Learning Objectives
- Understand the modern periodic table and its organization
- Learn about electronic configuration and block classification
- Study periodic trends in properties
- Understand atomic and ionic radii, ionisation enthalpy, electron gain enthalpy, and electronegativity
- Apply periodic trends to predict element properties
Key Concepts
Modern Periodic Law and Table
Modern Periodic Law: Properties of elements are a periodic function of their atomic numbers.
The periodic table has 18 groups (columns) and 7 periods (rows). 118 elements total.
Blocks: s-block (groups 1-2), p-block (groups 13-18), d-block (groups 3-12), f-block (lanthanoids and actinoids).
Electronic Configuration and Position
Period number = number of the outermost shell (highest principal quantum number).
Group number: s-block = number of valence electrons; p-block = 12 + number of valence electrons; d-block = varies.
Periodic Trends
Atomic Radius:
- Across a period (left to right): decreases (increasing nuclear charge pulls electrons closer).
- Down a group: increases (addition of new shells).
- Types: Covalent radius, Van der Waals radius, metallic radius.
Ionic Radius:
- Cation < parent atom (fewer electrons, same nuclear charge).
- Anion > parent atom (more electrons, same nuclear charge).
- Isoelectronic species: more nuclear charge → smaller radius (e.g., O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺).
Ionisation Enthalpy (IE):
- Energy required to remove the outermost electron from a gaseous atom.
- Across a period: generally increases (with exceptions at half-filled/fully-filled subshells).
- Down a group: decreases (electron farther from nucleus, more shielding).
- IE₁ < IE₂ < IE₃ (successive IEs increase).
- Exceptions: IE of Be > B (2s² more stable than 2s²2p¹), IE of N > O (half-filled 2p³ is extra stable).
Electron Gain Enthalpy (EGE):
- Energy change when an electron is added to a gaseous atom.
- Generally becomes more negative across a period (halogens have most negative).
- Becomes less negative down a group (except Cl more negative than F due to small size of F causing electron repulsion).
- Noble gases, alkaline earth metals, and nitrogen have positive EGE.
Electronegativity:
- Tendency of an atom to attract shared electron pairs in a bond.
- Pauling scale: F = 4.0 (most electronegative).
- Increases across a period, decreases down a group.
Other Periodic Properties
Metallic character: Decreases across period, increases down group.
Non-metallic character: Opposite of metallic character.
Oxide nature: Basic → amphoteric → acidic (across a period).
Valence/Oxidation states: Increase from 1 to 7 across a period (groups 1-7), then decrease.
Summary
The modern periodic table organizes elements by atomic number, revealing periodic trends. Atomic radius decreases across periods and increases down groups. Ionisation enthalpy generally increases across periods with exceptions at half/fully-filled subshells. Electron gain enthalpy becomes more negative toward halogens. Electronegativity follows the same trend. These trends help predict chemical properties of elements.
Important Terms
- Periodicity: Repetition of properties at regular intervals of atomic number
- Ionisation Enthalpy: Energy to remove outermost electron from gaseous atom
- Electron Gain Enthalpy: Energy change on adding electron to gaseous atom
- Electronegativity: Ability to attract bonding electrons (Pauling scale)
- Shielding Effect: Reduction in nuclear attraction due to inner electrons
- Effective Nuclear Charge: Z_eff = Z - σ (shielding constant)
Quick Revision
- Across period: radius ↓, IE ↑, EGE more negative, EN ↑
- Down group: radius ↑, IE ↓, EGE less negative, EN ↓
- IE exceptions: Be > B, N > O (half-filled stability)
- EGE: Cl > F (small size of F causes repulsion)
- Isoelectronic: more Z → smaller radius
- Most electronegative: F (4.0); Least IE: Cs/Fr