Chemical Bonding Lessons | Chemistry

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Bond Types

Metal + Nonmetal = Ionic · Two Nonmetals = Covalent

Electronegativity difference >1.7 = ionic · <1.7 = covalent · 0 = nonpolar covalent

The fastest bond-type rule: metal + nonmetal → ionic (electron transfer). Two nonmetals → covalent (electron sharing). Metals together → metallic (electron sea). For precision: electronegativity difference >1.7 = ionic; 0.4–1.7 = polar covalent; <0.4 = nonpolar covalent.

📖 Full Lesson → Difficulty: Beginner

Ionic bonding

Complete transfer of electrons. Metal loses electrons (cation); nonmetal gains (anion). NaCl: Na⁺ and Cl⁻. Crystal lattice held together by electrostatic attraction. High melting point, conducts electricity when dissolved.

Covalent bonding

Sharing of electrons. Two nonmetals share pairs to both achieve noble gas configuration. H₂O, CO₂, CH₄. Usually lower melting point than ionic. Doesn't conduct electricity.

Metallic bonding

"Electron sea" model: metal cations in a lattice surrounded by delocalized electrons. Explains conductivity, malleability, and ductility of metals.

ΔEN boundary note

The 1.7 cutoff is a guideline, not a law. Many chemists use 1.8 or 2.0. The ionic/covalent distinction is a continuum — highly polar covalent bonds share properties of both.

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🃏 Bond Types

Ionic vs covalent bonds — how can you tell?

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🃏 Answer

Metal + Nonmetal = Ionic · Two Nonmetals = Covalent

Ionic bonding — Complete transfer of electrons. Metal loses electrons (cation); nonmetal gains (anion). NaCl: Na⁺ and Cl⁻. Crystal lattice held together by electrostatic attraction. High melting point, conducts electricity when dissolved.

Covalent bonding — Sharing of electrons. Two nonmetals share pairs to both achieve noble gas configuration. H₂O, CO₂, CH₄. Usually lower melting point than ionic. Doesn't conduct electricity.

Metallic bonding — "Electron sea" model: metal cations in a lattice surrounded by delocalized electrons. Explains conductivity, malleability, and ductility of metals.

ΔEN boundary note — The 1.7 cutoff is a guideline, not a law. Many chemists use 1.8 or 2.0. The ionic/covalent distinction is a continuum — highly polar covalent bonds share properties of both.

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Lewis Structures

SECS (S=Structure/skeleton, E=Electrons counted, C=Connect with bonds, S=Satisfy octets)

Step-by-step Lewis dot structure method

SECS: Structure (draw skeleton — least electronegative atom = central), Electrons (count total valence electrons), Connect (single bonds first — uses 2e each), Satisfy (complete octets with lone pairs, use double/triple bonds if electrons run out). Hydrogen always gets exactly 2 electrons. Octet rule: most atoms want 8 valence electrons.

📖 Full Lesson → Difficulty: Intermediate

Counting valence electrons

Use group number. Add 1e per negative charge; subtract 1e per positive charge. H₂O: 2(1) + 6 = 8e total. CO₂: 4 + 2(6) = 16e total.

Formal charge

FC = (valence e) - (lone pair e) - ½(bonding e). Best structure minimizes formal charges and puts negative FC on most electronegative atom.

Octet exceptions

Expanded octet: Period 3+ atoms (S, P, Cl) can have 10 or 12 electrons — they have d orbitals. Incomplete octet: B and Al can have 6 electrons (BF₃). Odd electrons: NO, NO₂ (free radicals).

Resonance

When multiple Lewis structures are equally valid, draw all of them. True structure is the average (resonance hybrid). Benzene's C-C bonds are all equal length — neither single nor double.

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🃏 Lewis Structures

SECS

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🃏 Answer

SECS (S=Structure/skeleton, E=Electrons counted, C=Connect with bonds, S=Satisfy octets)

Counting valence electrons — Use group number. Add 1e per negative charge; subtract 1e per positive charge. H₂O: 2(1) + 6 = 8e total. CO₂: 4 + 2(6) = 16e total.

Formal charge — FC = (valence e) - (lone pair e) - ½(bonding e). Best structure minimizes formal charges and puts negative FC on most electronegative atom.

Octet exceptions — Expanded octet: Period 3+ atoms (S, P, Cl) can have 10 or 12 electrons — they have d orbitals. Incomplete octet: B and Al can have 6 electrons (BF₃). Odd electrons: NO, NO₂ (free radicals).

