Periodic table of elements and the periodic law

Introduction

By now you know that atoms are built from protons, neutrons and electrons, and that electrons fill orbitals in a predictable order. Today we bring this together and look at the periodic table itself — how it's organised, why it's shaped the way it is, and why chemists call the patterns in it the periodic law. Understanding this structure will let you predict properties of elements without memorising hundreds of facts.

What is the periodic law?

The periodic law states that when elements are arranged by increasing atomic number, their physical and chemical properties repeat in a regular, predictable pattern. This repetition is called periodicity. It happens because electron configurations repeat in a cyclical way as you move through the table — elements in the same group have the same number of valence electrons, so they behave similarly (Talbot, p.233).

Periods, groups and blocks

• Periods are the horizontal rows. The period number tells you the highest energy level (shell) occupied by electrons in that element's ground-state configuration.
• Groups are the vertical columns. Elements in a group share the same number of valence electrons, which is why they show similar chemical behaviour.
• Blocks (s, p, d, f) refer to which sublevel is being filled last, according to the aufbau principle.

Worked example 1: Period lengths and orbitals filled

Why does period 4 contain 18 elements instead of 8?

Across period 4, electrons fill the 4s, 3d, and 4p orbitals in turn — that's 2 + 10 + 6 = 18 electrons, hence 18 elements, from potassium (1s²2s²2p⁶3s²3p⁶4s¹) to krypton (Talbot, p.239). Compare this with period 2 or 3, where only s and p orbitals (2 + 6 = 8) are filled, giving 8 elements.

(Talbot, p.239)

Periods 6 and 7 are especially long (32 elements each) because the f-block (lanthanoids and actinoids) is also being filled — these are usually pulled out and shown as a separate strip below the main table just to keep it compact (Talbot, p.239).

Trends down a group

As you go down a group, atoms gain extra shells, and this affects metallic vs non-metallic character:

• Down group 1, metallic character increases — elements become more reactive as metals (easier to lose their single valence electron).
• Down group 17, non-metallic character decreases — elements become less reactive as non-metals (harder to attract one more electron as atoms get larger) (Talbot, p.233).

There isn't a sharp line between "metal" and "non-metal" — instead there's a continuum, reflected in how oxides behave: metal oxides tend to be basic, non-metal oxides tend to be acidic, and elements in between (like aluminium) form amphoteric oxides that can react as both acid and base (Talbot, p.233).

Worked example 2: Predicting oxide behaviour

Question: Would you expect sodium oxide (Na₂O) or sulfur trioxide (SO₃) to be more acidic?

Answer: Sodium is a metal on the far left of the table, so Na₂O is basic. Sulfur is a non-metal on the right, so SO₃ is acidic. This fits the general left-to-right trend from basic → amphoteric → acidic oxides.

A note on symbols and new elements

Chemical symbols were devised by Berzelius so that chemistry could be understood across languages — for example, tungsten's symbol W comes from the German name wolfram, not the English name (Talbot, p.235). The table is still growing too: element 118 (oganesson) was synthesized in 2002, and scientists are currently attempting to create element 119 (Talbot, p.242).

Key takeaways

• The periodic law: properties of elements repeat periodically when arranged by atomic number.
• Period number = highest occupied energy level; group = number of valence electrons.
• Period lengths (2, 8, 8, 18, 18, 32, 32) come directly from the order in which sublevels (s, p, d, f) are filled.
• Down group 1, metallic character increases; down group 17, non-metallic character decreases.
• Metal/non-metal character is a continuum, shown by the basic → amphoteric → acidic trend in oxides.