Periodicity — OCR A-Level Chemistry

This guided-reading row introduces the layout of the periodic table as a framework for later study of periodicity. Read the specification section and your textbook together, and learn how elements are arranged in order of increasing atomic number into periods (horizontal rows) and groups (vertical columns). Note that elements in the same group have the same number of outer-shell electrons and therefore similar chemical properties, while elements across a period show a gradual change in properties. Identify the s-, p- and d-blocks and the position of metals, non-metals and the noble gases. Build a labelled diagram and practise locating elements by group and period so you can use the table confidently when explaining trends.

This row introduces the periodic table as a way of arranging the elements, and it is the frame for the sub-points that follow, including ordering by increasing atomic (proton) number. Read it as a map: elements sit in periods (horizontal rows) and groups (vertical columns), and their position links to proton number and to patterns in physical and chemical properties. When you study, practise reading positions off a full table, noting the group number and period of an element, and connecting position to electron configuration and to trends such as ionisation energy across a period. Treat the statement as the organising idea for the section rather than a single fact to memorise, and check each later sub-point against this arrangement.

This sub-point states the ordering principle of the modern periodic table: elements are arranged by increasing atomic (proton) number, that is, the number of protons in the nucleus. Atomic number is the same as proton number, so an element with 11 protons (sodium) comes after one with 10 (neon) and before one with 12 (magnesium). This replaced older orderings based on relative atomic mass, which placed some elements out of sequence. In a multiple-choice question you may be asked which quantity increases across the table, or to identify the correct ordering principle. Read each option carefully: proton number, nucleon number and relative atomic mass are different quantities, and only proton number gives the modern sequence.

Periodicity is the repeating pattern in physical and chemical properties as you move across a period and then down to the next period. Across Period 2 or Period 3, nuclear charge rises while the outer electrons stay in the same shell, so atomic radius falls, first ionisation energy generally rises (with dips at Al and S in Period 3), and electronegativity increases. Metallic character falls and non-metallic character rises, so the oxides change from basic (Na₂O, MgO) through amphoteric (Al₂O₃) to acidic (SO₂, SO₃, and the oxides of P, S and Cl). Melting points trace the structure: giant ionic and giant covalent solids are high, simple molecular solids low. The pattern repeats in the next period because the same outer-shell configuration recurs.

Elements in the same group share the number of outer-shell electrons, so they show similar chemical properties that change gradually down the group. In Group 2, reactivity increases down the group as the first and second ionisation energies fall and the atoms become larger, so the metals react more vigorously with water and oxygen. In Group 7, reactivity decreases down the group because the outer shell is further from the nucleus and more shielded, so oxidising power falls and displacement reactions show a more reactive halogen displacing a less reactive one from its salt. Down Group 2 the hydroxides become more soluble and the thermal stability of carbonates and nitrates increases; down Group 7 the boiling points rise as the molecules get larger and van der Waals forces strengthen.

This row is a study guide, not an assessed answer target. Read it as a map of how Period 2 and Period 3 electron configurations link to the trend in first ionisation energy. Work across each period from Group 1 to Group 18 and write the configuration of each element, for example Na as 1s² 2s² 2p⁶ 3s¹ and Mg as 1s² 2s² 2p⁶ 3s². Note that the outer electron is removed from the highest-energy occupied shell, so the shell number and subshell matter. Then explain the general rise across a period by the increasing nuclear charge and the similar shielding, which pulls the outer electron in more tightly. Treat the dips at Group 3 and Group 6 as the interesting cases and prepare to explain each with subshell energies and electron pairing. Use the specification cross-reference to 2.2.1 d to revise subshell energies and spin-pair repulsion before you attempt the trend.

This criterion asks you to describe how electron configurations change as you move across Period 2 (Li to Ne) and Period 3 (Na to Ar). The key pattern is that each element adds one proton and one electron, and the extra electron enters the same outer shell until that shell is full. In Period 2 the 2s subshell fills at Be and the 2p subshell fills from B to Ne; in Period 3 the 3s subshell fills at Mg and the 3p subshell fills from Al to Ar. Write configurations in subshell notation, for example Mg as 1s² 2s² 2p⁶ 3s² and Cl as 1s² 2s² 2p⁶ 3s² 3p⁵. Note that the 4s subshell is not occupied in these periods, so K and Ca lie beyond the trend described here. Use the cross-reference to 2.2.1 d to check subshell energies and the order in which orbitals fill.

Block classification depends on which subshell receives the differentiating electron in the ground-state atom. An element belongs to the s-block when its highest-energy electron enters an s subshell, to the p-block when it enters a p subshell, and to the d-block when it enters a d subshell. So hydrogen and helium sit in the s-block, groups 13 to 18 form the p-block, and the transition metals plus scandium and zinc occupy the d-block. Learn to read a configuration such as 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p¹ and identify the last-filled subshell as p, placing the element in the p-block. Remember that helium is s-block despite sitting above neon, and that zinc is d-block even though it forms no partly filled d subshell in its common ion.

First ionisation energy is the energy needed to remove 1 mol of electrons from 1 mol of gaseous atoms, forming 1 mol of gaseous 1⁺ ions. The definition requires gaseous atoms and gaseous ions, so a solid or liquid sample would give a different value. Successive ionisation energies measure removal of the second, third and later electrons from the same gaseous species, each step costing more because the remaining ion holds the electrons more tightly. Large jumps between successive values reveal the number of electrons in the outer shell, so a jump after the second value indicates two outer electrons. Trends across periods and down groups follow from nuclear charge, shielding and atomic radius.

First ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms, forming one mole of gaseous 1⁺ ions. Across Periods 2 and 3, the nuclear charge rises by one proton per element while the outer electrons enter the same shell, so atomic radius falls and the outer electron is held more strongly; first ionisation energy therefore rises overall. Explain the small dips at Group 3 (outer electron in a p orbital, slightly higher in energy and shielded by the s electrons) and at Group 6 (two electrons paired in one p orbital, so electron–electron repulsion eases removal). Down a group, the outer electron occupies a shell further from the nucleus and is shielded by more inner shells, so atomic radius increases and attraction weakens; first ionisation energy falls. Link each trend explicitly to attraction, nuclear charge and atomic radius.

Successive ionisation energies are the energies needed to remove electrons one at a time from a gaseous atom. Each removal takes more energy than the last because the remaining electrons are held by a progressively larger effective nuclear charge relative to fewer electrons. A large jump in energy marks the start of a new shell: the electrons removed before the jump were in the outer shell, and the jump occurs when an electron is taken from a shell closer to the nucleus. Count the electrons removed before the first big jump to find the number in the outermost shell, which equals the group number for main-group elements. Repeat for each subsequent jump to assign electrons to each shell. For example, three electrons removed before a large jump indicate three outer electrons and Group 3.

Across Period 2 and Period 3, the bonding and structure of the elements change from metallic to giant covalent to simple molecular. Sodium, magnesium and aluminium form giant metallic lattices; silicon forms a giant covalent (macromolecular) lattice; phosphorus, sulfur, chlorine and argon exist as simple molecular or monatomic substances. Melting point reflects the forces broken on melting: strong metallic bonds in the metals, many strong covalent bonds in silicon, and only weak intermolecular forces between simple molecules. Melting point therefore rises from sodium to magnesium to aluminium, peaks sharply at silicon, then falls abruptly to phosphorus and sulfur, with chlorine and argon lower still. A question may ask you to identify an element from a melting-point value, explain a rise or fall, or predict the structure from the trend.

This row is a course-guidance entry rather than a standalone assessed statement: the published heading 3.1.1(d) introduces an explanation topic whose sub-clauses carry the actual content, so read the sub-clauses beneath it before revising. Work through the specification in order, note each sub-clause the heading introduces, and check that your notes explain the structure and bonding of the relevant element or compound rather than merely describing it. Use the official specification and your centre's scheme of work to confirm exactly which sub-clauses fall under this heading, and treat any summary that assigns different content to 3.1.1(d) as unreliable. Build explanation-style answers that link structure to properties, and test yourself by writing short explanations for each sub-clause.

Metallic bonding is modelled as the strong electrostatic attraction between cations (positive ions) arranged in a regular lattice and a sea of delocalised electrons that are free to move throughout the structure. Each metal atom loses its outer electrons to become a cation, and those released electrons are shared across the whole lattice rather than belonging to any one ion. The attraction acts in all directions, holding the lattice together. Because the electrons are delocalised, they can carry charge when a potential difference is applied, which explains why metals conduct electricity in the solid state. The strength of the attraction rises with the charge on the cation and the number of delocalised electrons per atom, so magnesium (2⁺ ions, two electrons per atom) is stronger than sodium (1⁺ ions, one electron per atom).

A giant metallic lattice is a three-dimensional regular arrangement of metal cations surrounded by delocalised electrons, extending throughout the whole crystal with no discrete molecules. Every metal, from sodium to iron, adopts some form of this giant lattice, so the structure is described as giant because it has no finite size and as a lattice because the ions sit in a repeating pattern. The regular packing means the layers of cations can slide over one another when a force is applied, because the delocalised electrons continue to hold the ions together in their new positions. This explains why metals are malleable and ductile rather than brittle. The same delocalised electrons allow conduction of electricity and heat. Contrast this with a simple molecular substance, which consists of small discrete molecules held by weak intermolecular forces.

A giant covalent lattice is a continuous three-dimensional network of atoms joined by strong covalent bonds, with no discrete molecules. In diamond each carbon atom bonds tetrahedrally to four others, giving a rigid three-dimensional network. In graphite each carbon atom bonds to three others in a hexagonal layer, and the layers are held by weak intermolecular forces, with delocalised electrons between layers. Graphene is a single layer of graphite, one atom thick. Silicon forms a similar three-dimensional network to diamond, with each silicon atom covalently bonded to four neighbours. The strong covalent bonds, not intermolecular forces, explain the very high melting points of these solids.

Giant metallic lattices consist of positive metal ions in a regular arrangement surrounded by delocalised electrons. Giant covalent lattices are networks of atoms joined by strong covalent bonds. Melting and boiling points are high in both because strong attractions must be overcome: metallic bonds in metals, and covalent bonds in giant covalent solids. Metals generally conduct electricity because delocalised electrons move through the lattice; most giant covalent solids do not conduct, except graphite and graphene, which have delocalised electrons between or along layers. Metals are generally insoluble in water because the metallic bonding is strong, while giant covalent solids are also insoluble because their covalent networks are not broken by water.

Explain the melting-point trend across Periods 2 and 3 by linking each element's structure and bonding to the energy needed to overcome its forces. Sodium, magnesium and aluminium form giant metallic lattices; melting point rises from Na to Al as ionic charge and the number of delocalised electrons per atom increase, strengthening metallic bonding. Silicon (and boron in Period 2) is giant covalent, with many strong covalent bonds to break, giving very high melting points. Phosphorus (P₄), sulfur (S₈), chlorine (Cl₂) and argon are simple molecular or monatomic, held by weak London dispersion forces; melting points are low and depend on molecular size, so S₈ melts higher than P₄ and Cl₂. Argon is monatomic with only weak dispersion forces, giving the lowest melting point.