Short Notes: Classification and Periodicity

3.1 Modern Periodic Law

Modern Periodic Law: Properties of elements are periodic functions of their atomic numbers. This statement expresses that when elements are arranged in order of increasing atomic number, a periodic repetition of properties is observed.

• Historical context: Dmitri Mendeleev arranged elements by increasing atomic mass and predicted missing elements; Henry Moseley showed that atomic number (Z) - not atomic mass - is the fundamental ordering principle, leading to the modern periodic law.
• Periodic table structure: There are 18 groups (vertical columns) and 7 periods (horizontal rows).
• Blocks of the periodic table: Elements are grouped into blocks according to their outermost electron subshell being filled:
  • s-block: Groups 1-2
  • p-block: Groups 13-18
  • d-block: Groups 3-12 (transition elements)
  • f-block: Lanthanides and actinides (placed separately at the bottom)
• Basis of periodicity: Periodic repetition arises from the recurring pattern of valence electronic configuration as atomic number increases.
• Group similarity: Elements in the same group have similar chemical properties because they have the same number of valence electrons and similar valence-shell configurations.

3.2 Periodic Trends

Trends in properties across a period and down a group can be explained using two key concepts: effective nuclear charge (Z_eff) felt by valence electrons and the shielding effect of inner electrons. Penetration of orbitals and subshell energies also influence anomalies.

• Atomic radius (size of atom): Generally decreases across a period (left → right) due to increasing Z_eff which pulls electrons closer; increases down a group due to addition of electron shells.
• Ionic radius: Cations are smaller than their parent atoms (loss of an electron reduces electron-electron repulsion and may remove an outer shell); anions are larger (gain of electron increases repulsion). Across a period, radii of isoelectronic species decrease with increasing nuclear charge.
• Ionisation enthalpy (first ionisation energy, IE1): Energy required to remove one electron from an isolated gaseous atom or ion. IE1 generally increases across a period (higher Z_eff) and decreases down a group (outer electron further from nucleus and more shielded).
• Electron affinity (EA): Energy change when an electron is added to an isolated gaseous atom. EA generally becomes more negative across a period (atoms more readily accept electrons) and varies down a group (often becomes less negative).
• Electronegativity: A measure of an atom's ability to attract shared electrons in a bond. Electronegativity increases across a period and decreases down a group.
• Metallic and non-metallic character: Metallic character decreases across a period and increases down a group. Non-metallic character shows the opposite trend.
• Oxidation states and valency: s- and p-block elements show predictable valencies based on valence electrons; d-block (transition) and f-block elements show variable oxidation states due to participation of (n-1)d and nf electrons.
• Reactivity trends: Alkali metals (Group 1) become more reactive down the group; halogen reactivity as oxidising agents generally decreases down the group (F₂ > Cl₂ > Br₂ > I₂).
• Anomalies: Small deviations occur due to subshell configurations and electron pairing (for example, the ionisation enthalpy of boron is lower than that of beryllium; oxygen shows a smaller electron affinity than nitrogen in some cases).

Explanations-key points:

• Effective nuclear charge (Z_eff): Net positive charge experienced by valence electrons after shielding by inner electrons; larger Z_eff pulls electrons closer and raises ionisation energies.
• Shielding effect: Inner shell electrons reduce attraction between nucleus and outer electrons; shielding increases down a group, reducing Z_eff on outermost electrons.
• Orbital penetration and pairing: s orbitals penetrate closer to nucleus than p, d, f and so influence trends; electron pairing in orbitals causes extra repulsion and explains certain anomalies.

Representative examples:

• Across Period 3: atomic radius decreases from Na → Ar; ionisation enthalpy increases; metallic character decreases.
• Down Group 1 (alkali metals): atomic radius increases from Li → Cs; IE1 decreases; reactivity increases.
• Isoelectronic species: for example, O²⁻ < F⁻ < Na⁺ < Mg²⁺ in size when compared with same number of electrons but increasing nuclear charge leads to decreasing ionic radius.

3.3 Key Definitions

• Atomic number (Z): Number of protons in the nucleus of an atom; uniquely identifies an element.
• Mass number (A): Total number of protons and neutrons in the nucleus.
• Isotopes: Atoms of the same element (same Z) having different mass numbers (different numbers of neutrons).
• Isobars: Nuclei of different elements having the same mass number (A) but different Z.
• Isotones: Nuclei of different elements having the same number of neutrons.
• Electronic configuration: Distribution of electrons among atomic orbitals following Aufbau principle, Pauli exclusion principle and Hund's rule; determines valence and chemical behaviour.
• Valence electrons: Electrons in the outermost shell that participate in chemical bonding.
• Valency: Combining capacity of an atom, usually equal to the number of electrons lost, gained or shared to attain a noble gas configuration.
• Oxidation state (oxidation number): Formal charge on an atom in a compound assuming ionic distribution of electrons.
• Atomic radius (types): Covalent radius, metallic radius and van der Waals radius - different definitions depending on bonding and measurement method.
• Ionic radius: Effective radius of an ion in an ionic solid; depends on coordination number and charge.
• Ionisation enthalpy (IE): Minimum energy required to remove an electron from a gaseous atom or ion.
• Electron affinity (EA): Energy change when an electron is added to a gaseous atom.
• Electronegativity: Relative tendency of an atom to attract shared electrons; commonly given on the Pauling scale.
• Effective nuclear charge (Z_eff): Net positive charge experienced by valence electrons after accounting for shielding by inner electrons.
• Shielding (screening): Reduction in effective nuclear attraction on valence electrons due to inner electrons.

Summary: The modern periodic law organises elements by increasing atomic number, producing a table in which chemical and physical properties repeat periodically. Understanding electronic configuration, effective nuclear charge and shielding explains the major periodic trends - atomic and ionic sizes, ionisation enthalpy, electron affinity, electronegativity and metallic character - which are essential for predicting element behaviour and chemical reactivity.