Physical Properties of s-Block Elements (Alkali Metals)
- All are metals with one electron in their outermost orbital and thus they form unipositive M+ ions (where M is an alkali metal) but hydrogen is an exception as it is not a metal even though it has s1 configuration.
- At normal temperature, all the metals adopt a body-centred cubic type of lattice with a coordination number of 8.
- The elements are soft and low melting. This is due to the contribution of only one electron per atom towards the metallic bonding. In solids, atoms are held with low values of cohesive energies.
- Their atomic and ionic radii increase on descending the group. This is due to the addition of a higher valence shell (which lies farther away from the nucleus) as we descend the group. The ionic radii are considerably smaller than the corresponding atomic radii. This is due to the complete removal of outermost shell containing only electron. Within in period, alkali atoms have the largest atomic radii.
- Their melting and boiling points decrease on descending the group.
- Their densities increase on descending the group with the exception of K whose density is less than that of Na. Because of the larger atomic radii, densities are low and lithium, sodium and potassium are even lighter than water.
- Their ionization energies decrease on descending the group. The values are low indicating the high reactivity of these metals. Cesium is the most reactive among the alkali metals. Francium being radioactive, is not studied in detail.
- Because of low ionization energies, the metals are strong reducing agents.
- The elements impart characteristics colours to the flame and can be detected (by a flame test) or determined (by flame photometry). In the flame, the electrons are excited to a higher energy level. When the electron drops back to its original energy level, it emits the absorbed energy as radiation whose wavelength lies in the visible region (400 nm to 750 nm).
- Because of low ionization energies, the elements when irradiated by light, The emitted electrons are called photoelectrons. Based on this phenomenon, potassium and cesium find use in photoelectric cells.
- The electro-negativity values for the elements are very small. When these elements react with other elements to form compounds, a large electro-negativity difference between the two atoms exists leading to the formation of ionic bonds.
- The standard electrode potentials. E0(M+ 1M), are highly negative indicating the powerful reducing nature of the metals. The high negative values are largely due to high hydration energy of their ions, which is in the order Li+ > Na+ > K+ > Rb+ > Cs+.
- Alkali metals dissolve in liquid ammonia and form a deep blue solution when dilute. The solution is conducting and acts as strong reducing agents. The reducing property is due to the presence of solvated (or ammoniated) cations and solvated electrons. Metals can be recovered by evaporating ammonia.
The blue colour is due to the solvated electrons. The paramagnetic behaviour of solution decreases with increasing concentration. This suggests that the electrons can associate to form diamagnetic electrons pairs.( 2e-e2-2 ).
The deep blue solution is moderately stable at temperatures where ammonia is still liquid. In the presence of light, the following reaction can also occur.
Na + NH3(I) NaNH2 + H2
This reaction is catalysed by transition metal salts (e.g iron (III) chloride)
- The conducting ability of alkali metal ions follows the order
Li+ < Na+ < K+ < Rb+ < Cs+