Hydrogen bonding

“Ortho-hydroxybenzaldehyde folds so its OH hydrogen points to the carbonyl oxygen, forming an internal hydrogen bond.”

Hydrogen bonds form when H attached to N, O or F interacts with a lone-pair donor; they can be intermolecular or intramolecular if geometry allows. A common mistake is to label H-bonds only between separate molecules—intramolecular H-bonds in ortho-substituted aromatics are classic examples.

Memory trick: H-bond: H-N/O/F + lone pair + near-linear → strong

Hydrogen bonding — common mistake

“Thinking H-bond is a covalent link is like assuming Velcro sticks as strongly as a welded joint.”

Hydrogen bonds are directional, weaker than covalent bonds (~5–30 kJ/mol) yet stronger than typical dipole–dipole forces. They dictate properties like water’s high ΔHvap. The exam trap is to count H-bonded H as part of the molecule’s valence.

Memory trick: H-bond ≠ covalent; it's a bridge, not a bond.

Fajans' rules

“A small, highly charged metal ion is like a strong magnet pulling electron clouds of the anion toward itself.”

Fajans' rules state that high charge density (small radius, high charge) on the cation polarises the anion, giving partial covalent character to an otherwise ionic bond. The nuance: even salts like AlCl₃ show covalent behaviour.

Memory trick: Charge⁺/size⁻ → covalent.

Fajans' rules — common mistake

“Assuming NaCl is purely ionic is like believing a steel ball is never magnetic; high-charge ions can distort electron clouds.”

Not all compounds written with ionic formulas are 100 % ionic; cations with high charge/small size (Al³⁺, Be²⁺) polarise large anions, leading to covalent character. The trap is to ignore cation size and charge.

Memory trick: Small⁺, high⁺ → covalent tilt.

Bond order = (N_b - N_a)/2

“Like counting occupied seats in a theater, each bonding electron adds a seat, each antibonding electron removes one, then halve the net.”

Bond order = (Nb – Na)/2 gives the number of effective bonds. Only valence electrons in the MO diagram are counted; core electrons stay in non-bonding orbitals. When Nb–Na is odd the result is half-integer (e.g., O2⁺ = 2.5). Trap: forgetting electrons in degenerate π* levels or using total electrons leads to a wrong bond order.

Bond order = (N_b - N_a)/2

Memory trick: Bond order = (bonding – antibonding)/2 → net handshakes ÷2

Steric number = sigma bonds + lone pairs

“A central atom’s party guest list counts each handshake (sigma) and each solo dancer (lone pair) to decide the dance floor layout.”

Steric number = sigma bonds + lone pairs; it equals the number of electron-pair regions and gives the hybridisation (sp, sp², sp³…). Trap: a double or triple bond contributes only one sigma, while each lone pair counts as one region.

Steric number = sigma bonds + lone pairs

Memory trick: SN = σ bonds + LPs, not π.

Dipole moment mu = q * d

“A tiny battery inside a molecule pushes charge across a distance, just like a flashlight’s current flows from the negative cell to the positive tip.”

Dipole moment μ equals charge magnitude q times separation d, directed from the negative to the positive charge. For homonuclear diatomics μ = 0 because charges cancel; forgetting the direction causes sign errors in vector addition.

Dipole moment mu = q * d

Memory trick: μ = q·d, arrow from – to +

Formal charge = V - N - B/2

“Balancing a checkbook: valence credits minus nonbonding debits minus half the shared expenses gives the formal charge.”

Formal charge = V − N − B/2, where V = valence electrons, N = non-bonding electrons, B = bonding electrons around the atom. Divide B by 2 because each bond is shared. Trap: forget to halve the electrons in multiple bonds; the most stable resonance has smallest |FC|.

Formal charge = V - N - B/2

Memory trick: FC = V – (N + B/2).

% ionic character rises with electronegativity difference

“When a tug-of-war team has one much stronger player, the rope pulls more towards that side, showing higher ionic character.”

% ionic character increases with electronegativity difference ΔEN, roughly given by %I = (1 − e^(-0.25ΔEN²)) × 100 (Pauling). The relation is exponential, not linear, and never reaches 100 % even for large ΔEN. Trap: assuming any ΔEN > 2 gives a fully ionic bond.

% ionic character rises with electronegativity difference

Memory trick: Ionic % ≈ (1-e^(-0.25ΔEN²))×100.

sigma bond vs pi bond

“A handshake (σ) lets you turn your hand freely, but a high-five (π) locks both arms together, stopping rotation.”

A σ bond forms by head-on overlap, is stronger and allows free rotation; a π bond forms by side-on overlap, is weaker and restricts rotation. A double bond = 1 σ + 1 π, a triple = 1 σ + 2 π. Trap: assuming extra bonds are all σ; only the first is σ.

Memory trick: σ = head-on, free spin; π = side-on, lock.

valence bond theory vs molecular orbital theory

“When arranging a family photo, VBT seats each person in a fixed spot, while MOT lets them mingle across the whole frame.”

VBT confines electrons to hybrid orbitals, giving accurate shapes but cannot explain paramagnetism or bond order. MOT spreads electrons over the whole molecule; bond order = (bonding-antibonding)/2 predicts stability and magnetism. Trap: O₂ is diamagnetic in VBT yet paramagnetic in MOT – always check bond order for O₂-type molecules.

Memory trick: VBT = V for fixed, MOT = M for mobile

ionic bond vs covalent bond

“Like a basketball being handed completely from one player to another (ionic) versus two friends gently passing it back and forth (covalent).”

