonmetals. Nonetheless, there is wide variation in the densities of metals. Lithium (Li) is the least dense solid element, and osmium (Os) is the densest. The metals of groups IA and IIA are referred to as the light metals because they are exceptions to this generalization. The high density of most metals is due to the tightly packed crystal lattice of the metallic structure.

Electrical Conductivity: Why Are Metals Good Conductors?

In order for a substance to conduct electricity, it must contain charged particles (charge carriers) that are sufficiently mobile to move in response to an applied electric field. In the case of ionic compounds in water solutions, the ions themselves serve this function. The same thing holds true of ionic compounds when melted. Ionic solids contain the same charge carriers, but because they are fixed in place, these solids are insulators.

In metals, the charge carriers are the electrons, and because they move freely through the lattice, metals are highly conductive. The very low mass and inertia of the electrons allows them to conduct high-frequency alternating currents, something that electrolytic solutions cannot do.

Electrical conductivity, as well as the electrons’ contribution to the heat capacity and heat conductivity of metals, can be calculated from the free electron model, which does not take the detailed structure of the ion lattice into account.

Mechanical properties

Mechanical properties of metals include malleability and ductility, meaning the capacity for plastic deformation. Reversible elastic deformation in metals can be described by Hooke’s Law for restoring forces, in which the stress is linearly proportional to the strain. Applied heat, or forces larger than the elastic limit, may cause an irreversible deformation of the object, known as plastic deformation or plasticity.

Metallic solids are known and valued for these qualities, which derive from the non-directional nature of the attractions between the atomic nuclei and the sea of electrons. The bonding within ionic or covalent solids may be stronger, but it is also directional, making these solids brittle and subject to fracture when struck with a hammer, for example. A metal, by contrast, is more likely to be simply deformed or dented.

Although metals are black due to their ability to absorb all wavelengths equally, gold (Au) has a distinctive color. According to the theory of special relativity, increased mass of inner-shell electrons that have very high momentum causes orbitals to contract. Because outer electrons are less affected, blue-light absorption is increased, resulting in enhanced reflection of yellow and red light.

GoldGold is a noble metal; it is resistant to corrosion and oxidation.