What is the difference between solubility and rate of dissolving?

Solubility is how much can dissolve — a maximum amount at a given temperature, measured in grams per 100 g of water. Rate of dissolving is how fast it gets there. Stirring, grinding the solid into a powder, and heating all speed up the rate by increasing contact between solute and solvent particles, but only temperature actually changes the maximum amount.

How can I tell if a solution is saturated or supersaturated just by looking at it?

Look for undissolved solid. A saturated solution that has extra solute added will have solid sitting at the bottom that refuses to dissolve no matter how long you stir. A supersaturated solution is completely clear with no solid present, but drop in a seed crystal or scratch the glass and crystals will form immediately. That reaction to a disturbance is the reliable test.

Why does solubility get reported per 100 g of water instead of per liter?

Using a fixed mass of solvent keeps the comparison fair, since the volume of water changes slightly with temperature while its mass does not. The 100 g basis also makes scaling easy: for 50 g of water halve the value, for 300 g triple it. Concentration measures like molarity, which you will use next, do go by volume of solution.

Does pressure affect how much sugar dissolves in water?

Essentially not at all. Solids and liquids are already tightly packed and nearly incompressible, so squeezing the system does not change how much solid can dissolve. Pressure matters only for gases dissolved in liquids, where Henry's law applies — which is why sealed soda holds so much carbon dioxide and open soda loses it.

CHEM-8.1

Solutions & Solubility

Learn why substances dissolve, how to read solubility in grams per 100 g of water, and how to classify unsaturated, saturated, and supersaturated solutions.

What you'll do in this lesson

A voice-first session with the Crimsora tutor on Solutions & Solubility, then targeted practice and FRQs — with the tutor adapting to where you get stuck.

Start Solutions & Solubility free

What this lesson covers

Drop a spoonful of sugar into iced tea and it sits on the bottom; stir it into hot tea and it vanishes. Nothing about the sugar changed — what changed is the outcome of a tug-of-war between particle attractions. Dissolving is not magic and it is not unlimited. Every solute has a measurable ceiling in a given solvent at a given temperature, and chemists report that ceiling in grams of solute per 100 g of water.

In this lesson you will learn to describe dissolving as a competition between solute-solute, solvent-solvent, and solute-solvent attractions; to use solubility values to label a solution as unsaturated, saturated, or supersaturated; and to predict what happens when you heat a solution, cool it, or change the pressure of a gas above it. These ideas set up the concentration calculations you will do next.

Dissolving as a Competition Between Attractions

For a solute to dissolve, three sets of attractions matter. First, solute particles are attracted to each other — in a crystal of

\text{NaCl}

these are ionic bonds holding

\text{Na}^+

and

\text{Cl}^-

in a lattice. Second, solvent particles are attracted to each other — water molecules hold together through hydrogen bonding. Both of those attractions must be partly overcome, which absorbs energy. Third, solute particles are attracted to solvent particles, and forming those new attractions releases energy.

Dissolving happens readily when the third attraction is comparable to or stronger than the first two. Water's partially negative oxygen surrounds each

\text{Na}^+

and its partially positive hydrogens surround each

\text{Cl}^-

; these ion-dipole attractions pay back the energy cost of breaking the lattice apart. The process of solvent particles surrounding solute particles is called solvation, or hydration when the solvent is water.

This competition explains the rule like dissolves like. Polar solvents dissolve polar and ionic solutes because the new solute-solvent attractions are strong. Nonpolar solutes such as oil are held together only by weak dispersion forces, and water molecules would rather hydrogen-bond with each other than surround an oil molecule, so oil and water separate.

A common misconception is that dissolving destroys the solute or turns it into a new substance. It does not. The particles are separated and dispersed, not chemically changed — evaporate the water and the salt returns. Another frequent error is confusing how fast something dissolves with how much can dissolve. Stirring, heating, and grinding into powder all speed up dissolving by increasing contact between particles, but only temperature changes the actual limit.

Solubility Values and the Three Solution Types

Solubility is the maximum mass of solute that dissolves in a fixed amount of solvent at a stated temperature, almost always reported as grams of solute per 100 g of water. Solubility must always come with a temperature attached; a value with no temperature is incomplete.

At the solubility limit, dissolving and crystallizing occur at the same rate. Particles keep leaving the crystal and rejoining it, but the amount in solution stops changing. That dynamic balance defines a saturated solution.

Type	Amount dissolved	Undissolved solid present?	Stability
Unsaturated	Less than the solubility limit	No — more would still dissolve	Stable
Saturated	Exactly the limit	Yes, if extra was added; dissolving and crystallizing are balanced	Stable
Supersaturated	More than the limit	No, but the excess is held precariously	Unstable

A supersaturated solution is made by saturating a solution at a high temperature and then cooling it slowly and without disturbance. The extra solute stays dissolved because there is no surface for crystals to start forming on. Add a single seed crystal or scratch the container, and the excess crashes out until the solution is merely saturated. Hand warmers that snap into a hot solid use exactly this behavior with sodium acetate.

Where students go wrong: assuming that any solution with solid at the bottom is supersaturated. Undissolved solid sitting under a solution is the signature of a saturated solution, not a supersaturated one. A supersaturated solution looks perfectly clear until it is disturbed.

Temperature and Pressure Effects

For most solid solutes, solubility rises as temperature rises. Faster-moving water molecules collide with the crystal more energetically and can pull particles away from the lattice more effectively. Potassium nitrate is dramatic: roughly 32 g per 100 g of water at

20\ ^\circ\text{C}

but about 169 g per 100 g at

80\ ^\circ\text{C}

. Sodium chloride is nearly flat, going from about 36 g to about 38 g over that same range — a useful reminder that "solubility increases with temperature" describes a trend, not a law. A few solids, such as cerium sulfate, actually become less soluble when heated.

Gases behave the opposite way. Heating a solution gives dissolved gas molecules enough kinetic energy to escape into the air above, so gas solubility falls as temperature rises. Warm soda goes flat quickly, and warm river water holds less dissolved oxygen for fish.

Pressure has essentially no effect on solids and liquids dissolved in liquids, because those particles are already tightly packed and nearly incompressible. Pressure matters greatly for gases. Henry's law states that the solubility of a gas is proportional to the partial pressure of that gas above the liquid:

S = k P

. A sealed soda bottle has high-pressure carbon dioxide above the liquid, so a great deal stays dissolved. Open the cap, the pressure drops, and bubbles form immediately.

Change	Solid in water	Gas in water
Raise temperature	Usually increases solubility	Decreases solubility
Raise pressure above liquid	Essentially no effect	Increases solubility

Reading and Using a Solubility Curve

A solubility curve plots grams of solute per 100 g of water on the vertical axis against temperature on the horizontal axis. The curve itself is the set of saturated solutions. Any point below a curve represents an unsaturated solution — more solute could still dissolve at that temperature. Any point above the curve represents more solute than can normally stay dissolved, so it is either supersaturated or, more commonly, a saturated solution with the excess sitting on the bottom as solid.

To use a curve, first read the solubility at the given temperature, then scale it to the actual mass of water. Because the axis is per 100 g of water, working with 50 g of water means halving the value, and 250 g of water means multiplying by 2.5. Skipping this scaling step is the single most frequent error in solubility problems.

To find how much solid crystallizes on cooling, find the mass that can stay dissolved at the lower temperature and subtract it from the mass actually present:

\text{mass crystallized} = \text{mass present} - \left(\text{solubility at }T_{\text{final}}\right)\times\frac{\text{g of water}}{100}

If that difference is negative, nothing crystallizes and the solution is still unsaturated. This kind of controlled cooling is the basis of recrystallization, a standard purification technique: a hot saturated solution is cooled so the desired compound forms pure crystals while impurities, present in much smaller amounts, remain dissolved.

Key terms

Solute.: The substance being dissolved, usually the component present in the smaller amount, such as the salt in salt water.
Solvent.: The substance doing the dissolving, present in the larger amount; water is the most common solvent and gives aqueous solutions.
Solvation (hydration).: The process in which solvent particles surround and stabilize separated solute particles. When the solvent is water it is called hydration.
Solubility.: The maximum mass of solute that will dissolve in a given amount of solvent at a stated temperature, usually expressed in grams per 100 g of water.
Saturated solution.: A solution holding exactly its solubility limit, in which dissolving and crystallizing occur at equal rates.
Unsaturated solution.: A solution containing less solute than its solubility limit, so additional solute would still dissolve.
Supersaturated solution.: An unstable solution holding more dissolved solute than the solubility limit, made by careful cooling; a disturbance makes the excess crystallize out.
Henry's law.: The rule that the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas above the liquid, $S = kP$ .

Worked example

Potassium nitrate has a solubility of about 110 g per 100 g of water at

60\ ^\circ\text{C}

, 64 g per 100 g at

40\ ^\circ\text{C}

, and 32 g per 100 g at

20\ ^\circ\text{C}

. A student stirs 90 g of

\text{KNO}_3

into 200 g of water at

60\ ^\circ\text{C}

. (a) Is the solution unsaturated, saturated, or supersaturated? (b) The solution is cooled to

20\ ^\circ\text{C}

. How much solid crystallizes out?

Step 1: Scale the solubility to the actual mass of water. The table value is per 100 g of water, but the student used 200 g, which is twice as much. At

60\ ^\circ\text{C}

the water can hold

110\ \text{g}\times 2 = 220\ \text{g}

\text{KNO}_3

.

