Chemical Bonding — Science Sage Gr10

Why Do Atoms Bond?

Atoms bond to achieve a more stable electron configuration — typically a full outer shell of 8 electrons (the octet rule), or 2 electrons for hydrogen and helium. When atoms bond, they reach a lower energy state, which is more stable. The type of bond that forms depends on the difference in electronegativity (ΔEN) between the two atoms.

Rule of thumb for bond type based on electronegativity difference (ΔEN):
ΔEN > 1.7 → Ionic bond (electron transfer)
ΔEN 0.4–1.7 → Polar covalent bond (unequal sharing)
ΔEN < 0.4 → Non-polar covalent bond (equal sharing)

Ionic Bonding

Ionic bonding occurs between a metal (low electronegativity) and a non-metal (high electronegativity). The metal transfers one or more electrons to the non-metal, forming oppositely charged ions. The electrostatic attraction between these ions is the ionic bond.

Example — sodium chloride (NaCl): Na has EN = 0.9, Cl has EN = 3.0, so ΔEN = 2.1 > 1.7 → ionic bond. Na loses 1 electron → Na⁺; Cl gains 1 electron → Cl⁻. The Na⁺ and Cl⁻ ions arrange into a crystal lattice.

Property	Explanation
High melting/boiling point	Strong electrostatic forces between ions require lots of energy to break
Brittle	Shifting the lattice lines up like charges → repulsion → shatters
Conduct electricity when molten or dissolved	Ions are free to move and carry charge
Do NOT conduct when solid	Ions are fixed in lattice positions
Hard	Strong lattice forces hold ions tightly in place

Examples: NaCl (table salt), MgO (magnesium oxide), CaCl₂, Al₂O₃

Covalent Bonding

Covalent bonding occurs between two non-metals. Instead of transferring electrons, the atoms SHARE one or more pairs of electrons. Each shared pair forms one bond.

Bond Type	Shared Pairs	Example	Bond Order
Single bond	1	H₂, HCl, Cl₂, H₂O	1
Double bond	2	O₂, CO₂, C₂H₄	2
Triple bond	3	N₂, CO, C₂H₂	3

Lewis dot structures show the electrons in a molecule. Bonding pairs (shared electrons) are drawn between atoms; lone pairs (non-bonding electrons) are drawn on individual atoms.

Dative (coordinate) covalent bond: both electrons in the shared pair come from the SAME atom. Example: when NH₃ donates a lone pair to H⁺ to form NH₄⁺, or when H₂O donates a lone pair to H⁺ to form H₃O⁺.

Property	Explanation
Low melting/boiling point (usually)	Only weak intermolecular forces between molecules
Do NOT conduct electricity	No charged particles (no ions, no free electrons)
Can be gas, liquid or solid at room temperature	Depends on the strength of intermolecular forces
Often soluble in non-polar solvents	Like dissolves like

VSEPR Theory — Predicting Molecular Shape

VSEPR stands for Valence Shell Electron Pair Repulsion. The key idea: electron pairs (both bonding pairs and lone pairs) around a central atom repel each other and arrange themselves to be as far apart as possible. Lone pairs repel more strongly than bonding pairs.

Bonding Pairs	Lone Pairs	Shape	Bond Angle	Example
2	0	Linear	180°	CO₂, BeCl₂
3	0	Trigonal planar	120°	BF₃, SO₃
4	0	Tetrahedral	109.5°	CH₄, CCl₄
3	1	Trigonal pyramidal	~107°	NH₃
2	2	Bent/Angular	~104.5°	H₂O

Note: lone pairs take up more space than bonding pairs, which is why the bond angle in NH₃ (107°) is less than tetrahedral (109.5°), and in H₂O (104.5°) is even less.

Molecular Polarity

A bond is polar if there is an electronegativity difference between the two atoms — electrons are pulled towards the more electronegative atom. But a molecule can be non-polar even if it has polar bonds, if the molecule is symmetrical (dipoles cancel).

Polar molecule: asymmetric charge distribution — has a net dipole moment. Examples: H₂O (bent), NH₃ (pyramidal), HCl
Non-polar molecule: symmetric — dipoles cancel. Examples: CO₂ (linear), CH₄ (tetrahedral), CCl₄
Test: draw the shape and add dipole arrows (δ+ to δ-); if they don't cancel, molecule is polar

Metallic Bonding

In metals, the outer valence electrons are not attached to individual atoms but are free to move throughout the entire structure — they form a "sea of delocalised electrons." The positive metal ions (cations) are arranged in a regular lattice, surrounded by this electron sea.

