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Chemistry 2e · Solutions and Colloids

Solubility

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On this page 8 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Key takeaway
  6. Check yourself
  7. Study tools
  8. Sources & references

In 30 seconds

is the maximum amount of solute that can dissolve in a given amount of solvent at a specified temperature — it is the quantitative version of the "like dissolves like" idea from topic 1. When that maximum is reached, the solution is saturated, and any extra solute sits undissolved in dynamic equilibrium with the dissolved particles.

Two variables dominate solubility:

  • Temperature. For most solid solutes, solubility increases with temperature (more kinetic energy and usually endothermic dissolution). For gases, solubility decreases with rising temperature — gas dissolution is exothermic, and heat drives dissolved gas out of solution.
  • Pressure. Pressure has essentially no effect on solid/liquid solubility, but gas solubility is directly proportional to the gas's partial pressure above the solution — : C = k P where C is the gas concentration in solution (mol/L), P the partial pressure (atm), and k the Henry's law constant for that gas–solvent pair at that temperature.

For ionic compounds, solubility can be predicted from a short set of and quantified at equilibrium with the solubility product, Ksp — the equilibrium constant for dissolution:

AgCl(s) ⇌ Ag+(aq) + Cl-(aq),   Ksp = [Ag+][Cl−] = 1.8 × 10-10

Small Ksp → low solubility; converting Ksp into a (mol/L) is a standard calculation.

Why this matters

  • Environmental and biological oxygen: Fish need dissolved O2. Henry's law explains why carbonated drinks fizz when opened (pressure drops) and why warm water holds less oxygen — a real concern in power-plant cooling discharges (thermal pollution) and summer fish kills.
  • Medical applications: Kidney and gallstones form when salts like calcium oxalate or calcium phosphate exceed their solubility in urine — the same equilibrium chemistry as Ksp. Gases used in anesthesia and deep-sea diving obey Henry's law (the bends is dissolved N2 coming out of solution when pressure drops).
  • Industry and pharmacy: Recrystallization purifies chemicals by exploiting temperature-dependent solubility; drug formulation must ensure active ingredients stay dissolved (or dissolve at the right rate) in the body.
  • Exam staple: Solubility rules, Henry's law calculations, and Ksp → molar solubility conversions appear in nearly every general-chemistry course.

The college version

Core Concepts

Saturation and the equilibrium view

  • Unsaturated solution: below the solubility limit — more solute can dissolve.
  • Saturated solution: at the limit — dissolved and undissolved solute are in dynamic equilibrium (rate of dissolving = rate of precipitating).
  • Supersaturated solution: temporarily above the limit; unstable, precipitates when disturbed (e.g., a seed crystal).

Solubility is always reported with a temperature: "the solubility of KNO3 is about 32 g per 100 g water at 20 °C, and about 246 g at 100 °C."

Temperature effects: solids vs gases

  • Most solids: solubility ↑ with temperature (the dissolution is usually endothermic; heating supplies energy and speeds equilibration). Exceptions exist — Ce2(SO4)3 becomes less soluble as temperature rises because its dissolution is exothermic.
  • All gases: solubility ↓ with temperature (gas dissolution is exothermic; warming shifts the equilibrium toward the gas phase). Warm soda goes flat faster; cold water holds more dissolved oxygen.

Pressure: Henry's law for gases

C = k P

Doubling the partial pressure of a gas above a solution doubles its equilibrium concentration in solution. This is why:

  • Soda is bottled under CO2 pressure; opening the bottle drops PCO2 to ~0.0004 atm (air's tiny CO2 partial pressure), so C plummets and the gas leaves — fizz.
  • Scuba divers must ascend slowly: N2 dissolved under high pressure comes out of solution if pressure drops too fast (decompression sickness).
  • Henry's law constants are specific to each gas–solvent pair and temperature (e.g., k ≈ 1.3 × 10-3 mol L-1 atm-1 for O2 in water at 25 °C).

