Hydrogen sulfide does form hydrogen bonds, but they are substantially weaker than the hydrogen bonds in water, and the forces holding them together have a different character. For decades, introductory chemistry courses either ignored H₂S hydrogen bonding entirely or stated flatly that it doesn’t occur. The real picture, sharpened considerably by recent spectroscopic and computational work, is more interesting: H₂S molecules interact through genuine hydrogen bonds, but those bonds behave in ways that challenge the textbook rules about what makes a “good” hydrogen bond in the first place.
Why the Textbook Answer Was Too Simple
The classic explanation for hydrogen bonding focuses on electronegativity. Fluorine, oxygen, and nitrogen are the three elements most commonly cited as capable of participating in hydrogen bonds, because they pull electron density away from hydrogen so strongly that a partial positive charge on the hydrogen can interact with a lone pair on a neighboring molecule. Sulfur is less electronegative than oxygen, so by this reasoning, H₂S shouldn’t form meaningful hydrogen bonds. That reasoning isn’t wrong exactly, but it’s incomplete. It treats hydrogen bonding as a binary phenomenon, present or absent, rather than a spectrum of interaction strengths.
The clearest evidence that H₂S does form hydrogen bonds comes from the low boiling point puzzle that chemistry students often encounter. H₂S boils at about −60 °C, which is higher than you’d predict based purely on its molecular weight if only weak van der Waals forces were at play. It doesn’t show nearly the same dramatic elevation that water does (water’s boiling point is anomalously high at 100 °C), but something beyond simple London dispersion forces is clearly going on. That “something” is hydrogen bonding, just much feebler than in water.
What Makes H₂S Hydrogen Bonds Different
Research using high-level quantum chemical calculations has shown that hydrogen bonds to sulfur differ from hydrogen bonds to oxygen in two fundamental ways: strength and geometry. Computational studies at advanced levels of theory found that hydrogen-bonded complexes involving sulfur bases are not only weaker than those involving oxygen but also show a distinct preference for a more perpendicular angle of approach between the hydrogen donor and the sulfur atom. In water, the hydrogen bond tends to line up more or less along the direction of the oxygen’s lone pairs. In H₂S, the approach angle is different, reflecting the way sulfur’s lone pairs are distributed in space.
This geometric difference matters because it changes how clusters of H₂S molecules organize themselves. Water molecules famously form extended networks of hydrogen bonds, creating the tetrahedral coordination that gives ice its open structure and liquid water its unusual properties. H₂S molecules don’t build the same kind of ordered network. Gas-phase infrared spectroscopy of H₂S clusters containing two to six molecules revealed that the S–H···S hydrogen bond networks are characterized by weak cooperativity and a substantial contribution from dispersion forces.1PubMed Central. Gas-Phase Infrared Spectroscopic Investigation of S–H···S Hydrogen Bonds in (H2S)n (n = 2–6) Clusters Cooperativity is the phenomenon where forming one hydrogen bond makes the next one slightly stronger, which is a key reason water’s hydrogen bond network is so robust. H₂S barely shows this effect.
Dispersion Forces Do the Heavy Lifting
One of the most revealing findings from recent computational chemistry is just how different the energy breakdown is for H₂S hydrogen bonds compared to water hydrogen bonds. In water, the electrostatic attraction between the partial positive charge on hydrogen and the partial negative charge on oxygen dominates the interaction energy. The hydrogen bond has a strongly directional character because of this electrostatic component. In H₂S, the balance shifts. Symmetry-adapted perturbation theory calculations on the H₂S dimer showed that the less directional character of its hydrogen bond, compared to water’s, results from a less dominant role of electrostatic interactions and a relatively larger contribution from dispersion forces, tied to sulfur’s greater polarizability.2Nature Communications. On the nature of hydrogen bonding in the H2S dimer
This isn’t a minor technical point. It changes what H₂S hydrogen bonds look like in practice. Because the interaction is more dispersion-driven, H₂S molecules aren’t as locked into specific orientations relative to each other. The hydrogen bond is “floppier” and less geometrically constrained. A separate computational study of H₂S-containing clusters confirmed that in these systems, the interaction energy is more dominated by dispersion and less dictated by highly directional electrostatic forces.3PubMed. Mixed H2S and H2O Clusters─New Insights into Dispersion-Dominated Hydrogen Bonding Another analysis of various dimer geometries found that while electrostatics still contributes, the influence of dispersion becomes quite pronounced in any dimer involving hydrogen sulfide.4PubMed. How Flexible Is the Hydrogen Sulfide Molecule Structure? Influence of Hydrogen Sulfide Molecule Geometry on Its Hydrogen Bonds
This distinction has practical consequences for how we model and predict H₂S behavior. Early density functional theory methods struggled to capture these weak, dispersion-heavy interactions properly.5The Journal of Chemical Physics. Theoretical investigation of weak hydrogen bonds to sulfur More recent dispersion-corrected methods have closed that gap. Benchmark studies now show that the best-performing functionals can reproduce sulfur hydrogen bond energies with errors of just a fraction of a kilocalorie per mole.6PubMed Central. Hydrogen Bond Benchmark: Focal‐Point Analysis and Assessment of DFT Functionals Getting the computational methods right matters because much of what we know about H₂S hydrogen bonds comes from calculations, and a method that can’t handle dispersion will underestimate the interaction.
