Classification of Elements and Periodicity in Properties

The One Thing

The modern periodic table arranges elements in increasing atomic number, and recurring patterns in their properties arise chiefly from recurring patterns in electronic configuration. Trends in size, ionisation enthalpy, electron gain enthalpy, electronegativity, metallic character and reactivity can therefore be predicted from an element’s position.

Reactions, Processes and Experiments

Key Terms

• Periodic law: The physical and chemical properties of elements are periodic functions of their atomic numbers.
• Modern periodic table: A table of elements arranged in increasing atomic number, consisting of 7 periods and 18 groups.
• Period: A horizontal row in the periodic table. The period number generally indicates the highest principal shell occupied by electrons.
• Group: A vertical column containing elements with similar valence-shell electronic configurations and related chemical properties.
• Block: A region classified according to the subshell receiving the last electron: s-block, p-block, d-block or f-block.
• Electronic configuration: The arrangement of electrons in shells, subshells and orbitals around the nucleus.
• Aufbau principle: Electrons fill orbitals in order of increasing energy.
• Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers; an orbital holds a maximum of two electrons with opposite spins.
• Hund’s rule: Electrons occupy degenerate orbitals singly with parallel spins before pairing.
• Valence electrons: Electrons in the outermost shell, or sometimes an incompletely filled inner shell, that participate in bonding.
• Effective nuclear charge: The net positive charge experienced by an electron after accounting for shielding by other electrons.
• Shielding effect: The reduction in nuclear attraction experienced by an outer electron because inner-shell electrons repel it.
• Atomic radius: An estimate of atomic size, often half the distance between the nuclei of two bonded identical atoms.
• Covalent radius: Half the distance between the nuclei of two identical atoms joined by a covalent bond.
• Van der Waals radius: Half the distance between the nuclei of two nearest non-bonded atoms in adjacent molecules.
• Ionic radius: The effective size of an ion in an ionic crystal.
• Ionisation enthalpy: The enthalpy required to remove the most loosely held electron from one mole of isolated gaseous atoms or ions.
• Electron gain enthalpy: The enthalpy change when an electron is added to one mole of isolated gaseous atoms to form gaseous anions.
• Electronegativity: The tendency of an atom in a chemical bond to attract the shared electron pair towards itself.
• Metallic character: The tendency of an element to lose electrons and form positive ions.
• Non-metallic character: The tendency of an element to gain or share electrons.
• Diagonal relationship: Similarity between certain diagonally placed elements, such as lithium and magnesium, caused by comparable charge density and size.
• Anomalous behaviour: Unusual properties of the first element of some groups, mainly due to its small size, high ionisation enthalpy, high electronegativity and absence of inner shells.
• Lanthanoid contraction: The gradual decrease in atomic and ionic radii across the lanthanoid series due to poor shielding by 4f electrons.

Easily Confused

• Atomic number and atomic mass: The modern periodic table is based on increasing atomic number, whereas Mendeleev’s periodic law was based on atomic mass.
• Period and group: A period is a horizontal row; a group is a vertical column.
• Covalent radius and van der Waals radius: Covalent radius concerns bonded identical atoms; van der Waals radius concerns nearest non-bonded atoms.
• Cations and anions: Cations are generally smaller than their parent atoms, whereas anions are generally larger.
• Ionisation enthalpy and electron gain enthalpy: Ionisation enthalpy concerns electron removal; electron gain enthalpy concerns electron addition.
• Metallic and non-metallic character: Metallic character involves electron loss, whereas non-metallic character involves electron gain or sharing.
• Electron gain enthalpy and electronegativity: Electron gain enthalpy is an enthalpy change for an isolated gaseous atom; electronegativity is an atom’s attraction for a shared pair in a chemical bond.
• Diagonal relationship and anomalous behaviour: A diagonal relationship describes similarities between diagonally placed elements, whereas anomalous behaviour describes the unusual properties of the first element of a group.
• Effective nuclear charge and shielding effect: Effective nuclear charge is the net attraction experienced by an electron; shielding effect is the reduction of nuclear attraction caused by inner electrons.
• General ionisation-enthalpy trend and its exceptions: Ionisation enthalpy generally increases across a period, but Be is higher than B and N is higher than O.
• Fluorine and chlorine electron gain enthalpy: Chlorine has a more negative electron gain enthalpy than fluorine because its larger 3p orbital causes less electron-electron repulsion.
• Lanthanoid contraction and ordinary atomic-radius trends: Lanthanoid contraction is the specific decrease across the lanthanoid series caused by poor 4f shielding; ordinary period and group trends arise from changes in shells, shielding and nuclear charge.
• Lanthanoids and actinoids: Both are f-block series placed separately at the bottom of the table; they are not identical series.

What Gets Asked

• State or apply a periodic law: Questions may distinguish Dobereiner’s triads, Newlands’ Law of Octaves, Mendeleev’s periodic law and the modern periodic law. The main mark-losing slip is giving atomic mass instead of atomic number as the modern basis.
• Write or identify electronic configurations: Questions may require the Aufbau sequence, shell and subshell capacities, orbital capacity, or the general configurations of the s-, p-, d- and f-blocks. Confusing the last subshell filled with the period or group is a common error.
• Predict periodic trends: Questions may ask how atomic radius, ionisation enthalpy, electron gain enthalpy, electronegativity, metallic character or non-metallic character changes across a period or down a group. The mark-losing slip is reversing the direction of a trend.
• Explain exceptions: Questions may ask why Be has a higher ionisation enthalpy than B, why N has a higher value than O, or why chlorine has a more negative electron gain enthalpy than fluorine. A response that states only the trend without the subshell-energy or electron-repulsion explanation loses the relevant mark.
• Order ions or atoms by size: Questions may use the isoelectronic series O2−, F−, Na+ and Mg2+. The required order is O2− > F− > Na+ > Mg2+, and the relevant explanation is increasing nuclear charge.
• Interpret successive ionisation enthalpies: A large jump indicates that all valence electrons have been removed and the next electron is from an inner shell. Treating every increase as the same kind of jump fails to identify the valence-shell boundary.
• Apply periodic position to chemical behaviour: Questions may ask about valency, bonding, metallic or non-metallic character, oxide character, noble-gas stability, hydrogen’s position, diagonal relationships or lanthanoid contraction. The specific supporting cause—valence configuration, charge density, small size or poor 4f shielding—must be matched to the example.