Chapter IV-2024_384f031ee4fdfc810ac36b1dc147eae4_241210_172101
Chapter IV: Periodic Classification of Elements
IV.1 Electronic Configuration
Definition: The electronic configuration for an atom in its ground state involves the distribution of electrons in its orbitals.
Key Rules for Poly-electronic Atoms:
Klechkowski's Rule: Orbitals are filled according to increasing values of (n + l). If two subshells have the same value of (n + l), the subshell with the lower principal quantum number (n) is filled first.
Pauli's Exclusion Principle: No two electrons in the same atom can have identical quantum numbers.
Hund's Rule: Maximum stability is reached when the number of electrons with parallel spins is maximized in a subshell.
IV.1.1 Klechkowski's Rule
Filling Order Example:
Order of orbitals filled:
3Li : 1s² 2s¹
9F : 1s² 2s² 2p⁵
11Na : 1s² 2s² 2p⁶ 3s¹
22Ti : 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d²
35Br : 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁵
Verification: Ensure the sum of exponents equals the atomic number for each element.
IV.1.2 Pauli's Exclusion Principle
Explanation: Electrons in the same atom cannot have all four quantum numbers identical; therefore, one electron must have a different spin.
IV.1.3 Hund’s Rule
Stability in Subshelss:
Most stable configuration is achieved when there are maximum unpaired electrons with parallel spins before pairing occurs.
IV.1.4 Paired and Unpaired Electrons
Definitions:
Paired: Two electrons in the same quantum box with opposite spins.
Unpaired: A single electron in a quantum box.
IV.2 Particular Electronic Configurations
Notable Exceptions:
Certain elements in the d-block deviate from Klechkowski's rule.
Example:
24Cr : [18Ar] 4s² 3d⁴ (Unstable) → More stable configuration: 24Cr : [18Ar] 4s¹ 3d⁵
29Cu : [18Ar] 4s² 3d⁹ (Unstable) → More stable configuration: 29Cu : [18Ar] 4s¹ 3d¹⁰
IV.3 Electronic Configuration of Ions
Formation of Ions:
Anions form by gaining electrons; Cations by losing them.
Example Configurations:
Na⁺: 1s² 2s² 2p⁶ 3s⁰
S²⁻: 1s² 2s² 2p⁶ 3s² 3p⁶
IV.4 The Periodic Table
Structure: Classified by increasing atomic number (Z).
Key Features:
Elements arranged in periods (rows) and groups (columns).
Elements in the same group share similar electronic configurations.
IV.4.1 Periods and Groups of the Periodic Table
Periods: Correspond to principal quantum numbers. New rows begin with the filling of s orbitals.
Groups: Elements with similar electronic configurations are grouped together.
IV.5 Periodicity of Physical and Chemical Properties
Definition: Elements exhibit periodic trends in properties such as atomic radius, ionization energy, electron affinity, and electronegativity.
IV.5.1 Atomic Radius (r_a)
Definition: Distance from nucleus to limit of the electron cloud.
Trends:
Increases down a group due to added electron layers.
Decreases across a period as effective nuclear charge increases.
IV.5.2 Ionic Radius
Trend: Cation radius < neutral atom radius < Anion radius.
Cations lose electrons, reducing their size; anions gain electrons, increasing size.
IV.5.3 Ionization Energy (E_i)
Definition: Energy required to remove an electron from an atom.
Trends: Increases across a period; decreases down a group.
IV.5.4 Electron Affinity (E_AE)
Definition: Energy change when an electron is added to a neutral atom, generally negative (exothermic).
IV.5.5 Electronegativity (χ)
Definition: Tendency of an atom to attract electrons in a bond.
Trend: Increases across a period, decreases down a group. Fluorine is the most electronegative element.
IV.5.6 Electronegativity Scales
Pauling Scale: Widely used; establishes a relative scale of electronegativity based on bond energies.
Mulliken Scale: Based on average of electron affinity and ionization energy.
Allred-Rochow Scale: Considers the nuclear charge and distance.
Key Configurations Summary
Cobalt: Co : [Ar] 4s² 3d⁷
Carbon: 6C: 1s² 2s² 2p²
Calcium: Ca: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²