Structure 3.1.2 – Periods, Groups, Classifications & Electron Configuration

The periodic table is arranged in periods (horizontal rows) and groups (vertical columns). This structure reflects the electron configurations of elements and helps predict their chemical behavior.

Period Number

The period number of an element corresponds to the number of occupied energy levels (electron shells) in its atoms.

• For example, all elements in Period 2 have electrons occupying two energy levels (n = 1 and n = 2).

Group Number and Valence Electrons

Elements in the same group have the same number of valence electrons—the electrons in the outermost shell. This is why they show similar chemical properties.

• Group 1 elements all have 1 valence electron.
• Group 17 elements all have 7 valence electrons.

Groups are numbered from 1 to 18 in modern periodic tables. The group number often indicates the number of valence electrons for main-group (s- and p-block) elements.

Important Classifications of Elements

1. Alkali Metals (Group 1)

These are soft, highly reactive metals with one valence electron. They react vigorously with water to form alkaline solutions and hydrogen gas.

• Examples: Lithium (Li), Sodium (Na), Potassium (K)
• 

2. Halogens (Group 17)

Halogens are reactive nonmetals with seven valence electrons. They tend to gain one electron to form halide ions (X⁻).

• Examples: Fluorine (F), Chlorine (Cl), Bromine (Br)
• 

3. Noble Gases (Group 18)

Noble gases are colorless, monatomic gases with full outer electron shells. They are chemically inert under standard conditions.

• Examples: Helium (He), Neon (Ne), Argon (Ar)
• 

4. Transition Elements (Groups 3–12)

Transition metals are d-block elements that can form multiple oxidation states and colored compounds. They often act as catalysts.

• Examples: Iron (Fe), Copper (Cu), Zinc (Zn)
• 

Deducing the Electron Configuration of an Atom (up to Z = 36)

Understanding the arrangement of electrons in atoms is essential to interpreting chemical behavior. The periodic table is designed in such a way that an element’s position (group and period) reveals its electron configuration, and vice versa. For atomic numbers up to 36 (Krypton), the configurations follow predictable filling of orbitals based on quantum mechanics.

Electron Configuration Order (Aufbau Principle)

Electrons fill orbitals in order of increasing energy:

• 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p
• 

Each type of orbital can hold:

• s: 2 electrons
• p: 6 electrons
• d: 10 electrons

Periodic Table Blocks

• s-block: Groups 1–2 (e.g., H, He, Na)
• p-block: Groups 13–18 (e.g., C, N, O, F)
• d-block: Transition metals (Groups 3–12)

From Position → Configuration

Steps:

1. Identify the period → number of energy levels
2. Identify the block (s, p, d) to find the sublevel being filled
3. Use the group number to determine valence electrons (main group)
4. Write configuration in order of orbital filling up to the atomic number

From Configuration → Position

Steps:

1. Count total electrons to find the atomic number
2. Identify highest principal quantum number (n) → period
3. For s- and p-blocks: use outermost electrons to determine group
4. For d-blocks: last electrons go into (n−1)d → identify as transition metal

Noble Gas Shorthand Notation

To simplify electron configurations, use the nearest noble gas before the element in brackets, followed by the rest of the configuration:

• Example: Calcium (Z = 20) → [Ar] 4s²