What Dissolves in Water and What Doesn’t?

Substances that carry an electrical charge or have regions of partial charge tend to dissolve in water, while those that lack any charge separation generally do not. Water itself is a polar molecule, meaning one end is slightly positive and the other slightly negative, and it preferentially pulls apart and surrounds other molecules that share that electrical character. This is why table salt vanishes in a glass of water but cooking oil sits on top. The full picture, though, involves more than just “polar dissolves, nonpolar doesn’t,” because plenty of substances break that simple rule in interesting ways.

Why Water Is So Good at Dissolving Things

Water earns its reputation as the “universal solvent” not because it dissolves everything, but because it dissolves more different kinds of substances than most other liquids. The key is its shape. Each water molecule has two hydrogen atoms bonded to one oxygen atom at an angle, creating a lopsided distribution of electrical charge. The oxygen side pulls electron density toward itself, becoming slightly negative, while the hydrogen side becomes slightly positive. This polarity lets water molecules orient themselves around charged or partially charged particles, effectively prying them apart from whatever solid, liquid, or gas they came from.

When you drop a crystal of table salt into water, the slightly negative oxygen ends of water molecules crowd around each positively charged sodium ion, while the slightly positive hydrogen ends surround each negatively charged chloride ion. This tug-of-war between water molecules and the crystal lattice is strong enough to pull the crystal apart ion by ion. The dissolved ions end up surrounded by a shell of water molecules, stabilized in solution. Any substance that can form similarly favorable interactions with water, whether through full charges or partial charges, has a good chance of dissolving.

What Dissolves Easily

The most straightforward category is ionic compounds, the salts formed when metals combine with nonmetals. Sodium chloride, potassium nitrate, calcium chloride, and most salts with alkali metals dissolve readily. Their charged ions interact strongly with water’s polar structure, and the energy released when water molecules surround those ions is enough to compensate for the energy it costs to break the crystal apart.

Polar organic molecules are the other big category. Sugars like glucose and sucrose dissolve easily because they are studded with hydroxyl groups, each consisting of an oxygen bonded to a hydrogen. These hydroxyl groups form hydrogen bonds with water, the same type of attraction that holds water molecules to each other. Molecular simulations show that every hydroxyl hydrogen on a sugar molecule forms a strong hydrogen bond with surrounding water molecules at very short distances, while the few methyl groups present on some sweeteners do not interact with water in the same way.

Small alcohols like ethanol and methanol are fully miscible with water for the same reason: their hydroxyl groups bond eagerly to water molecules. Short-chain organic acids like acetic acid (vinegar) dissolve well too. The pattern holds for any small molecule with enough polar groups relative to its overall size.

What Refuses to Dissolve

Oils, fats, waxes, and most hydrocarbons are the classic water-refusers. These molecules are made almost entirely of carbon-hydrogen bonds, which share electrons nearly equally, leaving no meaningful charge separation. Water molecules cannot form favorable interactions with them, so mixing would require breaking water’s own hydrogen-bond network without any energetic payoff.

The thermodynamics of this refusal are counterintuitive. You might expect that adding a nonpolar molecule to water would disrupt water’s structure in a messy, disordered way. In reality, the opposite happens: water molecules near a nonpolar surface become more ordered, forming a tight, cage-like arrangement around the intruder. This increased ordering actually decreases the entropy of the system, which makes the process thermodynamically unfavorable. The result is a driving force that pushes nonpolar molecules together and away from water, a phenomenon known as the hydrophobic effect.¹ This is why oil droplets in water spontaneously merge into larger droplets rather than staying dispersed.

Plastics, rubber, and many synthetic polymers also resist water. Their long carbon backbones are largely nonpolar, and even if they contain a few polar groups, those groups are vastly outnumbered by the hydrocarbon portions of the chain.

When Ionic Compounds Surprise You by Not Dissolving

The general rule that salts dissolve in water has enough exceptions to fill a chemistry reference table. Silver chloride, barium sulfate, and calcium carbonate are all ionic compounds, yet they are famously insoluble or nearly so. The simple explanation taught in many classrooms is that smaller, more highly charged ions hold their crystal lattices together more tightly. But the real picture is more complicated than charge density alone suggests.

Computational studies of ionic compounds show that silver chloride, for instance, has a bond roughly 8 kilocalories per mole stronger than silver chlorate and 12 kilocalories per mole stronger than silver nitrate. Meanwhile, sodium chloride’s bond is considerably weaker than silver chlorate’s. This means the same anion, chloride, can form either a highly soluble salt (sodium chloride) or a nearly insoluble one (silver chloride) depending on its partner. The strength of the specific bond between a given cation and anion, not just the generic charge density, determines whether water can pry the crystal apart.²

In practical terms, this is why solubility rules in chemistry involve memorizing specific combinations rather than applying a single formula. Lead iodide, for example, is barely soluble despite both lead and iodide being fairly large ions. The particular way those ions fit together in the crystal creates a lattice that water cannot efficiently dismantle.

