Sulfur dioxide dissolves readily in water, and it does so to a much greater extent than many other common gases. At room temperature and standard pressure, roughly 94 grams of SO2 can dissolve in a liter of water, making it one of the more water-soluble gases you are likely to encounter. The dissolution releases energy, which tells us the process is thermodynamically favorable under everyday conditions. But what happens after SO2 hits the water is where the chemistry gets genuinely interesting, because the gas does not simply sit there as dissolved molecules. It reacts.
What Happens When SO2 Meets Water
When sulfur dioxide gas contacts water, the first step is straightforward physical dissolution. The SO2 molecule slips into the liquid and exists briefly as a hydrated species. Thermodynamic measurements put the equilibrium constant for this dissolution at about 1.23 mol per kilogram of water per bar of SO2 pressure, with the process releasing roughly 27 kilojoules of energy per mole dissolved.1PubMed Central. Thermodynamics of Solution of SO2(g) in Water and of Aqueous Sulfur Dioxide Solutions That negative enthalpy means dissolving SO2 in water is exothermic, which has practical consequences for how temperature affects the process.
Almost immediately after dissolving, the hydrated SO2 begins reacting with water molecules. A portion of it forms sulfurous acid, which then loses hydrogen ions to produce bisulfite and, at higher pH, sulfite ions. The result is a pH-dependent equilibrium among three main species: dissolved SO2 itself, bisulfite ion, and sulfite ion.2PubMed. The roles of pH and ionic species in sulfur dioxide- and sulfite-induced bronchoconstriction In acidic conditions (low pH), dissolved molecular SO2 dominates. As the pH climbs toward neutral and above, bisulfite takes over, and in strongly alkaline solutions, sulfite becomes the major form. This is not a minor academic detail. It determines everything from how much SO2 a given body of water can absorb to how quickly that sulfur gets converted into other compounds.
Because each of these dissolved sulfur species behaves differently in chemical reactions, the pH of the water effectively controls the fate of the sulfur. A raindrop at pH 5 and a limestone-buffered stream at pH 8 will handle the same incoming SO2 in very different ways.
How Temperature and Pressure Shift Solubility
The relationship between temperature, pressure, and SO2 solubility is not as simple as the textbook rule “gases dissolve better in cold water.” That rule holds for many gases, and it holds for SO2 under certain conditions, but the full picture depends on whether the SO2 you are dealing with is in its gas phase or has been compressed into a liquid.
When SO2 is gaseous, increasing the pressure pushes more of it into solution, which is intuitive. Temperature, on the other hand, has a comparatively modest effect on solubility under these conditions. But the story flips when SO2 is compressed enough to become a liquid. Under those high-pressure conditions, higher temperatures actually increase the solubility of SO2 in water significantly.3Industrial & Engineering Chemistry Research. Solubility of SO2 in Water from 263.15 to 393.15 K and from 10 to 300 bar: Quantitative Raman Spectroscopic Measurements and PC-SAFT Prediction This reversal matters for anyone working with SO2 under industrial or geological conditions, where pressures and temperatures can be far removed from what you would experience standing outdoors.
For everyday environmental scenarios, though, think of SO2 as behaving like most other gases. A cold lake will absorb more of it than a warm one, all else being equal. Because the dissolution is exothermic, adding heat shifts the equilibrium back toward the gas phase, and warming water up drives some dissolved SO2 back out. This is why cold, rainy climates tend to be especially effective at scrubbing SO2 from the air through natural processes.
Dissolved Salts Change the Picture
Pure water in a laboratory beaker is one thing. The water SO2 actually encounters in the real world, whether in oceans, industrial scrubber solutions, or cloud droplets, contains dissolved salts and other solutes. Those solutes alter the gas’s solubility, sometimes substantially.
Systematic measurements of SO2 absorption in solutions containing common seawater salts, actual seawater, and concentrated brine show that dissolved ions affect how much SO2 the water can take up.4Fluid Phase Equilibria. Absorption of SO2 in different electrolyte solutions, seawater and brine The general trend is that higher salt concentrations reduce gas solubility, a phenomenon sometimes called “salting out.” The ions in solution compete for interactions with water molecules, leaving fewer available to solvate SO2. Seawater, with its roughly 3.5% salt content, absorbs less SO2 per liter than pure freshwater under identical temperature and pressure conditions. Brine is even less accommodating.
This matters for industrial processes that use seawater as a scrubbing medium, and it matters for understanding how much atmospheric SO2 the ocean can actually absorb. Coastal power plants that spray seawater through exhaust stacks to capture sulfur emissions need to account for the reduced absorption capacity compared to freshwater, often compensating with higher flow rates or additional chemical additives like limestone slurry.
Industrial Scrubbing and Why It Works
The high solubility of SO2 in water is the foundation of wet flue gas desulfurization, the dominant technology used worldwide to strip sulfur dioxide from power plant and factory exhaust. The basic concept is elegant: spray water, or more commonly a water-based slurry, into the rising stream of dirty gas, and the SO2 dissolves into the droplets. The cleaned gas exits the top; the sulfur-laden liquid collects at the bottom.
