Periods, Groups, and Blocks of the Periodic Table

In the periodic table, periods are the seven horizontal rows, groups are the 18 vertical columns, and blocks (s, p, d and f) are regions named after the type of orbital that holds the last electron added to the atoms. In brief: the period number of an element equals the number of its occupied electron shells, the group number of a main-group element tells its number of valence electrons, and the block tells which subshell is being filled.

Each position in the periodic table has a meaning. Reading it correctly allows you to find the electron configuration of an element without a table of configurations. This article explains periods, groups and blocks, compares the old and new group numbering systems, shows the block structure in a table, works through examples of locating elements and gives the history of the group names.

What are periods?

A period is a row of the table. Moving from left to right, the atomic number increases by one at each step, and the properties change gradually from strongly metallic to nonmetallic and end with a noble gas. After each noble gas, a new period starts with a highly reactive alkali metal. The length of a period depends on the number of orbitals that are filled: period 1 has 2 elements (1s), periods 2 and 3 have 8 each (s and p), periods 4 and 5 have 18 each (s, d and p) and periods 6 and 7 have 32 each (s, f, d and p).

What are groups?

A group is a column. The elements in a group have the same number of valence electrons (for the main groups) and form compounds with the same formulas. For example, the halogens in group 17 form compounds such as NaF, NaCl, NaBr and NaI. Properties change gradually down a group: atomic size increases, metallic character increases and the outer electrons become easier to remove.

What are the blocks?

• s-block: groups 1 and 2, plus helium. The last electron enters an s orbital.
• p-block: groups 13 to 18, except helium. The last electron enters a p orbital. It contains metals, metalloids and nonmetals.
• d-block: groups 3 to 12, the transition metals. The last electron enters a d orbital of the previous shell, (n − 1)d.
• f-block: the lanthanides and actinides. The last electron enters an f orbital, (n − 2)f.

Blocks and their properties

The number of columns in each block equals the maximum number of electrons that the subshell can hold: 2, 6, 10 and 14. The total of 2 + 6 + 10 + 14 = 32 explains why the longest periods have 32 elements.

How are groups numbered?

The International Union of Pure and Applied Chemistry (IUPAC) recommends numbering the groups 1 to 18 from left to right. Older systems used Roman numerals with the letters A and B (American) or other conventions (European), which sometimes caused confusion. In the old American system, the main groups were IA to VIIIA and the transition metals IB to VIIIB, so group 17 was VIIA. Note that in the old systems, the number of the main group equals the number of valence electrons, while in the new system the number of valence electrons for groups 13 to 18 is the group number minus 10.

Worked examples

1. From position to configuration. Selenium is in period 4, group 16. Period 4 means the highest shell is n = 4. Group 16 in the p-block has 6 valence electrons (4s² 4p⁴). The configuration is [Ar] 3d¹⁰ 4s² 4p⁴.
2. From configuration to position. The configuration [Ne] 3s² 3p³ has n = 3 as the highest shell, so it is in period 3. It has 2 + 3 = 5 valence electrons and ends in the p subshell, so it is in the p-block, group 15, phosphorus.
3. Block of a transition metal. Zirconium has the configuration [Kr] 4d² 5s², which ends with a d subshell filling, so it is in the d-block, period 5, group 4 (the second element of the d-block, so the group number is 2 + 2 = 4).
4. Counting elements in a block. The d-block has 10 columns in each of periods 4 to 7, so there are 10 d-block elements per period for periods 4 and 5, which makes 20 elements, and the same for periods 6 and 7 with the lanthanides and actinides counted separately.
5. Predicting formulas. Barium is in group 2, so it forms Ba²⁺. Its chloride is BaCl₂, like that of magnesium and calcium.

Where are the special cases?

Hydrogen is placed in group 1 because it has one electron in the s orbital, but it is a nonmetal and its properties are quite different from the alkali metals. Helium has the configuration 1s² and would belong to the s-block, but it is placed in group 18 because of its inertness. The lanthanides and actinides are usually shown below the main table, to avoid a very wide layout. Some tables place lutetium and lawrencium in group 3 under scandium and yttrium, while others place lanthanum and actinium there. This is a matter of continuing discussion, and IUPAC has published recommendations.

How did the layout develop?

Mendeleev’s table had eight groups with the main groups and subgroups drawn together in each, and no noble gases or transition-series separation. The long form of the table, with 18 columns, was developed in the early 20th century by Alfred Werner and others, and became standard after quantum theory explained its structure. IUPAC introduced the 1 to 18 numbering in 1988. In 1945 Glenn Seaborg proposed that the actinides form a series like the lanthanides, which led to the form of the table that we use today.

Which mistakes are common?

Mistake 1: Mixing up rows and columns

Periods run across, groups run down.

Mistake 2: Using the old group number as the IUPAC number

Group VIIA is group 17 in the IUPAC system.

Mistake 3: Thinking that the d-block is in the same shell as its period

The d subshell belongs to the previous shell: the 3d orbitals fill in period 4.

Mistake 4: Assuming that every group has a special name

Groups 3 to 12 are called transition metals as a whole, and some groups, such as group 4 or group 11, are named only after their first element in some textbooks.