Resonance — When multiple Lewis structures are equally valid, draw all of them. True structure is the average (resonance hybrid). Benzene's C-C bonds are all equal length — neither single nor double.

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VSEPR

Lone pairs compress — they take more space than bonding pairs

Electron pairs repel: lone pairs > bonding pairs in repulsion strength

VSEPR (Valence Shell Electron Pair Repulsion): electron pairs arrange to minimize repulsion. Lone pairs repel more than bonding pairs, compressing bond angles. H₂O has 2 bonding + 2 lone pairs — tetrahedral electron geometry but BENT molecular geometry (104.5° not 109.5°). NH₃: 3 bonding + 1 lone → trigonal pyramidal (107°).

📖 Full Lesson → Difficulty: Intermediate

Common shapes

2 groups = linear (180°). 3 = trigonal planar (120°). 4 = tetrahedral (109.5°). 5 = trigonal bipyramidal. 6 = octahedral (90°). These are electron geometries with no lone pairs.

Molecular vs electron geometry

Electron geometry = count all pairs (bonding + lone). Molecular geometry = only count bonding pairs. H₂O: electron = tetrahedral, molecular = bent.

Bond angle compression

Each lone pair compresses angles by ~2.5°. Methane (no lone pairs) = 109.5°. Ammonia (1 lone pair) ≈ 107°. Water (2 lone pairs) ≈ 104.5°.

Linear always nonpolar

CO₂ is linear — even though C=O bonds are polar, they point in opposite directions and cancel. Molecular geometry determines whether dipoles cancel.

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🃏 VSEPR

VSEPR — why do lone pairs change bond angles?

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🃏 Answer

Lone pairs compress — they take more space than bonding pairs

Common shapes — 2 groups = linear (180°). 3 = trigonal planar (120°). 4 = tetrahedral (109.5°). 5 = trigonal bipyramidal. 6 = octahedral (90°). These are electron geometries with no lone pairs.

Molecular vs electron geometry — Electron geometry = count all pairs (bonding + lone). Molecular geometry = only count bonding pairs. H₂O: electron = tetrahedral, molecular = bent.

Bond angle compression — Each lone pair compresses angles by ~2.5°. Methane (no lone pairs) = 109.5°. Ammonia (1 lone pair) ≈ 107°. Water (2 lone pairs) ≈ 104.5°.

Linear always nonpolar — CO₂ is linear — even though C=O bonds are polar, they point in opposite directions and cancel. Molecular geometry determines whether dipoles cancel.

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Polarity

Polar molecule = asymmetric shape + polar bonds that DON'T cancel

Two conditions needed: polar bonds AND geometry that doesn't cancel dipoles

A molecule is polar only if it has BOTH polar bonds AND a geometry where the dipoles don't cancel. CO₂ = polar bonds but linear → dipoles cancel → nonpolar. H₂O = polar bonds + bent → dipoles don't cancel → polar. Always check: are all identical groups surrounding the central atom? If yes, likely nonpolar. Asymmetric = polar.

📖 Full Lesson → Difficulty: Intermediate

Nonpolar examples

CO₂ (linear, cancels), CCl₄ (tetrahedral, all Cl same), BF₃ (trigonal planar, all F same), N₂, H₂. Perfect symmetry = all dipoles cancel.

Polar examples

H₂O (bent), NH₃ (pyramidal), HCl (diatomic, different atoms), CHCl₃ (3 Cl + 1 H = asymmetric), SO₂ (bent). One lone pair or different groups breaks symmetry.

Like dissolves like

Polar solvents (water) dissolve polar/ionic solutes. Nonpolar solvents (hexane) dissolve nonpolar solutes. Oil and water don't mix because oil is nonpolar and water is polar.

Dipole moment

Measured in Debyes (D). Points from positive to negative end. HF has a larger dipole moment than HCl because F is more electronegative and the bond is shorter.

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🃏 Polarity

What makes a molecule polar?

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🃏 Answer

Polar molecule = asymmetric shape + polar bonds that DON'T cancel

Nonpolar examples — CO₂ (linear, cancels), CCl₄ (tetrahedral, all Cl same), BF₃ (trigonal planar, all F same), N₂, H₂. Perfect symmetry = all dipoles cancel.

