Did you know that hydrogen sulfide, H₂S, is almost twice as heavy per molecule as water, H₂O, yet boils at about −60 °C while water boils at about 100 °C at ordinary atmospheric pressure? The larger H₂S molecule has stronger potential London dispersion interactions than a smaller molecule might suggest. But liquid water has something much more influential: an extensive network of intermolecular hydrogen bonds through its strongly polar O–H groups and oxygen lone pairs.
For JC students searching for an H2 Chemistry tutor in Hougang, water versus hydrogen sulfide boiling point, H2O H2S hydrogen bonding, Group 16 hydride boiling points, VSEPR molecular shape, intermolecular forces or A-Level Chemistry 9476 bonding comparisons, this is a classic example of why a formula and molar mass do not tell us every physical property.
I ask students to draw H–O–H and H–S–H, then identify which bonds are polar enough to create the strong, conventional hydrogen-bond interactions taught at H2 level. Both molecules are bent, both have lone pairs on the central atom, and both experience dispersion forces. Yet their condensed-phase attractions differ sharply. The fact that both have a similar two-hydrogen formula is not enough to make their boiling points remotely similar.
The long-term skill is to build an explanation from structure → electron distribution → intermolecular interactions → vapour pressure and boiling temperature. We will work through the two hydrides, compare other Group 16 compounds, identify common exam traps and end with changed practice questions.
At eduKate, our small groups of no more than three learners focus on verbal explanation and independent changed-task checks rather than copying an isolated sentence about hydrogen bonding. Families reading from Hougang should confirm current class venue, available JC lesson times and travel needs directly; this article does not by itself establish a physical classroom in Hougang.
Hougang Chemistry Tutor Library · Official SEAB H2 Chemistry 9476 syllabus
Start with the molecular formulae and masses
Water has the molecular formula:
H₂O.
Its relative molecular mass is approximately:
Mr(H₂O) = 2(1.0) + 16.0 = 18.0.
Hydrogen sulfide has molecular formula:
H₂S.
Its relative molecular mass is approximately:
Mr(H₂S) = 2(1.0) + 32.1 = 34.1.
So H₂S has a greater molar mass and more electrons, which typically increases its London dispersion contribution relative to a similar smaller molecule.
If molar mass were the only factor controlling boiling temperature, we might wrongly expect H₂S to boil higher.
But experiments show the opposite:
Water normal boiling point ≈ +100.0 °C.
Hydrogen sulfide normal boiling point ≈ −60.3 °C.
The difference is about 160.3 °C.
The US National Institute of Standards and Technology lists the normal boiling temperature of H₂S at about 212.87 K, equivalent to −60.28 °C. See NIST Chemistry WebBook hydrogen sulfide phase-change data.
That large difference demands an explanation beyond simply counting atoms.
What does a boiling point measure?
A liquid boils when its vapour pressure equals the external pressure under the relevant conditions.
At normal boiling point, the external pressure is the conventional standard atmospheric reference used for the quoted comparison.
Vapour pressure is related to the tendency of molecules in a condensed phase to escape into the gas phase.
When molecules in a liquid are held together by stronger intermolecular attractions, it generally takes a higher temperature to produce sufficient vapour pressure to reach the boiling condition.
This is why an intermolecular-force comparison is relevant.
A boiling-point question does not primarily ask which substance has the strongest covalent bond within an isolated molecule.
Ordinary boiling changes the distribution of molecules between liquid and vapour; it does not require wholesale cleavage of the H–O or H–S covalent bonds.
What is the difference between covalent and intermolecular bonds?
In H₂O, the bonds connecting hydrogen to oxygen are intramolecular covalent O–H bonds.
Between water molecules, there are intermolecular attractions, including strong hydrogen bonds.
In H₂S, the bonds connecting hydrogen to sulfur are intramolecular covalent S–H bonds.
Between H₂S molecules, there are intermolecular forces including London dispersion interactions and relatively weak permanent dipole–dipole interactions.
When a liquid boils normally, we are mainly overcoming or reorganising attractions between molecules as they enter the vapour phase.
