s-block

Template:Periodic table (s-block) The s-block is a block in the periodic table that consists of the first two groups, namely the alkali metals and the alkaline earth metals. The elements in the s-block generally exhibit well-defined trends in their physical and chemical properties, changing steadily moving down the groups. Their properties can be most readily explained in terms of their electron configuration, with their valence electrons occupying s-orbitals. By this definition, hydrogen and helium are sometimes also considered to be part of the s-block.

Properties

The modern periodic law states that an element's chemical and physical properties is a periodic function of its atomic number. The long form of the periodic table is based on modern periodic law. The long form is divided into four blocks, s, p, d, and f. In an atom of an s-block element,the last electron enters the s-orbital of the outermost electron shell:

Group 1:

Hydrogen (Z=1), Lithium (Z=3), Sodium (Z=11), Potassium (Z=19), Rubidium (Z=37), Caesium (Z=55), Francium (Z=87).

Group 2:

Beryllium (Z=4), Magnesium (Z=12), Calcium (Z=20), strontium (Z=38), Barium (Z=56), Radium (Z=88).

Anomalous properties of Lithium:

Among alkali metal cations, lithium has the smallest size and highest polarizing power. This separates lithium from its family members. Its covalence character is higher, which is responsible for its solubility in organic solvent.

Diagonal relationship

Lithium and Beryllium, the first elements in group 1 and group 2 respectively, show some properties which are different from other members of their respective groups. In these anomalous properties they resemble the second element of the succeeding group. Thus lithium shows similarities to magnesium and beryllium to aluminum in many of their properties.The cause of this diagonal relationship is this that as we go down the group atomic size increases(due to the increase in number of orbitals) and as we go from left to right in a period atomic size decreases(due to increasing number if electrons in the same orbital).due to this opposing factors the size of 2nd group elements is comparable to their respective diagonal elements of 3rd group(and not to the elements of their own group as going further down in a group electrons start entering d-orbital and therefore are very large as compare to the 2nd group elements of their same group)and therefore show some similarities.

       Group→      1    2   13    14
       Period 2   Li   Be   B     C
       Period 3   Na   Mg   Al    Si

This type of relationship in the periodic table is known as a diagonal relationship. The diagonal relationship is due to the similarities in ionic sizes and charge/radius ratio of the element. The similarity between lithium and magnesium is due to their same sizes. Lithium has a radius of 152pm and magnesium has a radius of 160pm.

Anomalous behaviour of lithium

Lithium - the first element of group 1 - differs from the rest of this group in many respects. This anomalous behaviour of lithium is due to the following reasons: a. small size of lithium atom and its ion. b. higher polarization power of Li⁺
(i.e. charge size ratio) resulting in increased covalent character of its compounds which is responsible for their solubility in organic solvents c. comparatively high ionisation enthalpy and low electropositive character of lithium as compared to other alkali metals d. non availability of d-orbitals in its valence shell e. strong intermetallic bonding Some of the properties in which lithium differs from other members of its group illustrating its anomalous behaviour are as follows: 1. Lithium is harder than sodium and potassium which are so soft that they can be cut by a knife. 2. The melting and boiling points of lithium are comparatively high. 3. Lithium forms monoxide with oxygen, other alkali form peroxide and superoxide. 4. Lithium combines with nitrogen to form nitrides, while other alkali metals do not. 5. Lithium Chloride is deliquescent and crystallizes as a hydrate LiCl.2H2O, whereas other alkali metal chlorides do not form hydrates.

See also

• Electron configuration