When H₂S Meets Water
Things get more interesting when H₂S and water molecules mix. In nature and in industrial settings, H₂S rarely exists in pure isolation. It dissolves in water, interacts with water vapor, and forms mixed clusters. Research on mixed pentamers (five-molecule clusters containing both H₂S and water) revealed important structural differences. In the 4:1 and 1:4 hydrogen sulfide-water pentamers, the hydrogen sulfide clusters reorganize to compensate for the weaker sulfur-centered hydrogen bonds.7PubMed Central. Hydrogen Bond Interaction Networks in the Mixed Pentamers of Hydrogen Sulfide and Water In other words, when you mix strong hydrogen-bonding molecules (water) with weak ones (H₂S), the system doesn’t just average out. The H₂S molecules rearrange themselves in ways that partially compensate for their weaker bonding ability.
This mixed-cluster behavior is directly relevant to gas hydrate formation. Gas hydrates are ice-like crystalline solids where water molecules form cage structures that trap guest molecules like methane or H₂S. The hydrogen bonding network of water provides the structural scaffolding, but the guest molecule influences how readily the hydrate forms. Neutron diffraction studies of hydrogen sulfide hydrate have shown that the hydrogen bonding network formed by the water cages changes with temperature, with the distribution of hydrogen atom positions shifting as the hydrate warms or cools.8ScienceDirect (Chemical Physics Letters). Observation of hydrogen bonding network in hydrogen sulfide hydrate using neutron diffraction
In the oil and gas industry, H₂S-containing gas hydrates are a serious pipeline safety concern. Hydrogen sulfide accelerates hydrate formation compared to sweet (H₂S-free) gas, and conventional chemical inhibitors used to prevent hydrate blockages perform poorly in sour gas reservoirs where H₂S is present.9Energy & Fuels. Molecular Insights into the Inhibition Mechanisms of Modified PVCap Kinetic Inhibitors in Hydrogen Sulfide-Containing Gas Hydrate Systems Understanding how H₂S hydrogen bonds interact with water’s hydrogen bond network is part of the effort to design better inhibitors for these conditions.
Sulfur Hydrogen Bonds in Proteins Are Surprisingly Strong
Perhaps the most counterintuitive finding in this field is that sulfur-centered hydrogen bonds in biological molecules can be just as strong as conventional hydrogen bonds involving oxygen and nitrogen. The amino acids cysteine and methionine both contain sulfur, and their side chains participate in hydrogen bonds with the protein backbone. Research combining gas-phase spectroscopy with high-level quantum calculations found that amide N–H···S hydrogen bonds are as strong as classical hydrogen bonds, despite sulfur’s lower electronegativity.10PubMed. Critical Assessment of the Strength of Hydrogen Bonds between the Sulfur Atom of Methionine/Cysteine and Backbone Amides in Proteins The same study concluded that the electronegativity of the acceptor atom is not the sole criterion for predicting hydrogen bond strength, directly challenging one of the most commonly taught rules in chemistry.
Why would sulfur hydrogen bonds in proteins be stronger than what you’d predict from the pure H₂S molecule? Part of the answer is context. In a protein, the sulfur atom in cysteine can exist in different chemical states: as a thiol (with a hydrogen still attached) or as a thiolate (having lost that hydrogen, leaving a negative charge). A study of the enzyme peroxiredoxin AhpE showed that sulfur hydrogen bonding can occur over longer distances than standard conventional hydrogen bonds while still reaching comparable strength. The hydrogen bonds to the thiolate form of the sulfur were both shorter (around 2.34 Å) and stronger than those to the thiol form (around 2.74 Å).11Scientific Reports. Revisiting sulfur H-bonds in proteins: The example of peroxiredoxin AhpE The charged thiolate is a much better hydrogen bond acceptor than neutral sulfur, which makes sense: the extra electron density gives the surrounding hydrogens more to interact with.
This finding has implications beyond academic interest. Cysteine residues play critical roles in enzyme active sites, redox chemistry, and protein folding. If the hydrogen bonds to sulfur in these contexts are stronger than previously assumed, models of protein structure and enzyme mechanisms that underestimate sulfur’s hydrogen bonding capacity may be getting the energetics wrong. Computational studies of protein stability and drug binding increasingly need to account for these interactions accurately.
