The Periodic Table
1. Atomic Structure and Electron Configuration
Atomic Number
The atomic number (Z) is the number of protons in the nucleus of an atom.
In a neutral atom, it is also equal to the number of electrons.
Example:
• Sodium (Na): Z = 11
Electron configuration:
1s² 2s² 2p⁶ 3s¹
Electron Configuration
Electrons occupy energy levels (shells) and orbitals around the nucleus.
Rules:
1. Electrons fill lowest energy orbitals first.
2. Orbitals hold maximum 2 electrons.
3. Orbitals fill singly before pairing (Hund's rule).
Example configurations:
Element Atomic Number Electron Configuration
Lithium 3 1s² 2s¹
Carbon 6 1s² 2s² 2p²
Oxygen 8 1s² 2s² 2p⁴
Sodium 11 1s² 2s² 2p⁶ 3s¹
These configurations help explain chemical behaviour.
2. Valence Theory
Valence Electrons
Valence electrons are electrons in the outermost shell.
They determine:
• Chemical reactivity
• Bond formation
• Position in the periodic table
Example:
, Element Valence Electrons Group
Na 1 Group 1
Mg 2 Group 2
Al 3 Group 13
Cl 7 Group 17
s-block and p-block Elements
s-block
• Groups 1–2
• Valence electrons in s orbital
• Highly reactive metals
Example: Na, Mg
p-block
• Groups 13–18
• Valence electrons in p orbitals
Example: C, N, O, Cl
(Transition metals are not included in this syllabus section.)
3. Periodic Trends
Periodic trends describe patterns of properties across periods and groups.
Across a Period (Left → Right)
As atomic number increases:
Property Trend Reason
Atomic radius Decreases Stronger nuclear attraction
Electronegativity Increases Nucleus attracts electrons more strongly
Ionisation energy Increases Harder to remove electrons
Metallic character Decreases Atoms hold electrons more strongly
Reason:
1. Atomic Structure and Electron Configuration
Atomic Number
The atomic number (Z) is the number of protons in the nucleus of an atom.
In a neutral atom, it is also equal to the number of electrons.
Example:
• Sodium (Na): Z = 11
Electron configuration:
1s² 2s² 2p⁶ 3s¹
Electron Configuration
Electrons occupy energy levels (shells) and orbitals around the nucleus.
Rules:
1. Electrons fill lowest energy orbitals first.
2. Orbitals hold maximum 2 electrons.
3. Orbitals fill singly before pairing (Hund's rule).
Example configurations:
Element Atomic Number Electron Configuration
Lithium 3 1s² 2s¹
Carbon 6 1s² 2s² 2p²
Oxygen 8 1s² 2s² 2p⁴
Sodium 11 1s² 2s² 2p⁶ 3s¹
These configurations help explain chemical behaviour.
2. Valence Theory
Valence Electrons
Valence electrons are electrons in the outermost shell.
They determine:
• Chemical reactivity
• Bond formation
• Position in the periodic table
Example:
, Element Valence Electrons Group
Na 1 Group 1
Mg 2 Group 2
Al 3 Group 13
Cl 7 Group 17
s-block and p-block Elements
s-block
• Groups 1–2
• Valence electrons in s orbital
• Highly reactive metals
Example: Na, Mg
p-block
• Groups 13–18
• Valence electrons in p orbitals
Example: C, N, O, Cl
(Transition metals are not included in this syllabus section.)
3. Periodic Trends
Periodic trends describe patterns of properties across periods and groups.
Across a Period (Left → Right)
As atomic number increases:
Property Trend Reason
Atomic radius Decreases Stronger nuclear attraction
Electronegativity Increases Nucleus attracts electrons more strongly
Ionisation energy Increases Harder to remove electrons
Metallic character Decreases Atoms hold electrons more strongly
Reason:
