Difficulty: Intermediate | Prerequisites: Intermolecular forces, thermodynamics basics (enthalpy, entropy, Gibbs free energy)
This topic covers the fundamentals of how and why substances dissolve. It builds on your understanding of intermolecular forces and introduces the thermodynamic reasoning behind solubility. If you are coming in cold, make sure you are comfortable with enthalpy, entropy, and the concept of Gibbs free energy before diving in. This material is foundational for everything that follows in solution chemistry, including colligative properties and equilibria.
Solutions are homogeneous mixtures where a solute dissolves in a solvent through intermolecular interactions (solvation). Whether something dissolves, and how much, depends on the thermodynamics of the process: the balance of enthalpy, entropy, and Gibbs free energy. A negative ΔG means the solute will dissolve; temperature and pressure shift solubility in predictable ways depending on whether the process is exothermic or endothermic.
Solution
A homogeneous mixture of two or more substances that combine to form a single phase (usually liquid, but not always).
In simple terms, this means everything is evenly mixed at the molecular level, so you cannot see separate components.
Solvent
The component of a solution present in the largest quantity. If both substances are present in equal amounts, the one most commonly used as a solvent takes that role.
Think of it as the substance doing the dissolving.
Solute
The substance that is dissolved in the solvent.
Think of it as the substance being dissolved.
Solvation
The electrostatic interaction between solute and solvent molecules, involving the breaking of old intermolecular interactions and the formation of new ones.
In simple terms, solvent molecules surround and interact with solute particles, pulling them into solution.
Hydration
Solvation in which water is the solvent.
Think of it as the specific name for solvation when water does the work.
Saturated solution
A solution in which the maximum amount of solute has dissolved at a given temperature, and the dissolved solute is in dynamic equilibrium with undissolved solute.
Think of it as the point where the solvent simply cannot hold any more solute.
Unsaturated solution
A solution that has not yet reached its saturation point, whether dilute or concentrated.
In simple terms, there is still room for more solute to dissolve.
Sparingly soluble salt
An ionic compound that dissolves only minimally in a given solvent.
Think of it as a salt that barely dissolves at all.
Spectator ions
Ions present in solution that do not participate in the chemical reaction. Sodium and nitrate ions are common examples.
In simple terms, they are bystanders in the reaction.
Solutions are homogeneous mixtures that form a single phase
They can be formed from combinations of all three phases of matter, not only solids dissolved in liquids
All solutions are mixtures, but not all mixtures are solutions
A mix of gases, for instance, is not a solution because the gases do not interact much chemically
In a same-phase mixture, the component present in the larger quantity is the solvent
Solutes interact with the solvent through intermolecular forces and can also interact with other dissolved compounds
Solvation is the process of breaking existing intermolecular interactions in both solute and solvent and forming new ones between them
When water is the solvent, solvation is specifically called hydration
Exothermic solvation
Occurs when the new solute-solvent interactions are stronger than the original ones
Favoured at low temperatures
Dissolution of gases into liquids is exothermic
Lowering the temperature of a liquid increases gas solubility (Le Chatelier's Principle)
Endothermic solvation
Occurs when the new interactions are weaker than the original ones
Favoured at high temperatures, because heat must be supplied to break the strong original interactions
Ideal solutions
The enthalpy of dissolution is zero
The strength of new solute-solvent interactions equals the strength of the old ones
Spontaneity depends on both enthalpy change (ΔH) and entropy change (ΔS)
Dissolution can be spontaneous whether it is exothermic or endothermic
At constant temperature and pressure, entropy always increases upon dissolution overall
Worked example: NaCl in water
The process is endothermic (ΔH is a small positive number)
More energy is needed to break NaCl ionic bonds and water's hydrogen bonds than is released when ion-dipole interactions form
Entropy of the NaCl increases when it breaks apart into ions
Entropy of the water decreases because water molecules become ordered around the ions
The net entropy change is still positive (the gain from NaCl breaking apart outweighs the loss from water ordering)
Because ΔS is positive and large enough, ΔG is negative, and dissolution is spontaneous
The maximum amount of a substance that can dissolve in a particular solvent at a given temperature
At saturation, dissolved solute is in dynamic equilibrium with undissolved solute
Adding more solute after saturation causes it to precipitate out
Both dilute and concentrated solutions are considered unsaturated if they have not yet reached the saturation point
Solubility and ΔG
If ΔG is negative: solute is soluble
If ΔG is large negative: very soluble
If ΔG is small negative: soluble but closer to being undissociated
If ΔG is positive: solute is insoluble
Gas solubility
Increased by decreasing temperature (exothermic dissolution)
Increased by increasing partial pressure above the solvent
Aqueous solutions use water as the solvent. Hydronium ions form when an acid reacts with water.