The Puzzle of Cellulose

Cellulose is one of the most abundant organic molecules on Earth, making up the structural walls of every plant cell. Chemically, it is a polymer of glucose, the very same sugar that dissolves so freely in water. Yet cellulose is stubbornly insoluble. You can soak a cotton shirt all day and it will never dissolve.

The reason is structural, not chemical. Cellulose chains are long and flat, and they stack together through extensive hydrogen bonding between neighboring chains. These crystalline regions are extremely stable, and breaking into them requires more energy than water can supply through solvation alone.³ The crystalline form of cellulose always has lower energy than its amorphous (disordered) form, making the crystalline regions especially resistant to dissolution.⁴

What makes this particularly interesting is that even amorphous cellulose, where the chains are not neatly stacked, still resists dissolving. The hydrogen bonding between chains, even in less-ordered regions, is apparently enough to keep water from pulling individual chains free. Dissolving cellulose requires specialized solvents or extreme conditions, which is why industries that process wood pulp into products like rayon or cellophane use powerful chemical baths rather than plain water.

Sugars Versus Cellulose and the Role of Shape

The contrast between glucose and cellulose drives home an important point: chemical composition is only part of the story. Glucose is a small, individual molecule with all its hydroxyl groups exposed and available to interact with water. Each hydroxyl group can form its own hydrogen bond with surrounding water molecules, and the molecule as a whole is small enough for water to completely surround it.

Simulations of sugar molecules in water show that the specific arrangement of hydroxyl groups affects how well different sugars dissolve and how they interact with surrounding water structure. The orientation of a single hydroxyl group, particularly the one at position four on the sugar ring, plays a pivotal role by participating in hydrogen-bond networks that involve other hydroxyl groups on the same molecule.⁵ This kind of cooperative bonding, where one hydrogen bond makes the next one easier to form, helps explain why sugars are so readily hydrated.

Cellulose, by contrast, buries many of those same hydroxyl groups in inter-chain bonds, hiding them from water. The molecule’s size and the tight packing of its chains create a physical barrier that prevents water from accessing the polar groups that would otherwise make it soluble. It is a reminder that solubility depends not just on what atoms are present but on how they are arranged and whether water can actually reach them.

How Dissolved Ions Change What Else Can Dissolve

Adding one dissolved substance to water can dramatically change how well other things dissolve. This is something anyone who has tried to dissolve sugar in heavily salted water has noticed, even if they did not know the chemistry behind it. Dissolved ions interact with water molecules, tying them up in hydration shells and leaving fewer free water molecules available to dissolve other things.

Different ions have different effects, and the ranking of how strongly various ions influence the behavior of dissolved substances has been studied for well over a century. This ordering, known as the Hofmeister series, shows up in a remarkably wide range of phenomena: it affects the solubility of nonpolar molecules, the stability of proteins, the behavior of polymers, and even how enzymes function.⁶ Some ions, like sulfate, are strong “salting-out” agents that reduce the solubility of other substances. Others, like thiocyanate, have the opposite effect and can actually increase solubility.

This matters in everyday life more than you might expect. Salting-out effects are why adding salt to ice water lowers its freezing point and why heavily salted fermentation brines behave differently from plain water. In medicine and biotechnology, choosing the right salt concentration is critical for keeping proteins dissolved and functional, or for deliberately crashing them out of solution during purification.

Proteins and the Biological Solubility Problem

Proteins are large, complex molecules with both polar and nonpolar regions, and their solubility in water is a matter of life and death for every cell. A protein that aggregates and falls out of solution can lose its function or, worse, form toxic clumps. Diseases like Alzheimer’s and sickle cell anemia are fundamentally problems of protein solubility gone wrong.

Not all polar amino acids contribute equally to keeping a protein dissolved. Studies measuring protein solubility under different conditions found a surprisingly wide range of contributions among the amino acids classified as hydrophilic. Aspartic acid, glutamic acid, and serine contribute significantly more to solubility than other hydrophilic amino acids like asparagine, glutamine, or threonine, particularly when the protein carries a high net charge.⁷ This finding has practical implications for bioengineering: if you want to make a protein more soluble, swapping in aspartic acid or glutamic acid at surface positions is a more effective strategy than simply adding any polar residue.

The general principle mirrors what happens with small molecules. Proteins fold so that their nonpolar amino acids are buried in the interior, away from water, while their polar and charged amino acids face outward, interacting with the surrounding solvent. When this arrangement is disrupted by heat, pH changes, or mutations, the protein can expose its hydrophobic core, lose solubility, and aggregate. This is exactly what happens when you cook an egg: heat unfolds the egg-white proteins, exposing their nonpolar regions, and they clump together irreversibly.

