Metals are a distinct class of elements unified by specific physical and chemical behaviors. They are characterized by their ability to readily conduct both heat and electricity, along with a tendency to exhibit a reflective surface, often called luster. Metals represent the majority of elements on the periodic table, forming the foundation for countless materials essential to modern technology and industry. Their defining characteristics stem from a unique atomic structure.
Where Metals Live on the Periodic Table
The elements on the periodic table are organized by their properties, with metals occupying the entire left side and center. A distinct, diagonal “stair-step” line, beginning at Boron, separates the metallic elements from the nonmetals found on the right side. Elements bordering this line are called metalloids, possessing properties of both categories.
The most reactive metals are found on the far left, comprising the Alkali Metals (Group 1) and the Alkaline Earth Metals (Group 2) in the s-block. Moving inward, the large central block is dominated by the Transition Metals, which span Groups 3 through 12 and form the d-block. These transition elements often form compounds with vibrant colors and exhibit multiple possible positive charges in reactions.
Below the main body of the table are the Inner Transition Metals, known as the Lanthanides and Actinides, which constitute the f-block elements. While most metals are grouped on the left, a few elements on the p-block’s lower-left side, such as aluminum, are also classified as metals. This placement structure visually demonstrates that over three-quarters of all known elements possess metallic properties.
Defining Physical Properties
The physical properties that define metals—such as electrical conductivity, malleability, and luster—are a direct consequence of their internal bonding structure. Metallic bonding is explained by the “sea of electrons” model. In this model, valence electrons are delocalized, meaning they are free to move throughout the entire solid structure rather than being tethered to a single atom. This creates a lattice of positively charged metal ions suspended within a mobile cloud of shared electrons.
This “sea” of electrons is responsible for the high electrical and thermal conductivity characteristic of metals. When an electrical voltage is applied, the delocalized electrons instantly flow, carrying the charge. Similarly, when heat is introduced, these energetic electrons rapidly transport thermal energy across the material.
The bonding model also accounts for malleability (the ability to be hammered into thin sheets) and ductility (the ability to be drawn into a wire). Since the positive ion cores are held together by the non-directional electron sea, layers of atoms can slide past one another without shattering the structure. The surrounding electron cloud reforms the bond in the new position, preventing strong repulsion.
Metals possess high densities and high melting and boiling points compared to nonmetals. The strong attractive force between the positive ions and the electron sea requires significant energy to break. Metallic luster is a result of the free electrons on the surface absorbing and quickly re-emitting light across the visible spectrum.
How Metals React Chemically
The chemical behavior of metals is governed by their tendency to lose electrons during a reaction. Metals generally have a low ionization energy, which is the energy required to remove an electron from an atom. This low energy requirement means their outer valence electrons are easily shed to achieve a more stable electron configuration.
When a metal atom loses one or more valence electrons, it becomes a positively charged ion, known as a cation. The number of electrons lost determines the charge of the cation. For example, Alkali metals lose one electron to form a 1+ ion, while Alkaline Earth metals lose two to form a 2+ ion. This process of losing electrons is defined as oxidation in a chemical reaction.
Because they readily lose electrons, metals are highly reactive with nonmetals, which tend to gain electrons. This exchange leads to the formation of ionic bonds, where the positively charged metal cation is attracted to the negatively charged nonmetal anion. Metals like sodium and potassium are highly reactive because their single valence electron is easily removed.
The chemical stability of a metal is inversely related to its ionization energy. Metals with higher ionization energies, such as noble metals like gold and platinum, are far less reactive. However, common metals are subject to chemical processes like corrosion, which is a form of oxidation. Rusting, for example, occurs when iron loses electrons to oxygen atoms in the presence of water, forming iron oxide.