Detecting H₂S Hydrogen Bonds Experimentally
Measuring hydrogen bonds in H₂S experimentally is harder than measuring them in water, precisely because they are weaker and less directional. Several complementary techniques have been used to pin down their properties. Gas-phase laser spectroscopy allows researchers to study isolated clusters of two or a few H₂S molecules, free from the complications of being surrounded by bulk solvent. Infrared spectroscopy is particularly useful because the stretching frequency of the S–H bond shifts when it participates in a hydrogen bond, just as the O–H stretch shifts in water. Those shifts are smaller in H₂S, consistent with a weaker interaction, but they are clearly detectable in clusters.1PubMed Central. Gas-Phase Infrared Spectroscopic Investigation of S–H···S Hydrogen Bonds in (H2S)n (n = 2–6) Clusters
Matrix isolation spectroscopy provides another window into these interactions. In this technique, H₂S and a partner molecule are trapped in a frozen noble gas matrix at very low temperatures, essentially freezing the complex in place so it can be studied with infrared spectroscopy. Recent experiments used this approach to investigate S–H···S–S hydrogen bonds between H₂S and diethyl disulfide, confirming that these sulfur-to-sulfur hydrogen bonds form reliably under controlled conditions.12PubMed. Matrix Isolation IR Spectroscopic Investigation of S-H···S-S H-Bond in H2S-Diethyl Disulfide Complex These experiments bridge the gap between what calculations predict and what physically happens when H₂S molecules interact.
Hydrogen Bonds Under Extreme Pressure and Superconductivity
In an entirely different context, hydrogen bonds in the hydrogen sulfide system have attracted intense attention from physicists. In 2015, compressed hydrogen sulfide (in the form H₃S at very high pressures) was found to superconduct at about 203 kelvin, which was the highest superconducting transition temperature ever recorded at the time. The hydrogen bonds in this material turn out to be central to the phenomenon.
Under normal conditions, the hydrogen atoms in H₃S sit closer to one sulfur atom than another, forming an asymmetric hydrogen bond. But at extreme pressures, quantum effects cause the proton to delocalize between two sulfur atoms, creating a symmetric hydrogen bond where the hydrogen sits at the midpoint. Calculations showed that this quantum hydrogen-bond symmetrization lowers the pressure needed for the key crystal structure by about 72 gigapascals, and the resulting symmetric phase exists across the entire pressure range where high-temperature superconductivity was measured.13Nature. Quantum hydrogen-bond symmetrization in the superconducting hydrogen sulfide system In effect, the quantum behavior of the hydrogen bond determines which crystal structure is stable, and that crystal structure is what produces the record-breaking superconductivity. The hydrogen bond in this case isn’t a weak intermolecular interaction at all: it’s a covalent-like bond whose symmetry controls an extraordinary material property.
This extreme-pressure context is a world away from the weak interactions between H₂S molecules at atmospheric conditions. But it illustrates how the hydrogen bond concept, when pushed to its limits, connects chemistry that seems straightforward to physics that is anything but. The same proton that barely holds two H₂S molecules together in a floppy gas-phase dimer can, under the right conditions, be the linchpin of a superconducting crystal.
Why Electronegativity Alone Doesn’t Tell the Whole Story
The traditional teaching that hydrogen bonds require highly electronegative atoms like F, O, or N is a useful simplification, but researchers have known for decades that it’s an oversimplification. Sulfur’s hydrogen bonds illustrate why. The strength of a hydrogen bond depends on several factors working together: the charge distribution on the donor and acceptor, the polarizability of the atoms involved, the geometry of approach, and the electronic environment created by neighboring atoms or the larger molecule.
Sulfur is more polarizable than oxygen, meaning its electron cloud deforms more easily in response to nearby charges. This larger polarizability is precisely why dispersion forces contribute more to H₂S hydrogen bonds than to water’s.2Nature Communications. On the nature of hydrogen bonding in the H2S dimer It also means that the local chemical environment around sulfur can dramatically modulate its hydrogen bonding ability. A neutral sulfur in free H₂S gas is a poor hydrogen bond acceptor. A thiolate sulfur in a protein active site, carrying extra electron density and flanked by polar backbone groups, can accept hydrogen bonds that rival those of oxygen.10PubMed. Critical Assessment of the Strength of Hydrogen Bonds between the Sulfur Atom of Methionine/Cysteine and Backbone Amides in Proteins Same element, very different behavior.
This is the takeaway that most textbook treatments miss. Asking whether H₂S “forms hydrogen bonds” as a yes-or-no question doesn’t capture what’s actually happening. The more useful framing is that sulfur can participate in hydrogen bonds across a wide range of strengths, from barely detectable in the gas-phase dimer to biologically significant in the interior of a folded protein. Where any particular sulfur hydrogen bond falls on that spectrum depends far more on the surrounding molecular context than on sulfur’s position in the periodic table.