Polar examples — H₂O (bent), NH₃ (pyramidal), HCl (diatomic, different atoms), CHCl₃ (3 Cl + 1 H = asymmetric), SO₂ (bent). One lone pair or different groups breaks symmetry.

Like dissolves like — Polar solvents (water) dissolve polar/ionic solutes. Nonpolar solvents (hexane) dissolve nonpolar solutes. Oil and water don't mix because oil is nonpolar and water is polar.

Dipole moment — Measured in Debyes (D). Points from positive to negative end. HF has a larger dipole moment than HCl because F is more electronegative and the bond is shorter.

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IMFs

London < Dipole–Dipole < H-bond — "LDH (L=London dispersion, D=Dipole-dipole, H=Hydrogen bonding) — weakest to strongest"

London dispersion (all) · Dipole-dipole (polar) · Hydrogen bonds (N, O, F with H)

Three types of intermolecular forces in order of increasing strength: London dispersion forces (all molecules), dipole-dipole (polar molecules), hydrogen bonds (H directly bonded to N, O, or F). Higher IMF strength = higher boiling point = more energy needed to separate molecules. Water's anomalously high boiling point (100°C for such a small molecule) is due to hydrogen bonding.

📖 Full Lesson → Difficulty: Intermediate

London dispersion forces

Present in ALL molecules. Caused by temporary dipoles from electron movement. Strength increases with molecular size (more electrons = bigger temporary dipoles). Noble gases only have LDF — hence very low boiling points.

Dipole-dipole forces

Between permanent dipoles in polar molecules. Positive end of one molecule attracted to negative end of another. Stronger than LDF for similar-sized molecules.

Hydrogen bonding

Special strong dipole-dipole: H must be bonded to N, O, or F (the three most electronegative small atoms). Explains water's surface tension, DNA base pairing, and protein secondary structure (α-helix).

Boiling point prediction

Compare IMF types first. If same type, larger molecule = more LDF = higher BP. H₂O boils at 100°C; H₂S (no H-bonds, only LDF) boils at -60°C despite being heavier.

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🃏 IMFs

Intermolecular forces (LDH) — weakest to strongest?

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🃏 Answer

Intermolecular forces, weakest to strongest — LDH

London < Dipole–Dipole < H-bond — "LDH (L=London dispersion, D=Dipole-dipole, H=Hydrogen bonding) — weakest to strongest"

London dispersion forces — Present in ALL molecules. Caused by temporary dipoles from electron movement. Strength increases with molecular size (more electrons = bigger temporary dipoles). Noble gases only have LDF — hence very low boiling points.

Dipole-dipole forces — Between permanent dipoles in polar molecules. Positive end of one molecule attracted to negative end of another. Stronger than LDF for similar-sized molecules.

Hydrogen bonding — Special strong dipole-dipole: H must be bonded to N, O, or F (the three most electronegative small atoms). Explains water's surface tension, DNA base pairing, and protein secondary structure (α-helix).

Boiling point prediction — Compare IMF types first. If same type, larger molecule = more LDF = higher BP. H₂O boils at 100°C; H₂S (no H-bonds, only LDF) boils at -60°C despite being heavier.

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Hybridization

Count electron groups → sp, sp², sp³, sp³d, sp³d²

2 groups=sp · 3=sp² · 4=sp³ · 5=sp³d · 6=sp³d²

Hybridization is determined by the number of electron groups (bonding pairs + lone pairs) around the central atom. Count all groups including lone pairs: 2 groups → sp (linear), 3 → sp² (trigonal planar), 4 → sp³ (tetrahedral). Multiple bonds count as ONE group — a double bond = 1 group. C in CO₂ has 2 groups → sp.

📖 Full Lesson → Difficulty: Intermediate

sp — 2 groups

Linear, 180°. Two sp orbitals + two unhybridized p orbitals. Examples: C in CO₂, C in HC≡CH (acetylene), Be in BeCl₂.

sp² — 3 groups

Trigonal planar, 120°. Three sp² + one unhybridized p orbital. Examples: C in ethylene (H₂C=CH₂), B in BF₃, carbons in benzene. The unhybridized p forms π bonds.

sp³ — 4 groups

Tetrahedral, 109.5°. Four sp³ orbitals. Examples: C in CH₄, N in NH₃ (4 groups including lone pair), O in H₂O (4 groups including 2 lone pairs).