Rocks, Minerals, and Geological Timescales

Water’s ability to dissolve things is not limited to the kitchen or the lab. Over geological time, water dissolves rocks, reshaping landscapes and controlling the chemistry of rivers, lakes, and oceans. The process is slow on a human timescale but relentless. Limestone caves form because slightly acidic rainwater (rain picks up carbon dioxide from the atmosphere, forming a weak acid) gradually dissolves calcium carbonate from the bedrock over thousands of years.

The rates at which different minerals dissolve in water follow a well-established hierarchy that mirrors what geologists observe in the field. Experimental studies of silicate mineral dissolution have quantitatively verified the stability series used in sedimentary geology, confirming that the minerals observed to weather fastest outdoors also dissolve fastest under controlled laboratory conditions.⁸ Quartz, for example, is among the most resistant common minerals, which is why sandy beaches persist while softer rocks erode.

One subtle limiting factor in mineral dissolution is the hydration of dissolved carbon dioxide. When CO₂ dissolves in water, it only slowly converts to carbonic acid, the species that actually attacks carbonate minerals. This rate-limiting step means that even when water is in full contact with a soluble mineral, the reaction can be bottlenecked by how fast the dissolved gas converts to its reactive form.

How Temperature Changes the Rules

For most solid substances, raising the temperature increases how much will dissolve. If you have ever made hot tea, you know that sugar dissolves much faster and in greater quantities in hot water than cold. The added thermal energy helps break apart the crystal lattice and gives water molecules more kinetic energy to surround the freed particles.

Gases, however, behave in the opposite direction. Carbon dioxide, oxygen, and nitrogen all become less soluble as water gets warmer. This is why a cold can of soda holds its fizz better than a warm one, and why warmer ocean waters hold less dissolved oxygen, creating stress for marine life in warming climates. The physics behind this reversal has to do with the nature of gas dissolution: dissolving a gas in water is an exothermic process (it releases heat), so adding more heat shifts the balance toward the gas escaping back out.

Pressure matters too, but mostly for gases. Increasing the pressure above a liquid forces more gas into solution, which is exactly how carbonated beverages are manufactured: CO₂ is dissolved under high pressure, and the familiar fizz occurs when you open the container and release that pressure. For solids, pressure has a negligible effect on solubility under normal conditions.

Drug Dissolution and Why Speed Matters

In medicine, it is not enough for a drug to be soluble; it has to dissolve at a rate fast enough to be absorbed by the body. A pill that dissolves too slowly may pass through your digestive tract before enough active ingredient reaches your bloodstream. A drug that dissolves too quickly might deliver a spike of the active compound followed by nothing.

The relationship between a drug’s solubility and how fast it dissolves is broadly linear, but the details involve a boundary layer of saturated solution that forms at the surface of the dissolving tablet. Fresh solvent must continuously replace this layer for dissolution to continue. Analysis of dissolution kinetics shows that resistance at the solid surface itself accounts for roughly a fifth of the total resistance to dissolution under standard testing conditions, with the rest coming from how fast the dissolved drug diffuses away from the surface.⁹

Pharmaceutical scientists use this knowledge to engineer better drug delivery. Grinding a drug into finer particles increases its surface area, speeding dissolution. Coating tablets with polymers that dissolve at specific pH levels can target where in the digestive tract a drug is released. Some poorly soluble drugs are reformulated as amorphous (non-crystalline) solids, which dissolve faster because, much like amorphous cellulose, they lack the stable crystal lattice that slows down the process, but unlike cellulose, their molecules are small enough for water to handle once freed.

The Phosphate Problem in Early Earth Chemistry

One of the more fascinating solubility puzzles has nothing to do with your kitchen or medicine cabinet. It concerns the origin of life itself. Phosphorus is essential for DNA, RNA, and the energy-carrying molecule ATP. But its most common geological form, orthophosphate, is poorly soluble and chemically unreactive under the conditions thought to exist on the early Earth. This creates what researchers call the “phosphate problem”: how did life get access to enough dissolved, reactive phosphorus to build its most critical molecules?

A growing body of research points to reduced phosphorus compounds as the solution. Phosphite, produced when certain iron-phosphorus minerals corrode in water, is far more soluble than orthophosphate. Even more promising are compounds formed when these phosphorus species react with dissolved ammonia under oxygen-free conditions. The resulting nitrogen-phosphorus species are both more soluble and more chemically reactive, capable of attaching phosphorus to organic molecules under mild environmental conditions.¹⁰ Solubility, in this context, was not just a convenience but possibly a gatekeeper that determined which chemical pathways could lead to living systems.