The Group 1 elements


Sodium and potassium are abundant in nature and can be found as salts such as chlorides. Lithium is a rare element that is found primarily in the mineral spodumene LiAlSi2O6. Rubidium and caesium are even rarer, but they can be found in reasonable amounts in minerals like zeolite pollucite (Cs2Al2Si4O12.nH2O). Electrolysis of molten metal chloride extracts sodium and lithium metals. Rubidium and caesium are obtained by reacting metal chloride with calcium or barium and potassium is obtained by reacting KCl with sodium metal.

The metals in Group 1 all have the valence electron configuration ns1. They are soft, conduct electricity and heat, and have low melting points as they progress through the group. Because each atom contributes only one electron to the molecular orbital band, their metallic bonding is weak, resulting in softness and low melting points (Section 3.19). This softness is especially noticeable in Cs, which melts at just 29°C. Because of their excellent thermal conductivities, liquid sodium and a sodium/potassium mixture have been used as coolants in some nuclear power plants. All of the elements have a body-centred cubic structure (Section 3.5), and their densities are low because this structure type is not closepacked and their atomic radii are large. Metals readily form alloys with one another, such as NaK, and with a variety of other metals, such as sodium/mercury amalgam. Table 11.1 summarizes some key characteristics.

Flame tests are commonly used to determine whether alkali metals and their compounds are present. Electronic transitions with energies in the visible part of the spectrum occur within the metal atoms and ions formed in flames, giving the flame its distinctive color:

A flame photometer can measure the intensity of the emission spectrum obtained from an alkali metal salt solution to provide a quantitative measurement of the element's concentration in the solution.

The trend in the atomic radii of Group 1 elements correlates with their chemical properties (Fig. 11.1). Because the valence shell becomes increasingly distant from the nucleus as the atomic radius increases from Li to Cs, the first ionization energy decreases down the group (Fig. 11.2; Section 1.9). The metals are reactive and form M ions more readily down the group because their first ionization energies are all low. Their reaction to water exemplifies this pattern:

The fact that both metals are denser than water, sink below the surface, and the sudden ignition of the hydrogen scatters the water violently is part of the reason for the explosive nature of the reaction of Rb and Cs with water.

The standard potentials of the couples M/M, which are all large and negative (Table 11.1), confirm the thermodynamic tendency to form M in aqueous conditions, indicating that the metals are readily oxidized. The surprising uniformity of alkali metal standard potentials can be explained by looking at the reduction half-reaction thermodynamic cycle (Fig. 11.3). As the radii of the ions increase, the enthalpies of sublimation and ionization both decrease (making oxidation more favorable); however, this trend is countered by a smaller enthalpy of hydration (making oxidation less favourable).


All of the elements must be stored in a hydrocarbon oil to avoid reactions with atmospheric oxygen, though Li, Na, and K can be handled in air for short periods of time; Rb and Cs must always be handled in an inert atmosphere.

Source : Shriver and Atkins' Inorganic Chemistry, Fifth Edition © 2010

Labels: Inorganic Chemistry

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