Ionic bond forms when ΔEN ≳1.7, causing near-complete electron transfer and lattice energy dominance; covalent when ΔEN is small, electrons are shared. Real bonds often have mixed character; high polarising cations (Al³⁺) make AlCl₃ covalent despite metal-nonmetal. Trap: Do not label every metal-nonmetal compound as purely ionic.

Memory trick: ΔEN >1.7 → ION, else COV

Memory trick: LP pushes harder: LP-LP > LP-BP > BP-BP

Enthalpy and Hess's law

“Making a stew, the total heat needed is the same whether you add carrots first or potatoes first—Hess’s law.”

Enthalpy is a state function; the overall ΔH for a reaction equals the sum of ΔH for any series of steps connecting reactants to products, provided all steps are at the same temperature and pressure. The path can be arbitrary, but every intermediate must be included correctly.

Memory trick: Hess = Add, any order, same T,P

Enthalpy and Hess's law — common mistake

“Reversing a treadmill’s direction makes the work you get back opposite to what you put in—sign flips in Hess’s law.”

When a step in a Hess cycle is used in reverse, its ΔH changes sign. Forgetting to flip the sign gives an incorrect total enthalpy, even if the overall reaction appears balanced.

Memory trick: Reverse step → flip ΔH sign

Entropy

“When ice cream melts in a bowl, the solid becomes a flowing liquid, increasing disorder.”

Entropy (S) counts the number of microscopic arrangements. It rises when a solid melts, a gas expands or particles mix because microstates increase. Trap: the system’s entropy can decrease (e.g., freezing) if the surroundings compensate, so only ΔS_total ≥ 0 is required.

Memory trick: S ↑ when microstates ↑; total ΔS ≥ 0 for spontaneity.

Entropy — common mistake

“When a pressure cooker whistles, steam rushes out while the water inside can condense, showing system entropy can drop.”

A frequent mistake is assuming ΔS is always positive. In reality only the sum of system and surroundings must increase. Processes like freezing lower system entropy but raise surroundings’ entropy enough to keep ΔS_total ≥ 0.

Memory trick: ΔS_sys may be negative; verify ΔS_total.

Gibbs free energy

“A battery powering a torch discharges because its ΔG is negative, turning chemical energy into light.”

Gibbs free energy G = H − T S predicts spontaneity at constant T and P. ΔG < 0 means the reaction proceeds without external work. Trap: at high temperature a positive ΔH can be overcome if ΔS is positive, making ΔG negative.

Memory trick: ΔG = ΔH − TΔS; negative = spontaneous.

Gibbs free energy — common mistake

“A cold pack feels cold as ammonium nitrate dissolves; the endothermic step still occurs because ΔG is negative at room temperature.”

Students often judge spontaneity from enthalpy alone. Even an endothermic reaction (ΔH > 0) can be spontaneous if the TΔS term is large enough, giving ΔG < 0. Always evaluate both ΔH and ΔS.

Memory trick: Endothermic + large ΔS → ΔG < 0.

Temperature dependence of spontaneity

“Ice cubes vanish in a warm drink but stay solid in a cold glass, showing ΔH>0, ΔS>0 needs enough T to be spontaneous.”

For a process with ΔH>0 and ΔS>0, ΔG=ΔH–TΔS. At low temperature the positive ΔH dominates, giving ΔG>0 (non-spontaneous). When T exceeds ΔH/ΔS, the –TΔS term outweighs ΔH, ΔG becomes negative and the process proceeds spontaneously. Trap: if ΔS is not positive, an endothermic reaction can never become spontaneous by heating.

Memory trick: Endo-entropy: spontaneous only if T > ΔH/ΔS

Temperature dependence of spontaneity — common mistake

“A freezer keeps chocolate solid even though melting is endothermic with ΔS>0; low T makes ΔG positive, not just slows the melt.”

Many think spontaneity is fixed, but ΔG varies with temperature. When ΔH and ΔS share the same sign, ΔG can change sign: ΔH>0, ΔS>0 → spontaneous only at high T; ΔH<0, ΔS<0 → spontaneous only at low T. If ΔH and ΔS have opposite signs, ΔG never flips. Kinetic slowdown does not alter ΔG’s sign.

Memory trick: Same-sign ΔH,ΔS → ΔG can flip; opposite signs → ΔG fixed.

First law for chemistry

“Like a battery powering a flashlight, a chemical system’s energy change equals heat absorbed plus work done.”

ΔU = q + w for any process. At constant pressure with only PV work, the heat exchanged q_p equals the enthalpy change ΔH. The trap: ΔH = q_p fails if non-PV work (e.g., electrical) is present or pressure isn’t constant.

Memory trick: ΔU = q + w; at P const, q = ΔH (PV-work only).

First law for chemistry — common mistake

“Mixing soda in a sealed bottle shows pressure rise; you can’t treat heat at constant volume the same as at constant pressure.”

A frequent mistake is to set q_p = q_v. In reality ΔH = q_p = ΔU + Δ(PV); at constant volume q_v = ΔU. Forgetting the Δ(PV) term gives wrong sign, especially for gases.

Memory trick: q_p ≠ q_v; ΔH = ΔU + Δ(PV).

Bond enthalpy estimate

“Breaking a LEGO tower needs effort, rebuilding releases it – just like breaking chemical bonds costs energy and forming bonds gives it back.”

Using average bond enthalpies, ΔH_rxn ≈ Σ D(broken) – Σ D(formed). This works when all bonds are similar; resonance or multiple bond types can make the estimate off by tens of kJ mol⁻¹.

Memory trick: ΔH ≈ Σ(broken) – Σ(formed).