The two rules to memorise first:
All salts containing ammonium (NH₄⁺) and alkali metal cations are water-soluble
All salts containing nitrate (NO₃⁻) and acetate are water-soluble
Additional solubility trends:
Halides (Cl⁻, Br⁻, I⁻) are soluble, except fluorides and halides with Ag⁺, Pb²⁺, and Hg₂²⁺
Sulfate salts (SO₄²⁻) are soluble, except those with Ca²⁺, Sr²⁺, Ba²⁺, and Pb²⁺
Metal oxides are generally insoluble, except those formed with alkali metals, ammonium, and CaO, SrO, BaO
Hydroxides are generally insoluble, except those formed with alkali metals, ammonium, Ca²⁺, Sr²⁺, and Ba²⁺
Carbonates (CO₃²⁻), phosphates (PO₄³⁻), sulfides (S²⁻), and sulfites (SO₃²⁻) are generally insoluble, except those formed with alkali metals and ammonium
Sodium and nitrate ions are usually spectator ions, so they are useful for delivering other ions into solution without participating in reactions.
Gibbs free energy of dissolution:
\Delta G = \Delta H - T \Delta SWhere ΔH is the enthalpy change, T is temperature in Kelvin, and ΔS is the entropy change. A negative ΔG means dissolution is spontaneous.
Solubility rules quick reference:
Ion type | Soluble? | Exceptions |
|---|---|---|
Na⁺, K⁺, NH₄⁺ salts | Always soluble | None |
NO₃⁻, acetate salts | Always soluble | None |
Cl⁻, Br⁻, I⁻ (halides) | Soluble | Except with Ag⁺, Pb²⁺, Hg₂²⁺ |
SO₄²⁻ (sulfate) | Soluble | Except with Ca²⁺, Sr²⁺, Ba²⁺, Pb²⁺ |
S²⁻ (sulfide) | Insoluble | Except with Na⁺, K⁺, NH₄⁺, Mg²⁺, Ca²⁺, Sr²⁺, Ba²⁺ |
O²⁻ (oxide) | Insoluble | Except with Na⁺, K⁺, Sr²⁺, Ba²⁺ |
OH⁻ (hydroxide) | Insoluble | Except with Na⁺, K⁺, Ca²⁺, Sr²⁺, Ba²⁺ |
CO₃²⁻, PO₄³⁻, SO₃²⁻ | Insoluble | Except with Na⁺, K⁺, NH₄⁺ |
Gas solubility decreasing with temperature is why a warm fizzy drink goes flat faster than a cold one. The CO₂ is less soluble in warmer liquid, so it escapes.
The solubility rules are the basis of qualitative analysis in chemistry labs. By adding specific reagents to an unknown solution and observing which precipitates form, you can identify which ions are present.
Students often think that endothermic dissolution cannot be spontaneous. It can. Entropy is the second driver of spontaneity, and a large positive ΔS can make ΔG negative even when ΔH is positive. NaCl dissolving in water is exactly this case.
Students often assume that "concentrated" and "saturated" mean the same thing. They do not. A concentrated solution can still be unsaturated if the saturation point has not been reached.
