Lesson One

Noble gases (He, Ne, Ar, Kr, Xe, Rn) are the most stable atoms because their outermost energy level is completely filled with electrons. Noble gas atoms don't undergo chemical reactions at normal conditions — their molecules are monatomic.

Noble gas	Electronic structure
Helium ₂He	1s²
Neon ₁₀Ne	[He], 2s², 2p⁶
Argon ₁₈Ar	[Ne], 3s², 3p⁶
Krypton ₃₆Kr	[Ar], 4s², 3d¹⁰, 4p⁶
Xenon ₅₄Xe	[Kr], 5s², 4d¹⁰, 5p⁶
Radon ₈₆Rn	[Xe], 6s², 4f¹⁴, 5d¹⁰, 6p⁶

All other elements are reactive — they undergo chemical reactions to complete their outermost shell by accepting, losing, or sharing electrons to acquire an electronic configuration similar to the nearest noble gas. As a result, bonds between atoms of reactant molecules are broken to form new bonds between atoms of product molecules → this is a chemical reaction.

Application: A mixture of iron filings with sulphur powder does NOT form a new compound (no chemical reaction). If heated to high temperature → a chemical reaction occurs (formation of chemical bond) between iron and sulphur → iron(II) sulphide compound.

Valence electrons play an important role in bond formation. Scientist Lewis introduced a method to represent valence electrons using dots around the element symbol.

For oxygen atom ₈O: electronic config = 1s², 2s², 2p⁴ → 6 valence electrons. Valence electrons are distributed singly as dots on four sides, then paired until fully distributed.

Group	1A	2A	3A	4A	5A	6A	7A	0
Element	₁₁Na	₁₂Mg	₁₃Al	₁₄Si	₁₅P	₁₆S	₁₇Cl	₁₈Ar
Config ends with	3s¹	3s²	3s²,3p¹	3s²,3p²	3s²,3p³	3s²,3p⁴	3s²,3p⁵	3s²,3p⁶
Valence e⁻	1	2	3	4	5	6	7	8

Lone pair: An electron pair found in one of the outer orbitals that doesn't share in bond formation.

Bond pair: An electron pair responsible for bond formation between two atoms.

Example: NH₃ has 1 lone pair and 3 bond pairs. Elements of the same group have similar numbers of lone pairs and bond pairs → they form similar compounds with other elements (e.g., H₂O and H₂S both have 2 bond pairs and 2 lone pairs).

Definition: The ionic bond is the bond resulting from the electrostatic attraction between the positive ion (cation) and the negative ion (anion) which are formed by the complete transfer of one or more electrons from the atom of a metallic element to the atom of a nonmetallic element.

Conditions for ionic bond formation:

1. The difference in electronegativity between the two bonded atoms must be more than 1.7
2. The ionization energy of the metallic element should be low
3. The electron affinity of the nonmetallic element should be high

Examples of ionic compounds:

• NaCl: Na (1s², 2s², 2p⁶, 3s¹) loses 1 electron → Na⁺ [like Ne]; Cl (1s², 2s², 2p⁶, 3s², 3p⁵) gains 1 electron → Cl⁻ [like Ar]
• MgO: Mg loses 2 electrons → Mg²⁺; O gains 2 electrons → O²⁻
• CaCl₂: Ca loses 2 electrons → Ca²⁺; each Cl gains 1 electron → 2Cl⁻

Compound	Elements	ΔEN	Ionic property
NaF	Na(0.9), F(4)	3.1	Strongest
NaCl	Na(0.9), Cl(3)	2.1	Strong
NaBr	Na(0.9), Br(2.8)	1.9	Less strong
NaI	Na(0.9), I(2.5)	1.6	Weakest

Conclusion: When horizontal distance between two bonded elements in the periodic table increases → electronegativity difference increases → ionic property increases.

Definition: The covalent bond is mostly formed between atoms of nonmetals with the same electronegativity (atoms of same element) or those of close electronegativities by sharing electrons, where each atom shares a certain number of electrons to complete the outermost shell.

Classification of covalent bonds (by electronegativity difference):

1. Pure covalent bond (ΔEN = 0): Formed between two atoms of the SAME nonmetal element. Both atoms attract bond pair equally → electron pair spends same time near each atom → net charge = zero. Examples: H-H in H₂, F-F in F₂.

2. Non-polar covalent bond (0 < ΔEN ≤ 0.4): Formed between two atoms of two DIFFERENT nonmetal elements. Example: C-H bond in methane CH₄ (2.5 − 2.1 = 0.4).

3. Polar covalent bond (0.4 < ΔEN < 1.7): More electronegative atom attracts bond pair more → acquires partial negative charge (δ⁻). Less electronegative atom acquires partial positive charge (δ⁺). As electronegativity difference increases → polarity increases.

Application — Polar covalent bond in HCl: In HCl, Cl is more electronegative → attracts electron pair more → Cl acquires δ⁻, H acquires δ⁺. ΔEN = 3 − 2.1 = 0.9.

Examples of polar compounds: Water H₂O: ΔEN = 3.5 − 2.1 = 1.4 → polar. Ammonia NH₃: ΔEN = 3 − 2.1 = 0.9 → polar.

Electronegativity difference scale:

• 0 → Pure covalent
• 0 to 0.4 → Non-polar covalent
• 0.4 to 1.7 → Polar covalent
• > 1.7 → Ionic

Worked Example: Determine type of bonds and arrange by polarity: H-Cl, C-O, H-H, N-O, P-Cl, C-H (Given: H=2.1, Cl=3, C=2.5, O=3.5, N=3, P=2.1)

Bond	ΔEN	Type
H-H	0	Pure covalent
C-H	0.4	Non-polar covalent
N-O	0.5	Polar covalent
H-Cl	0.9	Polar covalent
P-Cl	0.9	Polar covalent
C-O	1.0	Polar covalent

Order of increasing polarity: H-H < C-H < N-O < H-Cl = P-Cl < C-O