Wet scrubbing systems are widely used across SO2-emitting industries, and their efficiency hinges on the same physical and chemical principles that govern SO2 dissolution in any body of water.5Process Safety and Environmental Protection. Simulation of SO2 absorption and performance enhancement of wet flue gas desulfurization system The two main factors controlling how much sulfur a droplet can capture are the rate at which SO2 molecules cross from the gas phase into the liquid, and the chemistry inside the droplet once they arrive.6Open Physics. Numerical analysis of sulfur dioxide absorption in water droplets The mass transfer coefficient, essentially a measure of how quickly molecules can shuttle across the gas-liquid boundary, governs the speed of absorption. The in-droplet chemistry, particularly the reactions that convert dissolved SO2 into bisulfite and sulfate, controls how much total sulfur a droplet can hold before it becomes saturated.
This is why most modern scrubbers do not use plain water. Adding an alkaline reagent like limestone (calcium carbonate) or lime (calcium oxide) continuously neutralizes the acid formed when SO2 dissolves, keeping the pH elevated and shifting the equilibrium toward bisulfite and sulfite. That shift effectively makes room for more incoming SO2 because the dissolved molecular form is constantly being consumed. The end product is usually calcium sulfate, better known as gypsum, which can be sold as a construction material. A well-designed limestone scrubber can remove upward of 95% of the SO2 from exhaust gas.
SO2 in the Atmosphere and Acid Rain
Outside the controlled environment of an industrial scrubber, the same dissolution chemistry plays out every time SO2 encounters atmospheric moisture. Cloud droplets, fog, and rain are all effective at capturing gaseous SO2, though at wildly different efficiencies depending on conditions.
Experimental work on SO2 absorption by water droplets exposed to trace concentrations of the gas in air (ranging from 10 to 1,000 parts per billion) found that a typical urban cloud could remove between 0.08% and 2.0% of ambient SO2 per hour, depending on temperature and the dissolved metal content of the cloud water.7Atmospheric Environment. An experimental investigation of the absorption of sulphur dioxide by water drops containing heavy metal ions That range is wide for good reason. Trace metals like iron and manganese, common in urban and industrial aerosols, catalyze the oxidation of dissolved sulfur species, converting bisulfite into sulfate. This oxidation frees up capacity for more SO2 to dissolve, so metal-rich cloud water scrubs the atmosphere faster than pristine droplets would.
Hydrogen peroxide, another common atmospheric oxidant, also reacts rapidly with dissolved SO2. Research on this reaction in tiny aerosol particles at realistic humidity levels has shown that the high solute concentrations found in real atmospheric particles significantly speed up sulfate formation compared to what happens in dilute laboratory solutions.8PubMed Central. Fast oxidation of sulfur dioxide by hydrogen peroxide in deliquesced aerosol particles This finding helps explain why sulfate aerosol can build up so rapidly during severe haze events in polluted regions, a process that had long been underestimated by atmospheric models relying on dilute-solution chemistry.
Regardless of the specific oxidant involved, the ultimate product of SO2 reacting with atmospheric water is sulfuric acid or sulfate salts. When these fall to the ground as rain, that is acid rain in the classic sense. The entire acid rain problem, which drove major emissions-reduction legislation in the late twentieth century, is fundamentally a consequence of SO2’s eagerness to dissolve in water and the acidic products that result.
Sulfur Isotopes and Tracing Atmospheric Reactions
Scientists who study atmospheric chemistry have a clever way of figuring out which specific reaction pathway converted SO2 into sulfate in a given sample. Sulfur atoms come in several naturally occurring isotopes, and different oxidation reactions shuffle those isotopes at slightly different rates. Laboratory measurements of these fractionation factors show that aqueous oxidation of SO2 by hydrogen peroxide and by ozone produces characteristic sulfur-isotope signatures. The fractionation turns out to be controlled primarily by the acid-base equilibria of dissolved sulfur species, which is why the two different oxidants produce statistically indistinguishable isotope shifts within experimental uncertainty.9Atmospheric Chemistry and Physics. Sulfur isotope fractionation during oxidation of sulfur dioxide: gas-phase oxidation by OH radicals and aqueous oxidation by H2O2, O3 and iron catalysis Gas-phase oxidation by hydroxyl radicals, by contrast, leaves a different isotopic fingerprint.
This matters because it lets researchers analyze sulfate collected in ice cores, lake sediments, or aerosol filters and work backward to determine whether the sulfur was processed through aqueous chemistry (cloud and fog droplets) or gas-phase reactions. The technique has become a standard tool for understanding how atmospheric sulfur cycling has changed over decades and centuries.
