Not all solid substances that contain metal atoms conduct electricity. Electrical conductivity is defined by the ability of charged particles to move freely through a material when an electrical potential is applied. While pure metals are excellent conductors, the presence of a metal within a solid compound does not automatically guarantee this mobility. The determining factor is the chemical bonding structure, which dictates whether charge carriers—either electrons or ions—are available to move.
The Mechanism of Electrical Flow in Solids
For any solid material to conduct an electric current, it must possess mobile charged particles, which are typically electrons or ions. In metallic solids, the mechanism relies entirely on the movement of electrons. The atoms within a metal lattice share their outermost valence electrons, leading to the formation of a communal “sea” of delocalized electrons not bound to any single atom.
These delocalized electrons are free to roam throughout the entire crystal structure, acting as the charge carriers. When an external voltage is applied, this electron sea is set into motion. The positively charged metal ions remain fixed in their positions within the lattice, contributing to the structure but not the flow of current. This metallic bonding makes the movement of charge highly efficient.
Solids That Conduct Electricity
Solids that exhibit high electrical conductivity possess a vast supply of delocalized electrons. Pure metals, such as copper, aluminum, and silver, are prime examples of this structure. Copper is widely used in wiring because its single valence electron per atom is easily shared into the electron sea, allowing for minimal resistance to current flow.
Metal alloys, which are mixtures of two or more elements where at least one is a metal, also maintain this conducting structure. Substances like brass (a combination of copper and zinc) and various types of steel retain the metallic bond, ensuring a pool of mobile electrons remains available. These materials are utilized in applications where a specific combination of strength and conductivity is required.
Solids That Do Not Conduct Electricity
The reason many metal-containing solids do not conduct electricity lies in their specific bonding arrangement, which prevents the free movement of charge. The most common examples are solid ionic compounds, formed by the electrostatic attraction between positively charged metal ions and negatively charged non-metal ions. In a compound like solid sodium chloride (\(\text{NaCl}\)), the metal sodium has transferred its valence electron to chlorine, creating \(\text{Na}^{+}\) and \(\text{Cl}^{-}\) ions.
These charged ions are locked into a rigid, repeating crystal lattice structure by strong electrostatic forces. Because the electrons are localized and tightly held by the individual ions, they cannot move freely throughout the solid to carry a current. Furthermore, the ions themselves are fixed in their positions and cannot migrate under the influence of an electric field. Therefore, even though the solid contains metal atoms, the structure of the ionic bond makes it an electrical insulator.
Other metal compounds, such as certain metal oxides like magnesium oxide (\(\text{MgO}\)), also form these non-conducting ionic solids. The presence of the metal is irrelevant to conductivity in this state; the absence of delocalized electrons and the immobility of the ions are the deciding factors.
Conductivity Under Different States
While solid ionic compounds are insulators, their electrical properties change dramatically when their physical state is altered. When an ionic compound is heated above its melting point to become molten, the rigid crystal lattice breaks apart. This phase change frees the positively and negatively charged ions from their fixed positions.
The substance then becomes an excellent conductor of electricity because the now-mobile ions can migrate toward the oppositely charged electrodes when a voltage is applied. In this molten state, the conduction is due to the physical movement of the entire charged ions. A similar effect is observed when an ionic solid is dissolved in water, forming an aqueous solution. The solvent breaks the lattice, releasing the ions to move freely and carry charge through the solution. This is a key distinction from metallic conduction, which is carried solely by the movement of electrons.