Good electrical conductors — delocalised electrons carry charge
Good thermal conductors — delocalised electrons transfer kinetic energy
Malleable and ductile — layers of ions can slide without breaking bonds (the electron sea adjusts)
Lustrous — delocalised electrons interact with and reflect light
High melting points (generally) — strong electrostatic attraction between cations and electron sea

IEB Extension — Bond Order, Resonance & Formal Charge

Bond order is the number of shared electron pairs between two atoms. Higher bond order → shorter bond length, stronger bond (higher bond energy). For N₂ (triple bond, bond order 3): very short (110 pm), very strong (945 kJ·mol⁻¹). Compare: N–N single bond (145 pm, 163 kJ·mol⁻¹).

Resonance occurs when a molecule cannot be described by a single Lewis structure. For SO₃, three equivalent Lewis structures can be drawn — each showing a double bond to a different oxygen. The real structure is an average (resonance hybrid) with all S–O bonds equal in length (bond order 1.33). Other examples: CO₃²⁻, benzene (C₆H₆), NO₂⁻.

Formal charge = (valence electrons of atom) − (lone pair electrons) − ½(bonding electrons). The best Lewis structure has formal charges closest to zero. For CO₂: C has formal charge 0, each O has formal charge 0 — confirming the double-bond structure is correct.

Exceptions to the octet rule:

Electron-deficient molecules: BF₃ has only 6 electrons around B (yet is stable — B forms a 3rd bond with lone pairs from F in practice)
Odd-electron molecules: NO has 11 electrons (radical) — cannot satisfy octet for both atoms
Expanded octets: SF₆ (12 electrons around S), PCl₅ (10 electrons around P) — possible because Period 3+ elements can use d-orbitals

Bond Visualiser

Bond Type

Compound

Question 1 · 2 marks

The electronegativity values of Na and Cl are 0.9 and 3.0 respectively. What type of bond forms between them?

Question 2 · 3 marks

Which molecule has a TETRAHEDRAL shape?

Question 3 · 3 marks

Water (H₂O) has a bent shape. The bond angle is approximately:

Question 4 · 1 mark

Which property of metals is BEST explained by the sea of delocalised electrons model?

Question 5 · 2 marks

A molecule of CO₂ has two polar C=O bonds but is a non-polar molecule. Why?

Question 6 · 2 marks

In the formation of NH₄⁺ from NH₃ + H⁺, the bond formed between N and H⁺ is called:

Question 7 · 3 marks

Magnesium oxide (MgO) conducts electricity when melted, but tetrachloromethane (CCl₄, EN: C = 2.5, Cl = 3.0) does not conduct electricity even when it is a liquid. What is the BEST explanation for this difference?

Question 8 · 3 marks

Phosphorus trichloride (PCl₃) has a central P atom with 5 valence electrons, bonded to 3 Cl atoms. Using VSEPR theory, what is PCl₃'s shape and polarity?

IEB Extension Questions

IEB Question 1 · 2 marks IEB

The Lewis structure of SO₃ appears to show one S=O double bond and two S–O single bonds. However, all S–O bond lengths in SO₃ are equal (142 pm). The best explanation is:

IEB Question 2 · 1 mark IEB

Calculate the formal charge on each atom in CO₂. Given: C has 0 lone pair electrons and 8 bonding electrons; each O has 4 lone pair electrons and 4 bonding electrons.

Answer in full sentences. Draw Lewis structures where asked.

Question 1

Draw the Lewis (dot-and-cross) structure for each of the following molecules and identify all bonding pairs and lone pairs: (a) Cl₂ (b) H₂O (c) N₂

Question 2

Use VSEPR theory to predict the shape and bond angle of: (a) BF₃ (b) NH₃ (c) CCl₄. Explain your reasoning for each.

Question 3

Classify each bond as ionic, polar covalent, or non-polar covalent based on electronegativity differences: (a) H–F (EN: H=2.1, F=4.0) (b) C–H (EN: C=2.5, H=2.1) (c) Na–O (EN: Na=0.9, O=3.5)

Question 4

Explain why solid NaCl does not conduct electricity, but molten NaCl does. Draw a diagram (described in words) showing the lattice structure.

Question 5

Compare ionic and metallic bonding: (a) describe the nature of the bond in each case (b) give two physical properties that differ between ionic and metallic substances and explain why (c) name one example of each type.

Question 6

The table shows melting point and electrical conductivity data measured for three unknown substances, P, Q and R.

Substance	Melting point (°C)	Conducts as a solid?	Conducts molten or dissolved?
P	801	No	Yes
Q	−114	No	No
R	1085	Yes	Yes

(a) Using only the data in the table, identify the type of bonding (ionic, covalent, or metallic) present in each of P, Q and R.
(b) For each substance, justify your answer by referring to the conductivity pattern in the table.
(c) P is actually sodium chloride and R is actually copper. Explain, in terms of the particles present in each, why R conducts electricity even as a solid, while P only conducts once molten or dissolved.