σ and π bonds

Single bonds = 1σ. Double bonds = 1σ + 1π. Triple bonds = 1σ + 2π. σ bonds form by head-on overlap (along bonding axis). π bonds form by side-on overlap of p orbitals.

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🃏 Hybridization

Hybridization — how do you determine it?

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🃏 Answer

Count electron groups → sp, sp², sp³, sp³d, sp³d²

sp — 2 groups — Linear, 180°. Two sp orbitals + two unhybridized p orbitals. Examples: C in CO₂, C in HC≡CH (acetylene), Be in BeCl₂.

sp² — 3 groups — Trigonal planar, 120°. Three sp² + one unhybridized p orbital. Examples: C in ethylene (H₂C=CH₂), B in BF₃, carbons in benzene. The unhybridized p forms π bonds.

sp³ — 4 groups — Tetrahedral, 109.5°. Four sp³ orbitals. Examples: C in CH₄, N in NH₃ (4 groups including lone pair), O in H₂O (4 groups including 2 lone pairs).

σ and π bonds — Single bonds = 1σ. Double bonds = 1σ + 1π. Triple bonds = 1σ + 2π. σ bonds form by head-on overlap (along bonding axis). π bonds form by side-on overlap of p orbitals.

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Ionic Properties

Ionic = "Hard, High, Conduct when melted or dissolved"

High MP/BP · Hard but brittle · Conducts only when liquid or dissolved

Ionic compounds have a signature property set: hard (strong lattice) but brittle (ions shift, like charges align and repel), very high melting and boiling points (strong electrostatic forces), and they only conduct electricity when dissolved in water or melted (ions must be free to move — solid lattice locks them in place). NaCl dissolves in water; olive oil doesn't.

📖 Full Lesson → Difficulty: Beginner

Lattice energy

Energy released when gaseous ions form the crystal lattice. Higher lattice energy = stronger ionic bond = higher MP. Increases with higher charge and smaller ionic radius. MgO > NaCl in lattice energy (2+ and 2- vs 1+ and 1-).

Why brittle

When you strike an ionic crystal, layers shift. Like charges align → electrostatic repulsion → the crystal splits along crystal planes. Metals don't do this because electrons flow, maintaining attraction.

Electrolytes

Ionic compounds that dissociate completely are strong electrolytes. NaCl, KBr, MgCl₂. They produce solutions with high electrical conductivity. Important for biological function — nerves use Na⁺/K⁺ gradients.

Polyatomic ions

Common ones: SO₄²⁻ (sulfate), NO₃⁻ (nitrate), PO₄³⁻ (phosphate), OH⁻ (hydroxide), NH₄⁺ (ammonium), CO₃²⁻ (carbonate). Learn these — they appear constantly.

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🃏 Ionic Properties

Ionic compounds — the key properties?

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🃏 Answer

Ionic = "Hard, High, Conduct when melted or dissolved"

Lattice energy — Energy released when gaseous ions form the crystal lattice. Higher lattice energy = stronger ionic bond = higher MP. Increases with higher charge and smaller ionic radius. MgO > NaCl in lattice energy (2+ and 2- vs 1+ and 1-).

Why brittle — When you strike an ionic crystal, layers shift. Like charges align → electrostatic repulsion → the crystal splits along crystal planes. Metals don't do this because electrons flow, maintaining attraction.

Electrolytes — Ionic compounds that dissociate completely are strong electrolytes. NaCl, KBr, MgCl₂. They produce solutions with high electrical conductivity. Important for biological function — nerves use Na⁺/K⁺ gradients.

Polyatomic ions — Common ones: SO₄²⁻ (sulfate), NO₃⁻ (nitrate), PO₄³⁻ (phosphate), OH⁻ (hydroxide), NH₄⁺ (ammonium), CO₃²⁻ (carbonate). Learn these — they appear constantly.

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Metallic Bonding

Metals = "Electron Sea" — cations floating in mobile electrons

Delocalized electrons explain conductivity, malleability, and luster

The metallic bonding model: positive metal cations fixed in a lattice, surrounded by a "sea" of delocalized (free-moving) electrons. These mobile electrons explain every metallic property. Electrical conductivity: electrons carry charge. Thermal conductivity: same electrons transfer kinetic energy. Malleability/ductility: layers slide past each other while electrons maintain attraction throughout.

