3.6 Electronic Configurations and Types of Elements: s-, p-, d-, f- Blocks

• The aufbau (build-up) principle and the electronic configuration of atoms provide a theoretical foundation for the periodic classification.
• The elements in a vertical column of the Periodic Table constitute a group or family and exhibit similar chemical behaviour. This similarity arises because these elements have the same number and same distribution of electrons in their outermost orbitals. We can classify the elements into four blocks viz.,
  • \(s\)-block,
  • \(p\)-block,
  • \(d\)-block and
  • \(f\)-block depending on the type of atomic orbitals that are being filled with electrons.

This is illustrated in Fig. 3.3.

• We notice two exceptions to this categorisation. Strictly, helium belongs to the \(s\)-block but its positioning in the \(p\)-block along with other group 18 elements is justified because it has a completely filled valence shell \(\left(1 s^2\right)\) and as a result, exhibits properties characteristic of other noble gases.
• The other exception is hydrogen. It has only one \(s\)-electron and hence can be placed in group 1 (alkali metals). It can also gain an electron to achieve a noble gas arrangement and hence it can behave similar to a group 17 (halogen family) elements. Because it is a special case, we shall place hydrogen separately at the top of the Periodic Table as shown in Fig. 3.2 and Fig. 3.3. 

The \(s\)–Block Elements

• General electronic configuration is \({ns}^{1-2}\).
• The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have \(n s^1\) and \(n s^2\) outermost electronic configuration belong to the \(s\)-Block Elements.
• Group 1 elements are known as alkali metals because they react with water to form alkali. Group 2 elements are known as alkaline earth metals because their oxides react with water to form alkali and these are found in the soil or earth. The total number of \(s\)-block elements are 14.
• They are all reactive metals with low ionization enthalpies. They lose the outermost electron(s) readily to form \(1+\) ion (group-1, in the case of alkali metals) or 2+ ion (group-2, in the case of alkaline earth metals).
• The metallic character and the reactivity increase as we go down the group. Because of high reactivity, they are never found pure in nature. The compounds of the s-block elements, with the exception of those of lithium and beryllium are predominantly ionic.

The \(p\)–Block Elements

• General electronic configuration is \(n s^2 n p^{1-6}\).
• The \(p\)-Block Elements comprise those belonging to Group 13 to 18.
• The \(p\)-Block Elements together with the \(s\)-Block Elements are called the Representative Elements or Main Group Elements.
• The outermost electronic configuration varies from \(n s^2 n p^1\) to \(n s^2 n p^6\) in each period. At the end of each period is a noble gas element with a closed valence shell \(n s^2 n p^6\) configuration.
• Group 18 elements are the noble gases due to the completely filled valence shell. All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. The noble gases thus exhibit very low chemical reactivity.
• Preceding the noble gas family are two chemically important groups of non-metals. They are the halogens (Group 17) and the chalcogens (Group 16). These two groups of elements have highly negative electron gain enthalpies and readily add one or two electrons respectively to attain the stable noble gas configuration.
• The non-metallic character increases as we move from left to right across a period and metallic character increases as we go down the group.

The \(d\)–Block Elements (Transition Elements)

• General electronic configuration is \((n-1) {d}^{1-10} {~ns}^{0-2}\).
• These are the elements of Group 3 to 12 in the centre of the Periodic Table. These are characterised by the filling of inner \(d\) orbitals by electrons and are therefore referred to as d-Block Elements.
• These elements have the general outer electronic configuration \((n-1) d^{1-10} n s^{1-2}\) except for Pd where its electronic configuration is \(4 d^{10} 5 s^0\).
• They are all metals. They mostly form coloured ions, exhibit variable valence (oxidation states), paramagnetism and oftenly used as catalysts.
• However, \(\mathrm{Zn}, \mathrm{Cd}\) and \(\mathrm{Hg}\) which have the electronic configuration, \((n-1) d^{10} n s^2\) do not show most of the properties of transition elements because in these elements \(d\) orbitals are fully filled.
• They (transition metals) form a bridge between the chemically active metals of \(s\)-block elements and the less active elements of Groups 13 and 14 and thus take their familiar name “Transition Elements”.

