The arrangement of elements on the periodic table into vertical columns, known as groups or families, is based on similarities in their chemical behavior. These shared characteristics arise primarily from having the same number of electrons in their outermost shell. The highly reactive metals found in the first column are designated the alkali metals, a family defined by a unique set of physical and chemical attributes.
Group 1: The Alkali Metals on the Periodic Table
The elements formally designated as alkali metals are Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), and Francium (Fr). These six elements occupy Group 1, the far-left column of the periodic table. Their placement reflects a commonality in their electron configuration, which dictates their chemical identity.
Hydrogen (H) is also positioned at the top of Group 1, but it is not considered an alkali metal. Unlike all the other elements in the group, hydrogen exists as a diatomic gas at standard temperature and pressure, not a solid metal. Furthermore, hydrogen’s ionization energy, the energy required to remove its single electron, is significantly higher than that of the alkali metals. This difference means hydrogen does not readily form a positive ion like the other elements in Group 1.
Shared Physical and Electronic Characteristics
The defining electronic characteristic of every alkali metal is the presence of a single valence electron in its outermost shell. Because this single electron is easily lost, the alkali metals have the lowest first ionization energies of all elements in the periodic table.
The physical properties of these elements are closely related to this electronic structure. Alkali metals are the largest atoms within their respective periods, a factor that contributes to their low density. Lithium, sodium, and potassium, for example, are all less dense than water and will float on it.
Alkali metals are remarkably soft, easily cut with a knife, exposing a bright, silvery surface that tarnishes quickly in air. The metallic bonding is relatively weak due to having only one valence electron per atom. This weak bonding results in comparatively low melting and boiling points, which decrease as one moves down the group. For example, Cesium melts at only 28.4 °C, just above room temperature. They still exhibit characteristic metallic behavior by being excellent conductors of both heat and electricity.
Chemical Reactivity and Practical Applications
The high chemical reactivity of the alkali metals is a direct consequence of their single, loosely held valence electron. They readily lose this electron to form a cation with a +1 charge, achieving the stable electron configuration of a noble gas. This tendency to donate an electron makes them very strong reducing agents.
Their reactivity increases as one moves down the group, from Lithium to Cesium, because the single valence electron is progressively farther from the nucleus and is easier to remove. A common chemical behavior is their vigorous reaction with water, which produces a metal hydroxide and hydrogen gas. This reaction is so exothermic that the heat generated can ignite the hydrogen gas, especially for the heavier elements.
The chemical behavior of alkali metals is utilized in various applications:
- Lithium is a component in rechargeable batteries due to its low atomic mass and high energy density.
- Sodium and potassium are biologically important, playing a role in nerve impulse transmission and maintaining fluid balance in the body. Sodium chloride (table salt) is a common compound derived from sodium.
- Cesium is used in highly precise atomic clocks, which define the modern second.
- Francium, the heaviest alkali metal, is extremely rare and highly radioactive, meaning it has no practical applications outside of scientific study.