Volcanic Lakes and Natural SO2 Dissolution
Some of the most dramatic natural examples of SO2 dissolving in water occur at volcanic crater lakes, where hot, sulfur-rich gases bubble directly into standing bodies of acidic water. At El Chichón volcano in Mexico and Víti crater in Iceland, researchers have detected both hydrogen sulfide and SO2 being released from lake surfaces, even though SO2 would not normally be expected to survive in liquid water at pH 2 to 3 for very long.10Geophysical Research Letters. Sulfur Degassing From Steam‐Heated Crater Lakes: El Chichón (Chiapas, Mexico) and Víti (Iceland)
The explanation involves a chain of reactions. Hydrogen sulfide gas from subaquatic fumaroles enters the lake and undergoes a series of oxidation steps in the water column. One of the intermediate products is sulfite, which can then be converted to dissolved SO2 and eventually escape back to the atmosphere. At El Chichón, only about 0.2% to 5.4% of the hydrogen sulfide supplied by underwater vents makes it out of the lake as gas. The lake acts as a chemical reactor, absorbing volcanic gases and transforming them before releasing a different mix of sulfur compounds into the air. These crater lakes offer a natural laboratory for studying the same dissolution and reaction chemistry that operates in clouds and scrubbers, just at much higher concentrations and temperatures.
Measuring Dissolved SO2 at Very Low Levels
Detecting sulfur dioxide dissolved in water is straightforward when concentrations are high, as they are in industrial scrubber liquors or volcanic lake water. The challenge comes when you need to measure trace amounts, such as the parts-per-billion levels relevant to environmental monitoring or drinking water quality.
One sensitive approach uses a chemiluminescence reaction between dissolved sulfur species (specifically bisulfite) and potassium permanganate. This method can reliably measure SO2 concentrations as low as 0.8 nanograms per milliliter, with reproducibility in the range of 2 to 6 percent relative standard deviation.11Bunseki Kagaku. Continuous Measurement of PPB-Level Sulfur Dioxide Dissolved in Water by Flow Chemiluminescence Method An important practical detail: dissolved SO2 in slightly alkaline solutions containing formate is stable against oxidation by dissolved oxygen over the time scale of the measurement. Without that stabilizer, the sulfur species would oxidize before you could measure them, leading to falsely low readings. This illustrates a broader challenge in working with dissolved SO2. It is a reactive species, and if your sampling or storage conditions allow it to react with dissolved oxygen or metals before analysis, you will undercount it.
SO2 Solubility on Other Planets
The chemistry of SO2 dissolving in water is not purely a terrestrial concern. Planetary scientists have modeled what would happen if large amounts of volcanic SO2 entered an early Martian atmosphere when the planet still had liquid water on its surface. The results suggest that acid rain could have been a powerful mechanism for removing SO2 from Mars’s atmosphere. Modeling estimates indicate that roughly 40 centimeters per year of precipitation with a pH around 2.0 could have fallen in an early Martian environment, assuming surface temperatures near the freezing point of water and SO2 output consistent with the formation of the Tharsis volcanic province.12Planetary and Space Science. Sulfuric acid aerosols in the atmospheres of the terrestrial planets
This modeled acid rain would have been an efficient sink for volcanic SO2, removing it from the atmosphere quickly enough to prevent the gas from building up a significant greenhouse effect. That matters because some researchers have proposed high SO2 levels as a way to warm early Mars enough to keep water liquid. If acid rain kept stripping the SO2 back out, the greenhouse boost would have been short-lived. The solubility of SO2 in water, in other words, may have helped determine whether ancient Mars stayed warm or froze over. The same chemistry that scrubs smokestacks on Earth could have shaped the climate history of another planet.
Common Misconceptions About SO2 and Water
A few points of confusion come up frequently. One is the assumption that dissolved SO2 is the same thing as sulfuric acid. It is not. When SO2 dissolves, the initial product is sulfurous acid (or more precisely, hydrated SO2 and bisulfite), which is a weaker acid than sulfuric acid. Sulfuric acid forms only after the dissolved sulfur is further oxidized, a step that requires an additional oxidant like hydrogen peroxide, ozone, or catalytic metals. In the absence of those oxidants, dissolved SO2 can sit as bisulfite for extended periods without converting to sulfate.
Another common error is thinking that SO2’s high solubility means water bodies will just soak it all up indefinitely. In reality, absorption reaches equilibrium. Once the water is saturated at a given temperature, pressure, and pH, no more SO2 will dissolve unless something changes. In a natural setting, what typically “changes” is that dissolved sulfur gets oxidized to sulfate or consumed by microorganisms, freeing up capacity for more gas to enter. But in stagnant, acidic water with no oxidants, the absorption quickly plateaus. Industrial scrubbers solve this by continuously removing sulfur-laden liquid and replacing it with fresh alkaline slurry, keeping the system far from equilibrium.
A third misconception involves confusing SO2 with other sulfur gases. Hydrogen sulfide (H2S), the gas responsible for the rotten-egg smell, is far less soluble in water than SO2 and behaves quite differently once dissolved. The two gases share sulfur as an element but have distinct chemistry, different solubility profiles, and different environmental impacts. Lumping all “sulfur pollution” together obscures the specific hazards and remediation strategies appropriate to each compound.