Step 2: Compare to the amount present. Only 90 g was added, and 90 is less than 220, so all of it dissolves and the solution is unsaturated at

60\ ^\circ\text{C}

. You could add another 130 g before reaching the limit.

Step 3: Find the capacity at the new temperature. At

20\ ^\circ\text{C}

the solubility is 32 g per 100 g of water, so 200 g of water can hold

32\ \text{g}\times 2 = 64\ \text{g}

.

Step 4: Subtract. The solution contains 90 g but can only hold 64 g at

20\ ^\circ\text{C}

90\ \text{g} - 64\ \text{g} = 26\ \text{g}

So about 26 g of $\text{KNO}_3$ crystallizes, and the liquid left behind is a saturated solution containing 64 g of dissolved

\text{KNO}_3

.

Check the logic: if the cooling were done very slowly in an undisturbed container, all 90 g might stay dissolved, giving a supersaturated solution. The 26 g answer assumes ordinary cooling where crystallization actually occurs.

Practice questions

A clear solution contains 55 g of a salt dissolved in 100 g of water at

25\ ^\circ\text{C}

. The salt's solubility at

25\ ^\circ\text{C}

is 40 g per 100 g of water. Adding one small crystal causes solid to rapidly form throughout the liquid. Before the crystal was added, the solution was:

Unsaturated
Saturated
Supersaturated
Dilute but saturated

Answer: Supersaturated

The solution held 55 g when the limit is 40 g, so it contained more dissolved solute than the solubility value allows. It stayed clear only because no surface was available for crystals to begin growing on. The added crystal supplied that surface, and the excess 15 g crashed out, leaving a saturated solution. A saturated solution would have held exactly 40 g with any extra sitting undissolved on the bottom, and an unsaturated solution would have held less than 40 g and shown no reaction to the seed crystal.

Explain, in terms of particle attractions, why table salt dissolves readily in water but not in vegetable oil.

Answer: Water forms strong ion-dipole attractions with $\text{Na}^+$ and $\text{Cl}^-$ that compensate for the energy needed to break the ionic lattice and separate water molecules; oil molecules are nonpolar and can offer only weak dispersion forces, which cannot overcome the strong ionic attractions in the lattice, so the salt stays intact.

A complete answer names all three competing attractions. Breaking apart the

\text{NaCl}

lattice costs a large amount of energy because ionic bonds are strong. In water, that cost is repaid because each ion becomes surrounded by polar water molecules oriented by charge — oxygen toward the cation, hydrogens toward the anion. In oil, the only possible solute-solvent attraction is a weak induced-dipole interaction, nowhere near strong enough to pull ions out of the lattice, so the salt remains a solid. This is the reasoning behind 'like dissolves like'.

A sealed bottle of carbonated water is left in a hot car. When opened, it foams over far more violently than a cold bottle would. Give two reasons based on gas solubility.

Answer: Heating lowers the solubility of carbon dioxide in water, so more gas has already come out of solution and built up pressure in the headspace; and opening the bottle drops the pressure above the liquid, which by Henry's law further reduces how much gas can stay dissolved, releasing it rapidly.

Gases behave opposite to most solids with temperature: warmer molecules have more kinetic energy and escape the liquid more easily, so warm carbonated water holds less dissolved

\text{CO}_2

and the trapped gas raises the internal pressure. The second factor is pressure. Henry's law,

S = kP

, says solubility is proportional to the gas pressure above the liquid, so releasing the cap causes an immediate drop in the amount of gas the liquid can hold. Both effects push

\text{CO}_2

out at the same moment, producing the foam.

FAQ

What is the difference between solubility and rate of dissolving?: Solubility is how much can dissolve — a maximum amount at a given temperature, measured in grams per 100 g of water. Rate of dissolving is how fast it gets there. Stirring, grinding the solid into a powder, and heating all speed up the rate by increasing contact between solute and solvent particles, but only temperature actually changes the maximum amount.
How can I tell if a solution is saturated or supersaturated just by looking at it?: Look for undissolved solid. A saturated solution that has extra solute added will have solid sitting at the bottom that refuses to dissolve no matter how long you stir. A supersaturated solution is completely clear with no solid present, but drop in a seed crystal or scratch the glass and crystals will form immediately. That reaction to a disturbance is the reliable test.
Why does solubility get reported per 100 g of water instead of per liter?: Using a fixed mass of solvent keeps the comparison fair, since the volume of water changes slightly with temperature while its mass does not. The 100 g basis also makes scaling easy: for 50 g of water halve the value, for 300 g triple it. Concentration measures like molarity, which you will use next, do go by volume of solution.
Does pressure affect how much sugar dissolves in water?: Essentially not at all. Solids and liquids are already tightly packed and nearly incompressible, so squeezing the system does not change how much solid can dissolve. Pressure matters only for gases dissolved in liquids, where Henry's law applies — which is why sealed soda holds so much carbon dioxide and open soda loses it.

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