Arrangement of Elements

8.1 Arrangement of Elements

Definition:
The Periodic Table is a systematic arrangement of chemical elements, organised in periods (horizontal rows) and groups (vertical columns) in order of increasing proton number (atomic number).
Atomic number:
The number of protons in the nucleus of an atom; it determines the identity of the element.
Periods:
- Horizontal rows numbered from Period 1 to Period 7.
- Elements in the same period have the same number of electron shells.
- Across a period, the number of protons and electrons increases from left to right.
Groups:
- Vertical columns numbered Group I to Group VIII (or 1–18 in the IUPAC system).
- Elements in the same group have similar chemical properties because they have the same number of valence electrons.
Organisation rationale:
- The arrangement allows prediction of element properties.
- Elements are grouped so that patterns in reactivity, physical properties, and atomic structure are clear.
Special sections:
- Transition metals are placed in the middle block (Groups 3–12 in modern numbering).
- Lanthanides and actinides are placed separately to keep the table compact.

Metallic character:
- The tendency of an element to lose electrons and form positive ions (cations).
- Metals are typically on the left side of the Periodic Table.
Non-metallic character:
- The tendency of an element to gain electrons and form negative ions (anions) or share electrons.
- Non-metals are typically on the right side of the Periodic Table.
Trend across a period (left to right):
- Metallic character decreases.
- Non-metallic character increases.
- This is due to the increase in nuclear charge and decrease in atomic radius, making it harder for atoms to lose electrons but easier to gain them.
Example in Period 3:
Sodium (Na) → Magnesium (Mg) → Aluminium (Al) are metals,
Silicon (Si) is a metalloid,
Phosphorus (P) → Sulfur (S) → Chlorine (Cl) → Argon (Ar) are non-metals.

Group number (in Groups I–VIII) corresponds to the number of valence electrons.
The charge of the ion formed depends on whether the element loses or gains electrons to achieve a stable electronic configuration (usually a full outer shell).
Examples:
- Group I (e.g., Na): 1 valence electron → loses 1 electron → forms 1⁺ ions.
  Na → Na⁺ + e⁻
- Group II (e.g., Mg): 2 valence electrons → loses 2 electrons → forms 2⁺ ions.
  Mg → Mg²⁺ + 2e⁻
- Group VII (e.g., Cl): 7 valence electrons → gains 1 electron → forms 1⁻ ions.
  Cl + e⁻ → Cl⁻
- Group VI (e.g., O): 6 valence electrons → gains 2 electrons → forms 2⁻ ions.
  O + 2e⁻ → O²⁻

Reason:
Elements in the same group have the same number of electrons in their outer shell → undergo similar types of chemical reactions.
Examples:
- Group I (Alkali metals): React with water to produce hydrogen gas and an alkaline solution.
  2Na + 2H₂O → 2NaOH + H₂
- Group VII (Halogens): React with metals to form salts (metal halides).
  2Na + Cl₂ → 2NaCl
Trends in reactivity within a group:
- Metals (Group I): Reactivity increases down the group.
- Non-metals (Group VII): Reactivity decreases down the group.

Position tells:
- Period number → Number of electron shells.
- Group number → Number of valence electrons.
- Metal or non-metal nature based on side of the table.
Example:
- Element with atomic number 16 (Sulfur) is in Period 3, Group VI:
  - 3 electron shells.
  - 6 valence electrons.
  - Non-metal.
  - Likely to gain 2 electrons to form S²⁻.

Physical property trends:
- Group I: Melting and boiling points decrease down the group; density generally increases.
- Group VII: Melting and boiling points increase down the group; colours become darker.
Chemical property trends:
- Metals become more reactive down Group I.
- Non-metals become less reactive down Group VII.
Size trend:
- Atomic radius increases down a group due to more electron shells.
Electronegativity trend:
- Decreases down a group due to increased distance between nucleus and valence electrons.