Chemical Bonding Memory Tricks | Chemistry

🔗 Chemical Bonding — 9 Memory Tricks · Click any card to expand · Save favorites · Switch to Flashcard or Quiz mode below

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.

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.

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.

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.

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°).

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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).

🎓 Common Exam Questions

Q: What does SECS stand for and walk through drawing a Lewis structure using it.

A: SECS = Structure (draw skeleton — least electronegative atom in center, H always terminal), Electrons (count all valence electrons — add 1 per negative charge, subtract 1 per positive charge), Connect (place 2 electrons per bond between atoms), Satisfy (complete octets — start with outer atoms, then center; H needs only 2). Example: SO2 — S is center (less electronegative than O). Count: S(6) + O(6) + O(6) = 18 electrons. Connect: S-O-S uses 4. Remaining 14 fill octets. Check: S may need double bond. Result: O=S-O with lone pairs, S has 10 electrons (expanded octet allowed for period 3+).

Q: Explain VSEPR theory — what does it stand for and how do you predict molecular geometry?

A: VSEPR = Valence Shell Electron Pair Repulsion. Electron groups (bonding pairs AND lone pairs) repel each other and arrange to maximize separation. Count electron groups around central atom: 2 groups = linear (180°), 3 = trigonal planar (120°), 4 = tetrahedral (109.5°), 5 = trigonal bipyramidal, 6 = octahedral. Lone pairs compress bond angles more than bonding pairs. H2O: 4 electron groups (2 bonds + 2 lone pairs) = tetrahedral electron geometry but BENT molecular geometry (104.5°). NH3: 4 groups (3 bonds + 1 lone pair) = trigonal pyramidal (107°). Always distinguish electron geometry from molecular geometry.

Q: What are the three intermolecular forces (LDH) and what determines their relative strength?

A: LDH from weakest to strongest: London dispersion forces (L) — present in ALL molecules; caused by temporary dipoles from electron movement; strength increases with molecular size/surface area and number of electrons. Dipole-dipole forces (D) — in polar molecules; permanent dipoles attract; stronger than London for similar-sized molecules. Hydrogen bonding (H) — special strong dipole-dipole between H bonded to N, O, or F and a lone pair on N, O, or F on another molecule; explains water's unusually high boiling point. These forces explain boiling points, solubility ('like dissolves like'), and surface tension.

Q: What is hybridization and how do you determine it from a Lewis structure?

A: Hybridization is the mixing of atomic orbitals to form new equivalent orbitals for bonding. Count electron groups (bonding + lone pairs) around central atom: 2 = sp (linear, 180°), 3 = sp² (trigonal planar, 120°), 4 = sp³ (tetrahedral, 109.5°), 5 = sp³d (trigonal bipyramidal), 6 = sp³d² (octahedral). Key: double bonds count as ONE electron group. BF3 has 3 bonding groups = sp². CO2 (O=C=O) has 2 electron groups = sp. Ethylene (C2H4) each carbon has 3 groups (1 double + 2 single) = sp². Acetylene (C2H2) each carbon has 2 groups = sp. Hybridization determines geometry and bond angles.

Q: Compare ionic, covalent, and metallic bonding — properties of each.

A: Ionic bonding (metal + nonmetal): electron transfer creates cations and anions held by electrostatic attraction. Properties: high melting point, brittle, conducts electricity only when melted or dissolved (ions must be mobile), soluble in polar solvents. Covalent bonding (nonmetal + nonmetal): electron sharing. Molecular covalent: low melting point, nonconducting, variable solubility. Network covalent (diamond, SiO2): extremely high melting point, very hard, non-conducting. Metallic bonding (metal + metal): 'electron sea' — cations in mobile electron cloud. Properties: conducts electricity in solid state (mobile electrons), malleable (layers slide), lustrous (electrons reflect light), variable melting points.

Memory tricks for VSEPR, polarity & intermolecular forces

Memory Tricks