Periodic Table and Group Trends
The Development of the Periodic Table
Scope note: the trends in Groups 1, 7 and 0 are topic C4.1, examined on Chemistry Paper 2 (F2 / H4). They sit here because they follow on from the periodic table.
In the nineteenth century, scientists searched for a way to organise the growing number of known elements. Dmitri Mendeleev produced the most successful early arrangement around 1869. He ordered elements by atomic weight but made two crucial decisions that others had not: he left deliberate gaps for elements that had not yet been discovered, and he swapped the order of a few elements so that those with similar chemical properties lined up in the same vertical columns. When subsequently discovered elements — including gallium and germanium — matched the properties he had predicted for the gaps, his table gained wide acceptance. The modern periodic table arranges elements in order of atomic number (proton number) rather than atomic weight, which resolved the few anomalies in Mendeleev’s arrangement.
Elements in the periodic table are arranged in order of atomic number. Elements in the same group (column) have the same number of outer electrons and similar chemical properties. Elements in the same period (row) have the same number of electron shells. Metals are on the left and centre; non-metals are on the right.
Group 1 — The Alkali Metals
The alkali metals (lithium, sodium, potassium, rubidium, caesium and francium) all have one electron in their outer shell. They are soft, low-density metals with relatively low melting points compared to most metals. All react vigorously with water, producing a metal hydroxide solution and hydrogen gas. The general equation for the reaction with water is:
$$\text{metal} + H_2O \rightarrow \text{metal hydroxide} + H_2$$
For example, sodium reacts with water according to:
$$2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)$$
Going down Group 1, reactivity increases. This is because the outer electron is in a shell that is increasingly far from the nucleus, and so is more weakly attracted by the positive nuclear charge. The outer electron is therefore more easily lost. Potassium reacts faster and more vigorously with water than sodium, which in turn reacts faster than lithium. Melting points decrease and densities generally increase going down the group.
Flame tests allow you to identify metal ions by the colour they produce when heated in a Bunsen flame. A clean platinum or nichrome wire loop is dipped into the sample and held in the hottest part of the flame. Group 1 metal flame colours: lithium → crimson, sodium → yellow, potassium → lilac. This is a practical activity group (PAG C5) context — you should be able to describe the method, identify metals from flame colours, and explain why each loop must be cleaned between tests (to avoid contamination from previous samples).
Group 7 — The Halogens
The halogens (fluorine, chlorine, bromine, iodine and astatine) have seven electrons in their outer shell and exist as diatomic molecules (\(F_2\), \(Cl_2\), \(Br_2\), \(I_2\)). They need to gain just one electron to achieve a full outer shell, which makes them very reactive non-metals. At room temperature, chlorine is a pale yellow-green gas, bromine is a red-brown liquid, and iodine is a dark grey solid. Going down Group 7, the boiling point increases as the molecules get larger.
Going down Group 7, reactivity decreases. This is because the outer shell is further from the nucleus in atoms lower in the group, making it harder for the nucleus to attract an extra electron. A more reactive halogen can displace a less reactive halogen from a solution of its salt. For example, chlorine displaces bromine from potassium bromide solution:
$$Cl_2(aq) + 2KBr(aq) \rightarrow 2KCl(aq) + Br_2(aq)$$
If chlorine is added to potassium iodide solution, a brown colour develops as iodine is displaced. If bromine is added to potassium chloride solution, no reaction occurs because bromine is less reactive than chlorine.
Group 1: reactivity increases going down (outer electron lost more easily as it is further from the nucleus). Group 7: reactivity decreases going down (harder to gain an electron as the outer shell is further from the nucleus). These trends are opposite because metals lose electrons while non-metals gain them.
Group 0 — The Noble Gases
The noble gases (helium, neon, argon, krypton, xenon and radon) have completely full outer electron shells. This makes them very stable and extremely unreactive — they do not normally form compounds. They exist as single atoms (monatomic) rather than as molecules. Their boiling points increase going down the group as the atoms get larger and intermolecular forces increase. Noble gases have many practical applications: helium in weather balloons and airships (less dense than air, non-flammable), neon in display lighting (glows a distinctive red-orange), and argon in tungsten filament light bulbs and in welding (inert atmosphere prevents oxidation).
An element’s position in the periodic table lets you predict its reactions. Group number tells you how many outer electrons the atom has, which tells you how it bonds. Period number tells you how many electron shells it has, and going down a group the outer shell moves further from the nucleus. Mendeleev used exactly this logic to predict the properties of undiscovered elements — he predicted the density, melting point and typical compounds of elements such as gallium before they were discovered. The predictions were confirmed when the elements were isolated, which powerfully validated the periodic table as a predictive tool, not just a classification system.