The Periodic Table and Mendeleev
How Mendeleev Organised the Elements
Before atomic structure was understood, chemists tried to organise elements by their properties. Dmitri Mendeleev (1834–1907) published his periodic table in 1869. He arranged elements in order of atomic weight and grouped elements with similar chemical properties in the same column. Crucially, he left gaps where no known element fitted the pattern, predicting that new elements would be discovered to fill them. When elements such as gallium and germanium were later discovered and their properties closely matched his predictions, the scientific community accepted his table.
Mendeleev also swapped the positions of some elements so that their chemical properties aligned with their group, even if this meant they were not in strict order of atomic weight. This was a bold move, but it was vindicated when atomic numbers (rather than atomic weights) were later found to be the true basis for ordering the elements.
Mendeleev’s table (1869) organised elements by atomic weight, grouped similar properties in columns, and left gaps for undiscovered elements. The accurate prediction of the properties of those missing elements was key evidence that the table was correct.
The Modern Periodic Table
The modern periodic table arranges all known elements in order of atomic number (number of protons). Elements in the same vertical column share the same group. Elements in the same horizontal row share the same period. Metals are on the left and in the centre of the table; non-metals are on the right.
Electronic configuration determines chemical behaviour. The first electron shell holds up to 2 electrons; the second and third each hold up to 8. You can write the electronic configuration of an atom from its atomic number — for example, sodium (atomic number 11) is written as 2, 8, 1 (2 electrons in shell 1; 8 in shell 2; 1 in shell 3). Elements in the same group have the same number of electrons in their outermost shell, which is why they share similar chemical properties.
Write the electronic configuration of chlorine (atomic number 17). Fill shells in order: shell 1 takes 2 electrons (2 used, 15 remaining); shell 2 takes 8 (10 used, 7 remaining); shell 3 takes the remaining 7. Electronic configuration of chlorine: 2, 8, 7. Because chlorine has 7 outer electrons, it is in Group 7 and behaves as a halogen. From an atom’s electronic configuration, you can identify its group (= number of outer electrons, for groups 1–7) and period (= number of shells).
Group 0 — The Noble Gases
The noble gases (helium, neon, argon, krypton, xenon, radon) have full outer electron shells. This makes them extremely stable and chemically unreactive — they do not normally form compounds. Going down Group 0, the boiling points increase (from −269°C for helium to −62°C for radon) because the atoms become larger, with stronger intermolecular forces between them.
Group 1 — The Alkali Metals
The alkali metals (Li, Na, K, Rb, Cs, Fr) each have one electron in their outer shell. They react with water to produce a metal hydroxide and hydrogen gas — for example: $$2Na + 2H_2O \rightarrow 2NaOH + H_2$$
Going down Group 1, reactivity increases. This is because the outer electron is in a shell further from the nucleus, so the electrostatic attraction between the nucleus and the outer electron is weaker. The outer electron is more easily lost, making the metal more reactive. Going down the group, melting and boiling points decrease, and density generally increases.
Group 1 metals have one outer electron and react vigorously with water. Reactivity increases down the group because the outer electron is further from the nucleus and more easily lost. This is the opposite trend to Group 7.
Group 7 — The Halogens
The halogens (F, Cl, Br, I, At) have seven electrons in their outer shell. They exist as diatomic molecules (e.g. Cl2, Br2). Going down Group 7, reactivity decreases. This is because the outer shell is further from the nucleus, so the atom has a weaker pull on an incoming electron from a reaction partner — it is harder to gain the extra electron needed to complete the outer shell.
Physical state at room temperature changes down the group: fluorine is a pale yellow gas; chlorine is a yellow-green gas; bromine is an orange liquid; iodine is a grey-black solid. Boiling points increase down the group.
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: $$Cl_2 + 2KBr \rightarrow 2KCl + Br_2$$ The solution turns orange as bromine is formed. This reaction confirms the reactivity order: fluorine > chlorine > bromine > iodine. Iodine cannot displace chlorine from a chloride solution because iodine is less reactive than chlorine.
The transition metals occupy the central block of the periodic table. Compared with Group 1 metals they are much less reactive, have higher melting points, are harder and denser, and typically form coloured compounds. Many transition metals and their compounds act as catalysts — for example, iron is used as the catalyst in the industrial Haber process for making ammonia. Transition metals can form ions with different charges (e.g. iron can form Fe2+ and Fe3+).
Group 7 reactivity decreases down the group because the outer shell is further from the nucleus, making it harder for the atom to attract and gain an electron. This is the opposite of Group 1, where losing one electron gets easier down the group.