If a student says water boils at 100 °C because “it is harder to break the O–H covalent bond”, the explanation confuses physical phase change with chemical bond dissociation.
A stronger answer states exactly which interactions among molecules account for the difference.
Why water molecules are polar
Oxygen is highly electronegative relative to hydrogen. It attracts electron density in each O–H bond.
The bonding electrons are distributed unevenly, leaving oxygen partially negative and the O–H hydrogen atoms partially positive.
The water molecule is bent, so its two polar O–H bond dipoles do not cancel.
Therefore water has a substantial permanent molecular dipole.
The important combination is:
- Strongly polar O–H bonds.
- A bent shape giving a molecular dipole.
- Lone pairs on oxygen available for appropriate intermolecular interactions.
Those features permit strong hydrogen bonding in ordinary liquid water.
The molecule is not polar merely because there is oxygen somewhere in its formula. Shape and the vector combination of bond dipoles matter.
Why water forms strong hydrogen bonds
In the familiar H2 model, a hydrogen bond involves a sufficiently polarised H attached to a strongly electronegative atom such as oxygen, interacting with an available electron-pair region on another suitable atom.
Water has O–H donors: the hydrogen atoms attached directly to oxygen.
It also has oxygen acceptor sites: the oxygen lone pairs can participate in hydrogen-bond interactions.
One water molecule can therefore both donate and accept hydrogen bonds with neighbouring water molecules, creating a dynamic extended network.
A useful conceptual sketch is:
H–O–H ··· O–H,
where the dotted line is a hydrogen bond involving a partially positive O–H hydrogen and an oxygen acceptor on a neighbouring water molecule.
The dotted line is not a new O–H covalent bond. It represents a strong directional intermolecular attraction.
That network makes it energetically costly for water molecules to escape from the liquid into vapour.
Why H₂S does not form a comparable self-hydrogen-bond network
Sulfur lies below oxygen in Group 16. It is a larger atom and is less electronegative than oxygen.
Its S–H bonds are less polar than O–H bonds. The hydrogen atoms are consequently much weaker conventional hydrogen-bond donors.
In the standard H2 Chemistry description, H₂S does not form the same strong, extensive conventional hydrogen-bond network with itself as water does.
H₂S is still a polar molecule: its bonds are somewhat polar and its shape is bent, so it has permanent dipole interactions. All H₂S molecules also experience London dispersion forces.
It is therefore incorrect to write:
“Hydrogen sulfide has no intermolecular forces.”
The correct comparison is:
“Water forms particularly strong intermolecular hydrogen bonds, whereas liquid H₂S experiences mainly dispersion and relatively weaker dipole interactions without comparable strong self-hydrogen bonding.”
In specialised contexts, sulfur can participate in weaker nonclassical hydrogen-bond-related interactions. That advanced qualification does not overturn the standard explanation for the very different normal boiling temperatures.
The molecular shapes are similar but not identical
Both H₂O and H₂S have a central Group 16 atom bonded to two hydrogen atoms.
The central atom also has two lone pairs in the simple Lewis electron-dot model.
That gives four regions of electron density around the central atom and a bent molecular shape in VSEPR terminology.
Representative measured bond angles are approximately:
H–O–H ≈ 104.5°.
H–S–H ≈ 92.1°.
The angles differ because the central atoms’ sizes, orbital characteristics and bonding distributions differ. The exact measured angle is not obtained merely by saying both have “two lone pairs” and copying the water angle to sulfur.
The key H2 level conclusion is that both molecules are bent, while the electronic differences between O–H and S–H affect polarity and intermolecular interactions.
Shape alone is not enough to predict the boiling-point difference. We must connect shape with bond polarity and hydrogen-bonding ability.
Why H₂S can have stronger dispersion yet boil lower
London dispersion interactions arise from temporary fluctuations in electron distributions and induced dipoles.
A molecule with more electrons and greater polarizability often has stronger dispersion attractions than a smaller, otherwise similar molecule.
H₂S has a larger and more polarizable electron cloud than H₂O.
That gives H₂S an important dispersion contribution.
But boiling point depends on the combined condensed-phase interactions and other thermodynamic factors, not the strength of London forces alone.
