Classification of Elements and Periodicity

By the early 1800s chemists had isolated dozens of elements and needed a way to organise them. The breakthrough was noticing that properties repeat. Döbereiner (1817) spotted groups of three — triads — in which the middle element's atomic mass is roughly the average of the other two. Newlands (1865) arranged elements by increasing mass and found that every eighth element resembled the first — his musical law of octaves.

In 1869 Dmitri Mendeleev arranged the 63 known elements by increasing atomic mass into rows and columns, stating his periodic law: the properties of elements are a periodic function of their atomic masses. His genius lay in two bold moves — leaving gaps for undiscovered elements and even reversing a few mass orders to keep similar elements together.

In 1913 Henry Moseley measured the frequencies of X-rays emitted by elements and found they depended on atomic number \(Z\) — the nuclear charge — not atomic mass. This resolved every anomaly. The modern periodic law states:

The modern long-form table has 7 periods (horizontal rows) and 18 groups (vertical columns). A period number equals the principal quantum number \(n\) of the outermost shell being filled; elements in a group share the same valence-shell configuration, hence similar chemistry. By the subshell whose filling defines them, elements fall into four blocks.

For elements beyond \(Z=100\), before an official name is agreed, IUPAC assigns a systematic name built from the digits of the atomic number using numerical roots, ending in -ium.

An atom has no sharp edge, so "size" is defined operationally — as the covalent, metallic or van der Waals radius. Across a period, size decreases: electrons enter the same shell while nuclear charge rises, so the effective pull (\(Z_{\text{eff}}\)) tightens the cloud. Down a group, size increases: each new period adds a shell.

Ions differ from their parent atoms. A cation is always smaller (it loses a shell or feels more pull per electron); an anion is always larger (added electrons increase repulsion). For isoelectronic species — same electron count — size falls as nuclear charge rises.

The ionization enthalpy (IE) is the energy needed to remove the most loosely held electron from one mole of gaseous atoms. It rises across a period (greater \(Z_{\text{eff}}\), smaller size) and falls down a group (larger size, more shielding). Each successive IE is larger than the last because removing an electron from a more positive ion is harder.

The electron gain enthalpy (\(\Delta_{eg}H\)) is the energy change when an electron is added to a gaseous atom. A large negative value means the atom readily accepts an electron. It generally becomes more negative across a period and less negative down a group — with one celebrated twist.

Electronegativity (EN) is the tendency of an atom in a bond to attract the shared electron pair. Unlike IE or \(\Delta_{eg}H\), it is a relative property with no fixed units. On the Pauling scale, fluorine is the most electronegative element at \(4.0\). EN rises across a period and falls down a group — tracking atomic size and \(Z_{\text{eff}}\).

Valency repeats periodically: for main-group elements it equals the group number (for metals) or \(8\) minus the group number (for non-metals). Metallic character decreases across a period and increases down a group, and this is mirrored by the oxides: metal oxides are basic, non-metal oxides acidic, with amphoteric oxides in between.

Two finishing patterns: the diagonal relationship links an element to the one diagonally below-right (\(\ce{Li}\)–\(\ce{Mg}\), \(\ce{Be}\)–\(\ce{Al}\), \(\ce{B}\)–\(\ce{Si}\)) because falling charge density down a group is offset by rising charge density across it. And the first member of each main group (\(\ce{Li, Be, B, C, N, O, F}\)) behaves anomalously — small size, high EN, and no available \(d\)-orbitals set it apart from the rest of its family.

For each item, first decide which idea it tests — history, position, or a periodic trend — then apply the relevant rule. Difficulty rises down the list.