Water Versus Hydrogen Sulfide

Introduction

H₂O and H₂S are both bent group-16 hydrides, yet water's boiling point is far higher at the same pressure. H₂S is heavier, so a simple “larger mass means higher boiling” rule fails. Their different intermolecular interactions provide the stronger explanation.

Core explanation

Water has O–H bonds that create strong conventional hydrogen-bond donor sites and oxygen lone-pair acceptor sites. Water molecules can form multiple O–H···O contacts in a changing liquid network. Separating them into gas requires overcoming many favourable between-molecule interactions. Under 1 atm, water's normal boiling point is about 100°C.

H₂S also has two S–H bonds and a bent molecular geometry, and it has dispersion and some polar interactions. Sulfur is larger and less electronegative than oxygen, so S–H does not create the same strong conventional hydrogen-bond network as O–H. H₂S boils far below room temperature, around -60°C at atmospheric pressure. The exact numerical value varies slightly by source and pressure, but the large separation from water is robust.

Molecular mass still matters. H₂S has a greater molar mass and a more polarisable electron cloud, which tends to strengthen dispersion relative to water. That influence would push H₂S upward in a simple size-only comparison. Yet water's extensive hydrogen bonding outweighs the dispersion advantage in this pair. A good explanation mentions the competing factors rather than pretending H₂S has no attractions.

Their bent geometries mean both can have permanent dipoles, but dipole presence alone does not account for the magnitude of the boiling difference. Hydrogen-bond donor/acceptor quality and number of contacts are central. Also, vaporisation breaks or weakens intermolecular contacts, not the covalent O–H or S–H bonds inside the molecules.

The comparison is qualitative unless a problem supplies enthalpies or vapour-pressure data. One should not calculate a boiling-point difference by subtracting bond energies or counting hydrogen bonds as fixed permanent links. Liquid structures are dynamic, and boiling occurs when vapour pressure equals external pressure.

Step-by-step reasoning

1. Draw bent H₂O and H₂S molecules. 2. Identify common dispersion and possible dipole interactions in both. 3. Mark water's O–H donor and O acceptor sites. 4. Note the absence of an equally strong conventional network in H₂S. 5. Compare boiling at the same external pressure and preserve covalent bonds.

Visual explanation

Draw several water molecules joined by dotted O–H···O contacts and several H₂S molecules with only weaker general attractions indicated. Put a mass arrow favouring H₂S dispersion and a larger network arrow favouring water's boiling temperature.

Real-world analogy

Two groups of people may differ in average body weight, but one group repeatedly links hands while the other mostly stands separately. Separating the first group takes more effort despite its lower total weight per person. The analogy recalls multiple interactions, not literal molecular hand-holding.

Real-world example

Water is liquid under ordinary room conditions, while H₂S is a gas. That observation is a direct practical consequence of their very different boiling temperatures and underlines why chemical structure matters beyond formula mass.

Why?

Why does mass alone fail? Boiling depends on energy needed to separate molecules, and strong, numerous hydrogen-bond contacts can dominate over a heavier molecule's dispersion advantage in this comparison.

Common misconception

“H₂S has a lower boiling point because it has no intermolecular forces.” It has dispersion and polar interactions. It lacks a comparably strong conventional O–H-type hydrogen-bond network.

Worked example

Predict which of H₂O and H₂S is a liquid at 25°C and 1 atm. Water boils near 100°C, so 25°C lies below its boiling temperature and liquid water is stable at ordinary pressure. H₂S boils near -60°C, so 25°C lies well above its boiling temperature and it is gaseous. The structural explanation is water's strong hydrogen-bond network, despite H₂S being heavier and more dispersive.

Quick check

1. Does ordinary boiling of H₂O break its covalent O–H bonds? Answer: No. It overcomes intermolecular attractions among intact water molecules.

Exam focus

Compare at the same pressure and mention both mass/dispersion and hydrogen bonding. Do not claim H₂S has zero intermolecular attraction. Use structural donor/acceptor reasoning rather than a vague “water is polar” answer alone.

Advanced insight

Boiling-point trends within heavier group-16 hydrides can be interpreted with increasing dispersion as molar mass rises. Water sits above a simple extrapolation because hydrogen bonding adds a distinct strong interaction network.

Summary

Water boils much higher than H₂S because its O–H donors and O acceptors support extensive hydrogen bonding. H₂S's greater mass strengthens dispersion but does not compensate. Both have intermolecular attractions, and boiling leaves internal covalent bonds intact.

Practice questions

1. Which has greater molar mass, H₂O or H₂S? Answer: H₂S. 2. Which has the higher normal boiling point? Answer: H₂O, by a large margin. 3. What key intermolecular feature differentiates water? Answer: An extensive network of strong O–H···O hydrogen bonds. 4. Why is “H₂S has no forces” wrong? Answer: H₂S still has dispersion and dipole-related attractions.