Water’s strong, cooperative hydrogen-bond network dominates the comparison and supports its much higher normal boiling temperature.
This explains why a statement such as “larger molecules always boil at higher temperatures” is unreliable without checking the molecular structures.
Molar mass is a useful clue, not a replacement for identifying the relevant intermolecular forces.
What the comparison table should show
| Property | H₂O | H₂S |
|---|---|---|
| Relative molecular mass | About 18.0 | About 34.1 |
| Central atom | Oxygen | Sulfur |
| Molecular shape | Bent | Bent |
| Approximate bond angle | 104.5° | 92.1° |
| Bond polarity | Strongly polar O–H | Less polar S–H |
| Permanent molecular dipole? | Yes | Yes, generally smaller |
| Strong conventional self-hydrogen bonding? | Yes, extensive | Not comparable to water |
| Other intermolecular forces | Dispersion and dipole contributions | Dispersion and dipole contributions |
| Approximate normal boiling point | +100.0 °C | −60.3 °C |
The table provides a compact explanation, but students should be able to describe why the O–H group matters rather than simply memorising the last row.
The wider Group 16 trend
Group 16 hydrides include:
H₂O, H₂S, H₂Se and H₂Te.
Moving down from sulfur to selenium to tellurium, molecular size and polarizability increase. The ordinary London dispersion contribution becomes more important.
The normal boiling temperatures of H₂S, H₂Se and H₂Te generally rise down that sequence as the molecules become larger and more polarizable.
Water, however, sits dramatically higher than an extrapolation of that simpler trend would suggest because of its strong hydrogen bonding.
A student should therefore learn the pattern:
Among H₂S, H₂Se and H₂Te, boiling points broadly increase with larger molecular size and dispersion forces.
H₂O is an anomalously high-boiling member because it forms a strong intermolecular hydrogen-bond network.
This is a useful way to see that two general principles can operate within one family of molecules.
It does not mean hydrogen bonding is the only factor in water’s measured phase behaviour, or that the other hydrides have exactly identical shapes and polarities.
A compact trend table
| Hydride | Approximate relative mass | Main school-level attraction emphasis | General boiling trend |
|---|---|---|---|
| H₂O | 18 | Strong hydrogen-bond network | Much higher than expected from mass alone |
| H₂S | 34 | Dispersion and weaker dipole interactions | Lowest of the four in the ordinary comparison |
| H₂Se | 81 | Increasing dispersion with greater polarizability | Higher than H₂S |
| H₂Te | 130 | Further increase in dispersion | Higher than H₂Se |
The approximate molar masses use rounded atomic masses. Actual normal boiling points also depend on pressure and are measured experimental properties.
If a test supplies numerical data for the hydrides, use those data rather than estimating their exact boiling temperatures from mass alone.
For a readable qualitative account of the structures and boiling trends, see the University of Bristol’s Hydrogen Sulfide: Molecule of the Month.
Why water is a liquid at ordinary room temperature
At a typical room temperature around 25 °C and ordinary atmospheric pressure, water is liquid.
Its boiling point at roughly 1 atm is much higher than room temperature because its condensed-phase hydrogen-bond network creates a strong resistance to vaporisation at that pressure.
Hydrogen sulfide, by contrast, has a normal boiling point far below room temperature. Under typical room-temperature and atmospheric-pressure conditions, pure H₂S is gaseous.
That practical difference follows from the same intermolecular-force comparison.
But it does not mean a single water molecule is “liquid” while a single H₂S molecule is “gas”. Phase is a property of a collection of molecules under stated temperature and pressure, not a label carried permanently by an isolated molecule.
This is a useful bridge from molecular bonding to states of matter.
Why boiling water does not destroy the water molecules
Water vapour still contains H₂O molecules.
If boiling required every O–H covalent bond to break, the vapour could no longer be ordinary water vapour. It would contain different chemical species.
Instead, molecules separate from the liquid and enter the gas phase, with hydrogen bonding and other intermolecular interactions reorganising.
Some hydrogen-bonded clusters may exist under certain gas-phase conditions, but the overall condensed-to-vapour change is a physical phase transition.
