# Question #b5bf3

Oct 8, 2015

s, p, d and f-blocks are four divisions of the periodic table.

#### Explanation:

The whole periodic table is divided into four blocks. Groups 1 and 2 are called the s-block, because all the elements in those groups have electronic configurations ending is the s-orbital. For example, in Group 1, lithium has an electronic configuration of $1 {s}^{2} 2 {s}^{1}$ and Francium has an electronic configuration $\left[R n\right] 7 {s}^{1}$. In Group 2, Beryllium has electronic configuration $1 {s}^{2} 2 {s}^{2}$ and Radium has electronic configuration $\left[R n\right] 7 {s}^{2}$. As you can see, in both these groups, all the elements have electronic configuration such that they end in a s-orbital.

The transition metals, from Group 3 to Group 12, are called the d-block. The elements in these groups have electronic configuration such that they end in the d-orbital. Example:
$S c$ : $1 {s}^{2} 2 {s}^{2} 2 {p}^{6} 3 {s}^{2} 3 {p}^{6} 4 {s}^{2} 3 {d}^{1}$
$H g$ : $\left[X e\right] 5 {d}^{10}$

The elements from Groups 13 to 18 are called the p-block. They have electronic configuration that end in p-orbital. Example:
$B$ : $1 {s}^{2} 2 {s}^{2} 2 {p}^{1}$
$R n$ : $\left[X e\right] 6 {p}^{6}$
*n.b.: Helium is an exception. Although in group 18, its electronic configuration is $1 {s}^{2} 2 {s}^{2}$

In between Group 2 and the transition metals, there is a narrow band of elements in the 6th and 7th periods. They are called the Lanthanides and the Actinides respectively. These elements are the f-block. Their electronic configuration always end in the f-orbital. Example:
$L a$ : $\left[X e\right] 4 {f}^{1}$
$L r$ : $\left[R n\right] 4 {f}^{14}$

Note, I haven't written the full electronic configuration. I've only shown the last part of the configuration for elements with large atomic numbers, because that's the only relevant part in your question.