Classification of Elements and Periodicity in Properties

The One Thing

The modern periodic table arranges elements by increasing atomic number, so recurring valence-shell configurations produce recurring chemical properties. Periodic trends are explained mainly by effective nuclear charge, shielding effect, electron-shell number, and subshell arrangement.

Reactions, Processes and Experiments

Key Terms

• Modern periodic law: The physical and chemical properties of elements are periodic functions of their atomic numbers.
• Atomic number: The number of protons in the nucleus; it uniquely identifies an element and determines its position in the modern periodic table.
• Periodic table: An arrangement of elements in seven periods and eighteen groups based on increasing atomic number and similar electronic configurations.
• Period: A horizontal row in the periodic table. Its number generally indicates the highest principal quantum number or number of occupied electron shells.
• Group: A vertical column whose elements usually have similar valence-electron configurations and related chemical properties.
• Block: A section classified according to the subshell receiving the differentiating electron: s, p, d, or f.
• s-block elements: Elements in Groups 1 and 2 whose differentiating electron enters an s-subshell; they are generally reactive metals.
• p-block elements: Elements in Groups 13 to 18 whose differentiating electron enters a p-subshell; the block includes metals, metalloids, nonmetals, halogens, and noble gases.
• d-block elements: Elements in Groups 3 to 12 whose differentiating electron enters a d-subshell; they are commonly called transition elements.
• f-block elements: The lanthanoid and actinoid series, whose differentiating electron enters an f-subshell.
• Valence electrons: Electrons in the outermost shell, or electrons involved in bonding, that largely determine reactivity and valency.
• Effective nuclear charge: The net positive attraction experienced by an electron after shielding by other electrons has been considered.
• Shielding effect: The reduction in nuclear attraction experienced by an outer electron because inner-shell electrons repel it.
• Atomic radius: A measure of atomic size, commonly estimated as half the distance between the nuclei of two bonded identical atoms.
• Covalent radius: Half the internuclear distance between two identical atoms joined by a single covalent bond.
• Van der Waals radius: Half the distance between the nuclei of two identical, non-bonded atoms in contact.
• Ionic radius: The effective size of an ion in a crystal; cations are usually smaller and anions usually larger than their parent atoms.
• Ionization enthalpy: The enthalpy required to remove the most loosely held electron from one mole of isolated gaseous atoms to form gaseous cations.
• First ionization enthalpy: The energy required for $X(g)\to X^{+}(g)+e^{-}$.
• Successive ionization enthalpies: The energies required to remove additional electrons; they increase because each electron is removed from an increasingly positive ion.
• Electron gain enthalpy: The enthalpy change when an electron is added to one mole of isolated gaseous atoms to form gaseous anions.
• Electronegativity: The relative tendency of an atom in a chemical bond to attract the shared electron pair toward itself.
• Metallic character: The tendency to lose electrons and form positive ions.
• Non-metallic character: The tendency to gain or share electrons.
• Chemical reactivity: The tendency of an element to undergo chemical change.
• Diagonal relationship: Similarity between certain diagonally positioned elements, especially Li and Mg, Be and Al, and B and Si.
• Anomalous behaviour: The unusual properties of the first element of some groups, mainly due to its small size, high ionization enthalpy, high electronegativity, and absence of vacant d-orbitals.
• Periodic trends: Regular changes in elemental properties across periods and down groups caused mainly by changes in nuclear charge, shells, shielding, and electronic configuration.

Easily Confused

• Atomic number vs atomic mass: The modern periodic table is arranged by increasing atomic number, whereas Mendeleev’s table was mainly arranged by increasing atomic mass.
• Period vs group: A period is a horizontal row; a group is a vertical column.
• Covalent radius vs van der Waals radius: Covalent radius concerns bonded identical atoms; van der Waals radius concerns identical non-bonded atoms in contact.
• Ionization enthalpy vs electron gain enthalpy: Ionization enthalpy concerns removing an electron; electron gain enthalpy concerns adding an electron.
• Metallic vs non-metallic character: Metallic character describes electron loss, whereas non-metallic character describes electron gain or sharing.
• Electron gain enthalpy of fluorine vs chlorine: Chlorine has the more negative value because its larger 3p subshell produces less electron-electron repulsion for the added electron.
• Hydrogen vs Group 1 placement: Hydrogen has one valence electron like Group 1 elements but also resembles Group 17 elements because it needs one electron for a stable shell.
• Helium vs the general Group 18 configuration: Helium is $1s^2$, not $n{s}^{2}n{p}^6$, but it is placed in Group 18 because its shell is complete and it has noble-gas characteristics.
• Lanthanoids vs actinoids: Lanthanoids belong to Period 6; actinoids belong to Period 7.
• Across-period and down-group trends: Across a period, effective nuclear charge generally increases; down a group, electron-shell number and shielding increase.

What Gets Asked

• Classification history: Questions may require the sequence Dobereiner’s triads in 1829, Newlands’ law of octaves in 1866, Mendeleev’s table in 1869, and Moseley’s work in 1913. A common mark-losing error is stating that the modern table is arranged by atomic mass rather than atomic number.
• Table structure and configurations: Questions may ask for the numbers of periods, groups, and blocks, the capacities of periods, or the general configurations $n{s}^{1-2}$, $n{s}^{2}n{p}^{1-6}$, $(n-1)d^{1-10}n{s}^{0-2}$, and $(n-2)f^{1-14}(n-1)d^{0-1}n{s}^2$. Marks are lost by confusing a block with a period or group.
• Periodic trends: Questions may ask for trends in atomic radius, ionization enthalpy, electron gain enthalpy, electronegativity, metallic character, non-metallic character, or oxide character. The relevant distinction is that atomic radius generally decreases across a period but increases down a group, whereas metallic character follows the opposite broad pattern.
• Trend exceptions: Questions may test why Be has a higher first ionization enthalpy than B, why N has a higher first ionization enthalpy than O, or why chlorine has a more negative electron gain enthalpy than fluorine. Marks are lost when these are treated as violations without reference to subshell energy or electron-electron repulsion.
• Special placements: Questions may ask about hydrogen’s unique position, helium’s placement in Group 18, the separate placement of the f-block, or diagonal relationships involving Li and Mg, Be and Al, and B and Si. The specific electronic configurations and stated reasons must be included.
• Electron configuration principles: Questions may require use of $2n^2$, $2(2l+1)$, the Aufbau principle, Hund’s rule, or the Pauli exclusion principle. A frequent error is assigning the wrong capacity to s, p, d, or f subshells.