The same is true of H₂S during its normal boiling transition: the S–H covalent bonds are not routinely destroyed merely by crossing its boiling temperature.
This is why intramolecular bond dissociation energy should not be confused with enthalpy of vaporisation.
Both involve energy, but they describe different processes.
Why a bigger bond enthalpy does not directly equal a higher boiling point
The O–H covalent bond is generally stronger than the S–H covalent bond.
However, the fact that water boils at a much higher temperature should be explained primarily through intermolecular hydrogen bonding, not solely by comparing the strengths of isolated O–H and S–H covalent bonds.
Breaking an O–H covalent bond is a chemical bond-dissociation process. Converting liquid H₂O into vapour is a phase-change process.
A student can use O–H polarity as a reason that hydrogen bonds are possible, but must not say the boiling process breaks the covalent O–H bonds.
That connection between polar intramolecular bonds and intermolecular hydrogen bonds is the more precise H2 explanation.
What makes a conventional hydrogen-bond donor?
In the familiar school-level definition, an important hydrogen-bond donor contains hydrogen directly attached to a strongly electronegative atom such as N, O or F.
Examples include:
- Water’s O–H bonds.
- Alcohols’ O–H groups.
- Ammonia’s N–H bonds.
- Hydrogen fluoride’s H–F bond.
These structures can give a partially positive hydrogen able to interact strongly with a suitable lone-pair acceptor.
A hydrogen atom bonded to sulfur in H₂S is not an equivalent strong conventional donor in the standard comparison.
The presence of hydrogen in a molecular formula is therefore not sufficient.
Likewise, having an atom with lone pairs is not by itself enough to create an extensive strong self-hydrogen-bond network: both donor and acceptor geometry and electronic properties matter.
Contrast ethanol and dimethyl ether
Consider ethanol and dimethyl ether, both with formula C₂H₆O.
Ethanol, CH₃CH₂OH, contains an O–H bond and can donate hydrogen bonds between its own molecules.
Dimethyl ether, CH₃OCH₃, contains oxygen lone pairs but no O–H donor. It cannot support the same conventional strong self-hydrogen-bond network.
Thus ethanol boils much higher than dimethyl ether even though the molecules have the same molecular formula and Mr.
That is an especially clean example of hydrogen bonding because the molecular masses are equal.
In the H₂O versus H₂S comparison, the molecular masses differ, but the water still boils much higher because of the qualitative difference in molecular association.
Both comparisons teach the same method: first inspect donor and acceptor features, then consider all other forces.
Worked comparison 1: explain the anomaly in two sentences
Exam prompt:
Why does water have a much higher boiling point than hydrogen sulfide even though its Mr is lower?
A strong answer could be:
Water molecules have strongly polar O–H bonds and oxygen lone pairs, permitting extensive strong intermolecular hydrogen bonding. H₂S molecules are larger but lack comparable conventional self-hydrogen bonding, so their weaker overall condensed-phase associations allow a much lower normal boiling temperature.
This directly answers the surprising comparison.
An answer saying only “water has more bonds” is ambiguous, because both molecules have two main intramolecular H–X bonds and the relevant issue is intermolecular attraction.
An answer saying “H₂S is nonpolar” is also inaccurate in the familiar molecular model because H₂S is bent and has a permanent molecular dipole.
Worked comparison 2: a change in external pressure
Suppose a question asks whether water always boils at exactly 100 °C.
The answer is no.
The normal boiling point quoted near 100 °C refers to a conventional atmospheric pressure. Boiling happens when the vapour pressure matches the external pressure.
At lower external pressure, a liquid can reach its boiling condition at a lower temperature. At higher external pressure, the boiling temperature can rise.
Therefore the numerical water-versus-hydrogen-sulfide comparison must be understood as normal boiling points under the same pressure convention.
A student who reports 100 °C as an absolute fixed property independent of pressure is ignoring the thermodynamic definition of boiling.
Worked comparison 3: ratio and scale sanity check
Convert the approximate normal boiling points from Celsius to kelvin.
For water:
100.0 °C ≈ 373.15 K.
