Periodic Table

Classification of Elements Prior to Atomic Structure Knowledge:
- Initially, about 30 elements known in the early nineteenth century were classified into three categories based on physical properties: metals, nonmetals, and metalloids.
- Classification methods included:
- Dobereiner's Triads: Grouping of elements in threes exhibiting similar properties.
- Newlands’ Octaves: An arrangement where every eighth element shared similar properties based on atomic weights.
Key Scientists in Classification:
- Dmitri Mendeleev (1869): Formulated the periodic table based on atomic masses and properties of 63 known elements at that time.
- Henry Moseley (1913): Established that atomic number (Z) is a more fundamental property than atomic mass for arranging elements.

Definition:
- Mendeleev's periodic law states: "The physical and chemical properties of elements are periodic functions of their atomic masses."
Arrangement:
- Elements were arranged in increasing order based on atomic masses. Each element's position reflected its recurring chemical and physical properties.
Gaps: Mendeleev left gaps in the table for undiscovered elements, predicting their properties, which were later confirmed upon their discovery.

Following Moseley’s work, the periodic law underwent modification:
- Modern Periodic Law: "The physical and chemical properties of elements are periodic functions of their atomic numbers."
Revisions to the Table:
- The modern periodic table is a revised long form of Mendeleev's table incorporating noble gases and accounting for new elemental discoveries.

Cells:
- The modern periodic table features horizontal rows (periods) and vertical columns (groups).
Number of Periods and Groups:
- There are 7 periods (numbered 1 to 7) and 18 groups (numbered 1 to 18).
Total Elements:
- The table has 118 boxes representing all known elements, recognized by IUPAC.
Blocks of the Table:
- The table is divided into four blocks based on electron configurations:
- s-block (Groups 1 & 2): Elements with the last electron in an s subshell.
- p-block (Groups 13 to 18): Elements whose last electron enters a p subshell.
- d-block (Transition metals): Elements filling d orbitals.
- f-block (Lanthanoids and Actinoids): Elements with last electrons in f orbitals.

Quantum Numbers:
- Principal quantum number (n) indicates the main energy level (shell).
- Azimuthal quantum number (l) indicates the shape of the orbital.
Elasticity in Configuration:
- As atomic number increases in a period, one electron is added to the outermost shell, maintaining the unique electronic configuration.

First Period (Z=1 to 2):
- Filled: 1s (H: 1s¹, He: 1s²).
- Contains 2 elements.
Second Period (Z=3 to 10):
- Starts with Li: 1s² 2s¹ and ends with Ne: 1s²2s²2p⁶.
- Contains 8 elements.
Third Period (Z=11 to 18):
- Begins: Na: 1s²2s²2p⁶3s¹ to Ar: 1s²2s²2p⁶3s²3p⁶.
- Contains 8 elements.
Fourth Period (Z=19 to 36):
- First two filled: K (4s¹) and Ca (4s²). Then 3d subshell fills.
- Total 18 elements (2 in s, 10 in d, 6 in p).
Fifth Period (Similar to Fourth):
- Contains 18 elements (5s, 4d, 5p subshells).

s-Block Configuration: ns¹-² (Alkali and alkaline earth metals)
p-Block Configuration: ns²np¹-⁶
The trend keeps the same configuration down the group; in groups 13-18, valence configurations are maintained.

s-Block Elements:
- Groups 1-2, known for reactivity and forming cations.
- Typically found in nature in combined states; low ionization energies.
p-Block Elements:
- Groups 13-18, including metals, nonmetals, and metalloids.
- Significant variation in reactivity; inert gases exhibit low reactivity due to fully filled valence shells.
d-Block Elements:
- Transition metals found in groups 3-12, known for partially filled d-orbitals leading to a variety of oxidation states.
- Exhibit properties like paramagnetism, colored compounds, and catalytic abilities.
f-Block Elements:
- Comprising lanthanides and actinides, filling of f-subshells leads to unique properties. They exhibit similar behaviors within a series.

Effective Nuclear Charge (Zeff):
- The net positive charge experienced by an electron in a multi-electron atom.
- Zeff is calculated as Zeff = Z - σ where σ is the shielding constant.
- Increases across a period, decreases down a group due to increased shielding from additional inner-shell electrons.

Atomic Radius:
- Typically decreases across a period due to increasing Zeff and increases down a group due to added electron shells.
Ionic Radius:
- Generally, cations are smaller than their parent atoms, while anions are larger.

Ionization Energy:
- Energy required to remove an electron: generally increases across a period and decreases down a group due to changes in Zeff.
Electron Gain Enthalpy:
- Trend varies with elements gaining electrons more readily if their Zeff is higher.
Electronegativity:
- Tendency to attract shared electrons in covalent bonds, it increases across a period and decreases down a group.

The periodic table offers a systematic way to understand the properties of elements, their relationships, and trends based on atomic structure. The organization reflects both historical developments and modern understandings of atomic properties and electronic configurations. It continues to be a vital tool in chemistry for predicting behavior and reactivity of substances.