2.3.16 · D5Chemical Bonding
Question bank — Hydrogen bonding — intermolecular, intramolecular; consequences (boiling points, water density)
Before we start, three pieces of notation and geometry that every trap below leans on. Learn these first — the pictures do the heavy lifting.
Now the two conditions that define a hydrogen bond, because most traps break exactly one of them:
- Donor: an H covalently bonded to a small, highly electronegative atom (F, O, N) — see Electronegativity.
- Acceptor: a lone pair on another small, highly electronegative atom (F, O, N).
Both must hold. Drop either and you fall back to ordinary Dipole-dipole interactions or Van der Waals forces.
Two key structures you will be quizzed on — the 6-membered ring in o-nitrophenol and the 4-fold tetrahedral coordination of water — are sketched in figure s03; keep it open as you work.
True or false — justify
Every H-bond is just a strong dipole–dipole force with a special name.
Mostly true — it is electrostatic like a dipole–dipole force, but with extra closeness (bare proton, no inner electrons), strong angular directionality (≈180°), and a small covalent charge-transfer character, so it is best treated as its own category, not a mere rename.
A hydrogen bond is stronger than the O–H covalent bond inside a water molecule.
False — H-bonds are ~5–40 kJ/mol while the O–H covalent bond is ~460 kJ/mol; the H-bond is roughly ten times weaker, which is why water evaporates without decomposing.
Chloroform (CHCl₃) forms hydrogen bonds because it contains hydrogen and electronegative chlorine.
False as written — Cl is too large/diffuse to be a good acceptor, but CHCl₃'s C–H (acidified by three Cl) can weakly donate to strong acceptors like acetone; it is a borderline weak H-bond, not a textbook F/O/N one.
Intramolecular hydrogen bonding always raises a compound's boiling point.
False — an internal H-bond satisfies the molecule to itself, so it clings less to neighbours and boils lower than its intermolecular isomer (e.g. o- vs p-nitrophenol).
Ice floats because cold things are lighter than warm things.
False — mass is unchanged; the open tetrahedral H-bond lattice traps empty hexagonal holes, so the same mass occupies more volume, lowering density below that of liquid water.
HF has the highest boiling point of all hydrides because fluorine is the most electronegative element.
False — HF has the strongest individual H-bond but only one H to donate, so it forms chains (~2 bonds/molecule); water forms ~4 and boils higher, so H₂O > HF.
A single water molecule can participate in four hydrogen bonds.
True — it carries 2 O–H donors and 2 lone-pair acceptors, giving 4 "hands," which is exactly what builds ice's 3-D network (see Water — anomalous properties).
The strength of one H-bond in a big cluster equals its strength in an isolated pair.
False — H-bonds are cooperative: when H₂O acts as an acceptor it becomes a better donor, so bonds in a network reinforce each other and each is stronger than in an isolated dimer. This cooperativity underlies water's anomalies and the stability of protein and DNA structures.
Every solid is denser than its own liquid.
False — water is the famous exception because of its directional open lattice; most substances do contract on freezing, which is why this trap feels true.
Because HCl is polar, it should hydrogen-bond about as well as HF.
False — polarity alone isn't enough; the acceptor atom must be small so its lone pair is concentrated and reachable. Large Cl fails the size test, so HCl only does Dipole-dipole interactions.
Hydrogen bonding is what makes hydrocarbons insoluble in water.
False — it's the absence of H-bonding sites on hydrocarbons; they can't join water's network, so "like dissolves like" pushes them out (see Solubility and 'like dissolves like').
Spot the error
"NH₃ boils higher than H₂O because N has more lone pairs available."
Wrong ordering and wrong reason — water actually boils higher; NH₃ has 3 lone pairs but only 1 lone pair and 3 donor H's, limiting it to ~1 effective H-bond per molecule, so H₂O > HF > NH₃.
"o-nitrophenol is more soluble in water than p-nitrophenol."
Backwards — the ortho isomer locks its OH into an internal 6-membered ring (figure s03), leaving little to bond with water, so it is less water-soluble; para bonds outward to water and is more soluble.
"Water's density keeps rising all the way down to 0 °C."
Wrong — density peaks at 4 °C; below that, open ice-like clusters begin forming and expand the liquid, so density falls from 4 °C down to 0 °C.
