how water polarity enables hydrogen bonding

Question

Describe how the polarity of water molecules results in hydrogen bonding. Draw 3 water molecules and use a dotted line to show where hydrogen bonds are formed.​

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Answer

Chemistry concept with clear spatial structure — building a 3D interactive model.

The 3D model is ready — drag to rotate it, then slide the O···O distance control and watch the blue dotted hydrogen bonds fade and break.

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What you're looking at

Each of the three water molecules is drawn ball-and-stick: a red oxygen, two white hydrogens, joined by solid grey rods — those rods are the covalent O–H bonds inside a molecule. The blue dotted lines are hydrogen bonds, which form between molecules.

Step 1 — Why each O–H bond is polar

Oxygen's electronegativity is $3.44$, hydrogen's is $2.20$. Because O pulls the shared electron pair much harder, the electrons in the O–H bond spend more time near O. So the bond becomes polar covalent:

The charge is only shifted, not transferred — so this is not an ionic bond. Oxygen ends up partially negative $({\delta }^{-})$, hydrogen partially positive $({\delta }^{+})$.

Step 2 — Why the whole molecule is polar

Water's H–O–H angle is about $104.5^{\circ }$ — it is bent, not linear. Each O–H bond has its own little dipole, but because they sit at an angle, they don't cancel; they add into one net molecular dipole pointing from the hydrogen side $({\delta }^{+})$ toward the oxygen side $({\delta }^{-})$. That net dipole is what we call a polar molecule. It's why water dissolves salts and why it has such a high boiling point.

Step 3 — Where the hydrogen bond forms

A hydrogen bond is the electrostatic attraction between:

• the ${\delta }^{+}$ hydrogen of one water molecule, and
• a lone pair on the ${\delta }^{-}$ oxygen of a neighbouring water molecule.

Notice in the model how the H sits almost exactly on the straight line joining the two oxygens — real hydrogen bonds are close to linear. In the chain of three molecules you can see two hydrogen bonds. In general each water can donate 2 hydrogens and accept 2 lone pairs, so it can form up to 4 hydrogen bonds — that's the tetrahedral network that makes ice expand and float.

Step 4 — How strong is it?

Interaction	Energy
Covalent O–H bond	$\approx 460\text{kJ/mol}$
Hydrogen bond	$\approx 20\text{kJ/mol}$

A hydrogen bond is roughly 20–25 times weaker than the covalent bond inside the molecule. That's exactly what the slider shows: at the real liquid-water value $2.75\text{A˚}$ the dotted lines are solid; pull past $3.5\text{A˚}$ and they simply vanish — no covalent bonds were broken, only the weak intermolecular attraction.

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• Key concepts

  • Polarity — unequal sharing of electrons $(\Delta {\text{EN}}_{\text{O}-\text{H}}\approx 1.24)$ creates bond dipoles.
  • Bent geometry $(104.5^{\circ })$ — the two bond dipoles fail to cancel, giving a net molecular dipole.
  • Hydrogen bond — attraction between ${\delta }^{+}$ H and a lone pair on ${\delta }^{-}$ O; it is intermolecular.
  • Strength hierarchy — covalent (intramolecular) $\gg$ hydrogen bond (intermolecular), so H-bonds break first.

• Confusion points

  • "Polar bond = ionic bond." No — in a polar covalent bond electrons are shifted (δ⁺/δ⁻); in an ionic bond they are transferred (full +/− charges).
  • "The dotted line is a weak covalent bond." No — a hydrogen bond is a non-covalent attraction between two separate molecules; no electrons are shared between them.
  • "The two O–H bonds add to give a big dipole regardless of shape." Only the bent shape leaves a net dipole; a linear arrangement would make the two bond dipoles cancel exactly (this is why ${\text{CO}}_2$ is non-polar).
  • "Hydrogen bonds form between two hydrogens." No — always H(δ⁺) to O's lone pair; the H···H side is repulsive.

Understanding check: Carbon dioxide, ${\text{CO}}_2$, has two polar C=O bonds, yet the molecule is non-polar. In one or two lines, explain why ${\text{CO}}_2$'s polarity cancels but water's does not — and therefore why ${\text{CO}}_2$ cannot hydrogen-bond.