For hydrogen sulfide, using the NIST value:
212.87 K ≈ −60.28 °C.
The absolute-temperature ratio is approximately:
373.15/212.87 ≈ 1.75.
That arithmetic is correct as a temperature ratio, but it does not show that water’s hydrogen bonds are exactly 1.75 times stronger than H₂S’s intermolecular attractions.
Boiling temperatures are determined by the free energies of the liquid and vapour phases, and a simple ratio of temperatures is not a direct ratio of bond strengths or vaporisation enthalpies.
This is an important scientific warning: the units and physical meaning of a calculated ratio must match the interpretation you give it.
Does water have just one hydrogen bond per molecule?
Not in the full liquid description.
The O–H groups can donate hydrogen bonds and oxygen can accept hydrogen bonds, allowing a dynamic network in which each molecule may participate in several interactions.
In an idealised tetrahedral network, a water molecule can be described as having capacity to donate two and accept two hydrogen bonds.
The actual average number and arrangements in liquid water depend on temperature, pressure and the definition used for counting interactions.
It is therefore misleading to explain water’s high boiling point by drawing just one dotted line and saying “every water molecule has exactly one hydrogen bond”.
The important idea is networked intermolecular association, not a fixed one-to-one permanent pairing.
Water’s unusual solid structure: a useful extension
In ordinary ice, hydrogen-bond interactions help stabilise an open crystalline arrangement.
That structure contributes to ice having a lower density than liquid water under familiar everyday conditions, allowing ice to float.
This is a distinct property from boiling point, but both arise partly from the unusual intermolecular hydrogen bonding of H₂O.
H₂S solid and liquid phases do not exhibit the same familiar open ice network under corresponding ordinary comparisons.
Do not conclude that every substance that hydrogen-bonds must expand on freezing or have a boiling point of exactly 100 °C. Detailed packing and thermodynamics differ.
The broader H2 lesson is that intermolecular interactions can influence multiple macroscopic properties, but each property has its own equilibrium and structural details.
Common misconceptions and their first repairs
| Student answer | First missing concept | Repair question |
|---|---|---|
| H₂S must boil higher because its Mr is 34 | Ignored hydrogen bonding | What intermolecular interactions can H₂O make that H₂S cannot match? |
| H₂S has no intermolecular forces | Forgot dispersion and permanent dipole interactions | Do all molecules exhibit dispersion? |
| Both are bent, so their boiling points must match | Confused shape with interaction strength | How different are O–H and S–H bond polarities? |
| H₂S is nonpolar because sulfur is below oxygen | Ignored bond polarity and bent shape | Do the bond dipoles cancel completely? |
| Boiling breaks the O–H covalent bonds | Confused physical phase change with chemical decomposition | What molecules are present in water vapour? |
| All hydrogen-containing molecules hydrogen-bond strongly | Ignored the electronegative donor requirement | Which atom is directly bonded to H? |
| The boiling-point ratio equals the bond-strength ratio | Misinterpreted a temperature ratio | Which thermodynamic quantities determine phase equilibrium? |
| Every bent hydride has the water bond angle | Overused simple VSEPR geometry | How do orbital and central-atom properties differ? |
Seven independent practice questions
- Calculate the approximate Mr of H₂O and H₂S.
- State the molecular shape of both using the simple electron-domain model.
- Explain why water forms strong conventional hydrogen bonds with itself.
- Name two intermolecular-force types that still exist between H₂S molecules.
- Explain why H₂S can have a lower boiling temperature despite its greater Mr.
- Is boiling an ordinary covalent bond-breaking process?
- Explain the broad trend in boiling temperatures from H₂S to H₂Se to H₂Te and why H₂O does not follow the simple extrapolation.
Answers
1. H₂O ≈ 18; H₂S ≈ 34. 2. Both bent. 3. Strongly polar O–H donors can interact with oxygen lone-pair acceptors, giving a hydrogen-bond network. 4. London dispersion and permanent dipole–dipole interactions. 5. Water’s strong hydrogen-bonded condensed-phase network dominates the comparison. 6. No; normal boiling separates and redistributes molecules without destroying their molecular identities. 7. Greater size and polarizability strengthen dispersion down the sulfur-to-tellurium sequence, while water has anomalously strong intermolecular hydrogen bonding.