Students often think dissolving always involves a solid in a liquid. Solutions can form from any combination of phases: gas in liquid, liquid in liquid, or even solid in solid (alloys).
Students often confuse the effect of temperature on gas solubility vs. solid solubility. Gases become less soluble as temperature increases (exothermic dissolution). Most solids become more soluble as temperature increases (endothermic dissolution).
⚠️ Know the first two solubility rules cold: alkali metal and ammonium salts are always soluble; nitrate and acetate salts are always soluble.
⚠️ Be able to determine whether dissolution is spontaneous using ΔG = ΔH - TΔS. Expect a problem where you must evaluate the signs and magnitudes of ΔH and ΔS.
⚠️ Understand the NaCl-in-water example thoroughly. It is the go-to illustration of endothermic but spontaneous dissolution driven by entropy.
⚠️ Know how temperature affects gas vs. solid solubility and be able to explain why using Le Chatelier's Principle.
⚠️ Be ready to predict whether a precipitate forms when two aqueous solutions are mixed, using the solubility rules.
True or false: A concentrated solution is always saturated.
Answer: False. A solution can be concentrated but still unsaturated if it has not reached its saturation point.
Fill in the blank: When ΔG is ______, the solute is soluble.
Answer: Negative.
True or false: Dissolving a gas in a liquid is endothermic.
Answer: False. It is exothermic, which is why lowering temperature increases gas solubility.
Fill in the blank: Solvation in water is specifically called ______.
Answer: Hydration.
True or false: All sodium salts are water-soluble.
Answer: True.
Q: Explain why NaCl dissolves in water despite the process being endothermic.
A: Although ΔH is a small positive number (endothermic), the entropy increase from breaking the NaCl lattice into free ions is large enough that TΔS exceeds ΔH. This makes ΔG negative, so the dissolution is spontaneous.
Q: Will a precipitate form if you mix aqueous solutions of NaCl and AgNO₃? Explain using solubility rules.
A: Yes. Sodium nitrate (NaNO₃) is soluble (sodium and nitrate salts are always soluble). However, silver chloride (AgCl) is insoluble because halides with Ag⁺ are an exception to the general halide solubility rule. AgCl precipitates out.
Q: A student heats a carbonated drink and notices it goes flat. Explain this observation.
A: Dissolving CO₂ gas in water is exothermic. By Le Chatelier's Principle, increasing the temperature shifts the equilibrium away from the dissolved state, reducing gas solubility. The CO₂ escapes from solution.
Q: What is the difference between a saturated and a supersaturated solution?
A: A saturated solution holds the maximum amount of dissolved solute at equilibrium at a given temperature. A supersaturated solution contains more dissolved solute than the equilibrium amount, and any disturbance can cause the excess solute to precipitate.
Q: For a substance with a small negative ΔG of dissolution, describe its solubility behaviour.
A: The substance is soluble but only marginally so. It dissolves, but the equilibrium lies close to the undissociated state, meaning only a relatively small amount will dissolve before saturation is reached.
This connects to intermolecular forces because the type and strength of those forces determine whether solvation is exothermic or endothermic, and therefore whether a substance dissolves in a given solvent.
This connects to thermodynamics (ΔG, ΔH, ΔS) because solubility is fundamentally a question of whether dissolution is thermodynamically favoured. The same framework applies to chemical equilibrium more broadly.
This connects directly to colligative properties (covered in the next set of notes), which depend on the concentration of dissolved particles in solution.
Solutions, solubility, solvation, hydration, dissolution, saturated solution, unsaturated solution, supersaturated, solubility rules, aqueous solutions, ΔG, Gibbs free energy, enthalpy of dissolution, entropy of mixing, Le Chatelier's Principle, precipitate, spectator ions, sparingly soluble salts, ionic compounds in water, solubility table, CHEM 101, general chemistry, Michigan State University