The Periodic Table and Its Patterns
A System That Predicts Behaviour
The periodic table is one of the most powerful tools in chemistry. It does not just list the elements — it organises them so that patterns in their properties become visible and predictable. A chemist encountering an unfamiliar element for the first time can look at its position in the table and make reasonable predictions about its reactivity, the type of bonds it is likely to form, whether it is a metal or non-metal, and even what formula its compounds might take. Understanding why those predictions work requires knowing how electrons are arranged in atoms: the structure of the outer electron shell is the key to an element’s chemical behaviour.
How the Periodic Table Is Organised
Elements are arranged in order of increasing atomic number (number of protons). The table is divided into horizontal rows called periods and vertical columns called groups. Elements in the same group have the same number of electrons in their outer shell, which is why they share similar chemical properties.
The periodic table organises elements by atomic number. Periods (rows) share the same number of electron shells. Groups (columns) share the same number of outer electrons and therefore have similar chemical properties. Metals appear on the left; non-metals on the right.
Mendeleev and the Predictions of the Table
Mendeleev published his periodic table in 1869, arranging elements by atomic weight (not atomic number, as we use today) but grouping elements with similar properties in the same columns. Where no known element fitted, he left a deliberate gap and predicted the specific properties of the missing element — including its expected density, melting point, and the formula of its oxide. When gallium was discovered in 1875 and germanium in 1886, their measured properties closely matched his predictions. This predictive success led to wide scientific acceptance of his table and is a powerful example of how a scientific model gains credibility through making accurate, testable predictions about phenomena not yet observed.
Group 1 — The Alkali Metals
The alkali metals (Group 1: lithium, sodium, potassium, and below) each have one electron in their outer shell. They are soft metals with relatively low melting points, and they react vigorously with water to produce a metal hydroxide and hydrogen gas:
$$2Na + 2H_2O \rightarrow 2NaOH + H_2$$
Going down Group 1, reactivity increases. This is because the outer electron is in a shell further from the nucleus and is more shielded from the nuclear attraction by inner electron shells. It is therefore more easily lost. The elements also become denser and have lower melting points going down the group.
Group 7 — The Halogens
The halogens (Group 7: fluorine, chlorine, bromine, iodine) each have seven electrons in their outer shell. They exist as diatomic molecules (\(Cl_2\), \(Br_2\), \(I_2\)) and are highly reactive non-metals that react with metals to form ionic salts called halides.
Going down Group 7, reactivity decreases. The outer shell is further from the nucleus and more shielded, so it is harder to attract an extra electron to complete the shell. A more reactive halogen can displace a less reactive halogen from a solution of its salt: for example, chlorine water added to potassium bromide solution turns the solution orange, as bromine is released.
Group 1: reactivity increases down the group (outer electron further from nucleus, more shielded, more easily lost). Group 7: reactivity decreases down the group (outer shell further from nucleus, harder to gain an extra electron). These trends are opposite because metals lose electrons while non-metals gain them.
The noble gases (Group 0: helium, neon, argon, krypton, xenon) have completely full outer electron shells. This makes them very stable and extremely unreactive — they rarely form compounds under normal conditions. Boiling points increase going down the group. Noble gases are used where unreactive conditions are needed: argon in welding (shields the weld from oxygen), helium in balloons and airships, neon in lighting.
The transition metals occupy the central block of the periodic table. Compared with Group 1 alkali metals, they are much less reactive, harder, denser, have higher melting points, and are much better at acting as catalysts. A key distinguishing feature is that transition metals form coloured compounds (copper compounds are blue or green; iron(II) compounds are pale green; iron(III) compounds are orange-brown) and can often form ions with more than one charge (e.g. iron as Fe²⁺ or Fe³⁺).
Familiar examples: iron (Fe) as a catalyst in the Haber process for ammonia; copper (Cu) as an electrical conductor; titanium (Ti) in aircraft and biomedical implants; nickel (Ni) as the catalyst for hardening vegetable oils. Transition metals are useful precisely because they are less reactive than Group 1 metals, yet still able to participate in chemical reactions as catalysts.