The \(f\)–Block Elements (Inner-Transition Elements)

• General electronic configuration is \((n-2) f^{1-14}(n-1) d^{0-1}\) \({ns}^2\).
• The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids, \(\operatorname{Ce}(Z=58)-\operatorname{Lu}(Z=71)\) and Actinoids, \(\operatorname{Th}(Z=90)-\operatorname{Lr}(Z=103)\) are characterised by the outer electronic configuration \((n-2) f^{1-14}\) \((n-1) d^{0-1} n s^2\).
• The last electron added to each element is filled in \(f\)-orbital. These two series of elements are hence called the Inner Transition Elements ( \(f\)-Block Elements).
• They are all metals. Within each series, the properties of the elements are quite similar. The chemistry of the early actinoids is more complicated than the corresponding lanthanoids, due to the large number of oxidation states possible for these actinoid elements. Actinoid elements are radioactive. Many of the actinoid elements have been made only in nanogram quantities or even less by nuclear reactions and their chemistry is not fully studied.
• The elements after uranium are called Transuranium Elements.

Example 3.3: The elements \(Z=117\) and 120 have not yet been discovered. In which family / group would you place these elements and also give the electronic configuration in each case.

Answer: We see from Fig. 3.2, that element with \(Z\) \(=117\), would belong to the halogen family (Group 17) and the electronic configuration would be [Rn] \(5 f^{14} 6 d^{10} 7 s^2 7 p^5\). The element with \(Z=120\), will be placed in Group 2 (alkaline earth metals), and will have the electronic configuration [Uuo] \(8 s^2\).

Metals, Non-metals and Metalloids

In addition to displaying the classification of elements into \(\boldsymbol{s}\)-, \(\boldsymbol{p}\)-, \(\boldsymbol{d}\)-, and \(\boldsymbol{f}\)-blocks, Fig. 3.3 shows another broad classification of elements based on their properties. The elements can be divided into Metals and Non-Metals.

Metals:

• Metals comprise more than \(78 \%\) of all known elements and appear on the left side of the Periodic Table.
• Metals are usually solids at room temperature [mercury is an exception; gallium and caesium also have very low melting points (303K and 302K, respectively)].
• Metals usually have high melting and boiling points.
• They are good conductors of heat and electricity.
• They are malleable (can be flattened into thin sheets by hammering) and ductile (can be drawn into wires).

Non-metals:

• Non-metals are located at the top right-hand side of the Periodic Table. In fact, in a horizontal row, the property of elements change from metallic on the left to non-metallic on the right.
• Non-metals are usually solids or gases at room temperature with low melting and boiling points (boron and carbon are exceptions).
• They are poor conductors of heat and electricity.
• Most nonmetallic solids are brittle and are neither malleable nor ductile.
• The elements become more metallic as we go down a group; the nonmetallic character increases as one goes from left to right across the Periodic Table.
• The change from metallic to non-metallic character is not abrupt as shown by the thick zig-zag line in Fig. 3.3.
• Metalloids: The elements (e.g., silicon, germanium, arsenic, antimony and tellurium) bordering this line and running diagonally across the Periodic Table show properties that are characteristic of both metals and nonmetals. These elements are called Semi-metals or Metalloids.

Example 3.4: Considering the atomic number and position in the periodic table, arrange the following elements in the increasing order of metallic character: \(\mathrm{Si}, \mathrm{Be}, \mathrm{Mg}, \mathrm{Na}, \mathrm{P}\).

Answer: Metallic character increases down a group and decreases along a period as we move from left to right. Hence the order of increasing metallic character is: \(\mathrm{P}<\mathrm{Si}<\) \(\mathrm{Be}<\mathrm{Mg}<\mathrm{Na}\).