If a student gets the first two correct but struggles with Question 5, they need help moving from molecular structure to intermolecular force, not another set of numerical formula-mass exercises.
A seven-session JC1–JC2 learning route
- Session 1: draw water and hydrogen sulfide with lone pairs and bond dipoles.
- Session 2: identify bent shape and distinguish it from total molecular polarity.
- Session 3: apply the O–H donor and acceptor test to water.
- Session 4: compare conventional hydrogen bonding with dispersion and ordinary dipole interactions.
- Session 5: use the boiling-pressure definition to explain why stronger liquid association can raise boiling temperature.
- Session 6: compare other Group 16 hydrides and identify water as the hydrogen-bonding anomaly.
- Session 7: transfer the reasoning to ethanol versus methoxymethane or to carboxylic acids versus alcohols.
The final changed example matters most. It demonstrates whether students understand which molecular feature changes the dominant intermolecular interaction.
What parents choosing H2 Chemistry tuition near Hougang can ask
Bring a marked intermolecular-force or VSEPR question. Ask the tutor to identify whether the student has confused covalent bonds with intermolecular forces, drawn the wrong molecular shape, labelled H₂S as nonpolar or assumed larger Mr always produces higher boiling temperature.
In a three-student group, each learner should independently explain a different molecular comparison before the tutor reveals the model answers. Private tuition should also finish with a fresh structure.
Confirm syllabus experience, actual teaching venue, availability and travel requirements. A Hougang-focused Chemistry article is not a guarantee of a physical tuition classroom in Hougang.
Why this lesson extends into the wider world
Water’s unusual phase behaviour is a foundation of Earth’s climate, the water cycle and the chemistry of living systems.
Hydrogen sulfide is a very different substance: under ordinary room conditions it is a gas and is also highly toxic. Its physical behaviour cannot be predicted simply by knowing it has two hydrogen atoms like water.
The transferable scientific skill is to infer large-scale properties from molecular structure and interactions while paying attention to pressure, temperature and the limits of the model.
That skill supports later Chemistry, environmental science, chemical engineering and materials study, independently of any promise about grades or admission pathways.
Frequently asked questions
Why does H₂O boil much higher than H₂S?
Water forms a strong hydrogen-bonded network through its O–H groups and oxygen lone pairs, while H₂S does not support comparable conventional self-hydrogen bonding.
Is H₂S heavier than water?
Yes. H₂S has Mr about 34, compared with water’s 18, but greater Mr alone does not determine boiling point.
Is H₂S a nonpolar molecule?
No. It is bent and has a permanent dipole, although its S–H bonds are much less polar than water’s O–H bonds.
Does H₂S have intermolecular attractions?
Yes. It has London dispersion forces and dipole interactions, but not water’s strong extensive conventional hydrogen-bond network.
Why does H₂S boil near −60 °C?
Its liquid-phase intermolecular interactions are much weaker overall than those in water under the normal boiling-point comparison.
Are water’s boiling temperature and bond dissociation energy the same quantity?
No. Normal boiling concerns liquid–vapour equilibrium and intermolecular forces; O–H covalent bond dissociation is a different chemical energy process.
Is this relevant to H2 Chemistry 9476?
Yes. It reinforces molecular shape, bond polarity, intermolecular forces, hydrogen bonding and structure–property explanations.
Continue learning and sources
Hydrogen bonding, ice and water: a foundation guide · Why ethanol boils much higher than dimethyl ether · Why methanoic acid boils higher than ethanol despite equal Mr · Polar bonds and molecular dipole cancellation · NIST hydrogen sulfide boiling-point data · University of Bristol: hydrogen sulfide and water · Hougang Chemistry Tutor Library · eduKate Singapore Chemistry enquiries.
Safety and data note: The boiling points are approximate normal boiling points at ordinary atmospheric pressure. Hydrogen sulfide is an extremely toxic gas; no direct handling, smelling or demonstration is appropriate outside specialised professional controls. This article is entirely an educational discussion of molecular structure and physical properties.