📖 Full Lesson → Difficulty: Beginner

Why metals conduct

Delocalized electrons are free to move throughout the lattice. Apply a voltage → electrons flow → current. Unlike ionic compounds (fixed lattice) or covalent compounds (electrons localized in bonds).

Why metals are malleable

When a metal is struck, cation layers slide — but the electron sea maintains bonding throughout. No like-charge repulsion. Gold can be hammered into sheets a few atoms thick.

Alloys

Mixing two metals disrupts the regular lattice — makes it harder to slide layers (more strength) but usually reduces conductivity. Steel (Fe + C) is stronger than pure iron. Bronze (Cu + Sn) harder than pure copper.

Band theory

Modern view: atomic orbitals combine into bands. Metals have partially filled conduction bands → easy electron flow. Insulators have full valence band + large gap. Semiconductors have small gap — conductivity can be switched on.

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🃏 Metallic Bonding

Metallic bonding — what is the electron sea?

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🃏 Answer

Metals = "Electron Sea" — cations floating in mobile electrons

Why metals conduct — Delocalized electrons are free to move throughout the lattice. Apply a voltage → electrons flow → current. Unlike ionic compounds (fixed lattice) or covalent compounds (electrons localized in bonds).

Why metals are malleable — When a metal is struck, cation layers slide — but the electron sea maintains bonding throughout. No like-charge repulsion. Gold can be hammered into sheets a few atoms thick.

Alloys — Mixing two metals disrupts the regular lattice — makes it harder to slide layers (more strength) but usually reduces conductivity. Steel (Fe + C) is stronger than pure iron. Bronze (Cu + Sn) harder than pure copper.

Band theory — Modern view: atomic orbitals combine into bands. Metals have partially filled conduction bands → easy electron flow. Insulators have full valence band + large gap. Semiconductors have small gap — conductivity can be switched on.

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Bond Order

Higher bond order = shorter, stronger bond

Triple > double > single in strength and inversely in length

Bond order = number of bonding pairs between two atoms. Triple bonds (3) are shorter and stronger than double bonds (2), which are shorter and stronger than single bonds (1). C≡C is shorter than C=C is shorter than C-C. More electron density between nuclei = stronger attraction = harder to break = higher bond dissociation energy.

📖 Full Lesson → Difficulty: Intermediate

Bond length comparison

C-C ≈ 154 pm · C=C ≈ 134 pm · C≡C ≈ 120 pm. N-N ≈ 146 pm · N=N ≈ 122 pm · N≡N ≈ 110 pm. More bonds = shorter distance between nuclei.

Bond energy comparison

C-C ≈ 347 kJ/mol · C=C ≈ 614 kJ/mol · C≡C ≈ 835 kJ/mol. Note: double bond is NOT twice the single bond energy — π bond is weaker than σ bond.

Fractional bond orders

Resonance structures give fractional bond orders. Benzene C-C bond order = 1.5 (between 1 and 2). Ozone O-O bond order = 1.5. These bonds have intermediate length and strength.

MO theory bond order

In molecular orbital theory: bond order = ½(bonding e - antibonding e). O₂ bond order = ½(10-6) = 2 (double bond). Also predicts O₂ is paramagnetic (2 unpaired e in degenerate π* orbitals).

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🃏 Bond Order

Bond order — how does it affect length and strength?

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🃏 Answer

Higher bond order = shorter, stronger bond

Bond length comparison — C-C ≈ 154 pm · C=C ≈ 134 pm · C≡C ≈ 120 pm. N-N ≈ 146 pm · N=N ≈ 122 pm · N≡N ≈ 110 pm. More bonds = shorter distance between nuclei.

Bond energy comparison — C-C ≈ 347 kJ/mol · C=C ≈ 614 kJ/mol · C≡C ≈ 835 kJ/mol. Note: double bond is NOT twice the single bond energy — π bond is weaker than σ bond.

Fractional bond orders — Resonance structures give fractional bond orders. Benzene C-C bond order = 1.5 (between 1 and 2). Ozone O-O bond order = 1.5. These bonds have intermediate length and strength.

MO theory bond order — In molecular orbital theory: bond order = ½(bonding e - antibonding e). O₂ bond order = ½(10-6) = 2 (double bond). Also predicts O₂ is paramagnetic (2 unpaired e in degenerate π* orbitals).

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🎓 Common Exam Questions

Memory tricks for VSEPR, polarity & intermolecular forces

Memory Tricks