"Acetic acid dimers form through one hydrogen bond."
Wrong count — the cyclic dimer uses two O–H···O bonds simultaneously, which is why the apparent molar mass in non-polar solvents nearly doubles.
"In the DNA double helix, A pairs with T using three hydrogen bonds."
Wrong number — A–T uses two H-bonds and G–C uses three; the G–C pair is therefore harder to separate (see DNA structure).
"H₂S boils lower than H₂O because sulfur is bigger, giving weaker London forces."
Half-wrong — bigger S gives stronger London forces, which alone would raise its b.p.; H₂S boils lower because it can't hydrogen-bond, so it lacks the huge network-breaking energy that lifts water (see Boiling point trends of hydrides).
"Ethanol and dimethyl ether have the same formula (C₂H₆O), so they boil at the same temperature."
Wrong — ethanol has an O–H donor and boils ~78 °C; the ether has no O–H, cannot donate an H-bond, and boils ~−24 °C. The donor H, not the formula, decides.
"A hydrogen bond bent to 90° is just as strong as a straight one."
Wrong — H-bonds are directional; strength is greatest near 180° (X–H aimed straight at Y's lone pair) and drops sharply as the angle bends, as figure s02 shows.
Why questions
Why must the hydrogen atom, not just any atom, sit between the two electronegative atoms?
Because H uniquely has no inner electrons; once its lone electron is pulled toward X, its bare proton behaves like a point charge, letting the acceptor's lone pair approach unusually close, and Coulomb attraction () grows sharply at small distance.
Why is a 6-membered intramolecular ring (as in o-nitrophenol) so favourable?
The ortho geometry places the O–H hydrogen and an –NO₂ oxygen at just the right distance and near-straight angle to close a strain-free ring (figure s03); meta and para positions are too far to reach.
Why does water insulate fish through winter rather than freezing them solid from the bottom up?
The densest water (4 °C) sinks, so the coldest ice forms and floats on top, capping the lake; the layer beneath stays liquid, letting fish survive (see Water — anomalous properties).
Why does water have unusually high surface tension and viscosity?
Its 3-D, cooperative H-bond network resists both expansion of the surface and flow past neighbours; you must repeatedly break bonds to move or stretch the liquid.
Why do sugars and glucose dissolve readily in water while oils do not?
Sugars carry many –OH groups that H-bond to water and slot into its network; oils have only C–H and C–C, offering no H-bond sites, so they are excluded.
Edge cases
Can a C–H bond ever act as a hydrogen-bond donor?
Rarely and weakly — only when nearby electron-withdrawing groups make the H sufficiently (recall = a partial charge), e.g. CHCl₃ or HC≡CH; it is a real but marginal effect, not the standard F/O/N case.
What happens to H-bonding in water vapour (steam)?
The molecules are far apart and moving fast, so the H-bond network is essentially broken — that broken-bond energy is exactly the large enthalpy of vaporisation that boiling had to pay.
Does the noble-gas-like atom Ne, or a fully bonded N with no lone pair (as in ), act as an acceptor?
No — Ne isn't electronegative-and-lone-pair-donating in the required way, and 's nitrogen has no lone pair left, so neither can accept an H-bond; a lone pair on the acceptor is mandatory.
Between two molecules that can only donate (all their H's) but have no lone pairs, can an H-bond form?
No — a hydrogen bond needs both a donor H and an acceptor lone pair; a system with donors but zero acceptors cannot complete the X–H···Y linkage.
Does a single isolated H-bonded pair capture the full strength seen in bulk water?
No — because of cooperativity, a bond embedded in a network is stronger than the same bond in an isolated dimer, so lab measurements on gas-phase pairs underestimate the effective strength inside liquid water or ice.
At the exact instant liquid water is at 4 °C, why is density maximal rather than at 0 °C?
It is a tug-of-war: warming from 4 °C expands water normally (density down), while cooling below 4 °C builds open ice-like clusters (also density down), so the crossover peak sits at 4 °C.
Recall One-line self-test before you leave
State the two conditions for a hydrogen bond, and one consequence each of intermolecular vs intramolecular bonding. Answer ::: Donor = H on small electronegative F/O/N; acceptor = lone pair on F/O/N. Intermolecular → higher b.p. (molecules cling to each other); intramolecular → lower b.p. (molecule satisfies itself).