Solubility rules for ionic compounds (general)

These guidelines classify most common salts (memorized by generations of students):

  • Soluble: all nitrates (NO3−); all salts of group 1 metals and ammonium (NH4+); all chlorides, bromides, iodides except those of Ag+, Pb2+, Hg22+; most sulfates except CaSO4, SrSO4, BaSO4, PbSO4, Ag2SO4.
  • Insoluble: most carbonates, phosphates, sulfides, and hydroxides — except those of group 1 metals and NH4+ (and, for hydroxides, the heavier group 2 metals Ca2+, Sr2+, Ba2+ are sparingly soluble).

These rules come from measured equilibria; they summarize behavior and have a few genuine exceptions, so treat them as strong guidelines.

Ksp and molar solubility

For AgCl: Ksp = [Ag+][Cl−] = 1.8 × 10-10. If s = molar solubility (mol/L), then [Ag+] = s and [Cl−] = s, so:

Ksp = s2   ⇒  s = Ksp

For a 1:2 salt like PbCl2 (Ksp = 1.7 × 10-5):

Ksp = [Pb2+][Cl−]2 = s(2s)2 = 4s3

The general rule: always relate the ion concentrations back to s using the dissolution stoichiometry before substituting into Ksp.

How It Works / Step-by-Step Process

Worked example 1: Henry's law — dissolved oxygen in water

Problem. The Henry's law constant for O2 in water at 25 °C is k = 1.3 × 10-3 mol L-1 atm-1. (a) Find the O2 concentration when the gas above the water is pure O2 at 1.00 atm. (b) Find it when the water is open to air, where the O2 partial pressure is 0.21 atm.

Solution (a):

  1. Write Henry's law: C = k P
  2. Substitute: C = (1.3 × 10-3 mol L-1 atm-1)(1.00 atm) = 1.3 × 10-3 mol/L Dimensional check: (mol L⁻¹ atm⁻¹)(atm) = mol/L ✓

Solution (b):

  1. Same formula, new pressure: C = (1.3 × 10-3)(0.21) = 2.7 × 10-4 mol/L
  2. Air-exposed water holds only about one-fifth as much oxygen — the partial pressure, not the total pressure, is what matters.

Worked example 2: molar solubility from Ksp — AgCl

Problem. Ksp(AgCl) = 1.8 × 10-10. Find the molar solubility of AgCl in water.

Solution.

  1. Write the dissolution equilibrium: AgCl(s) ⇌ Ag+(aq) + Cl-(aq)
  2. Define s = molar solubility; then [Ag+] = s and [Cl−] = s.
  3. Write the Ksp expression and substitute: Ksp = [Ag+][Cl−] = s · s = s2
  4. Solve: s = Ksp = 1.8 × 10-10 = 1.3 × 10-5 mol/L
  5. Check the scale: ~1.3 × 10-5 M ≈ 0.002 g/L — genuinely "sparingly soluble."

Worked example 3: molar solubility from Ksp — PbCl2

Problem. Ksp(PbCl2) = 1.7 × 10-5. Find the molar solubility of PbCl2.

Solution.

  1. Dissolution: PbCl2(s) ⇌ Pb2+(aq) + 2Cl-(aq).
  2. With s = molar solubility: [Pb2+] = s, [Cl−] = 2s — the stoichiometry matters (two chlorides per formula unit).
  3. Substitute into the Ksp expression: Ksp = [Pb2+][Cl−]2 = (s)(2s)2 = 4s3
  4. Solve: s3 = Ksp4 = 1.7 × 10-54 = 4.3 × 10-6   ⇒  s = (4.3 × 10-6)1/3 = 1.6 × 10-2 mol/L
  5. Note: PbCl2 is far more soluble than AgCl — but you can only compare Ksp values directly between salts with the same ion stoichiometry; AgCl (s2) and PbCl2 (4s3) need the full conversion.

Common Confusions

Do Not ConfuseWithDifference
Solubility ↑ with temperatureAlways trueTrue for most solids but ALL gases become less soluble when heated
Pressure effect on solids/liquidsPressure effect on gasesHenry's law applies to gases; pressure barely changes solid/liquid solubility
Ksp magnitudeMolar solubility, directlyOnly comparable for identical stoichiometries: AgCl vs PbCl2 need conversion through s
[Cl−]sFor PbCl2, [Cl−] = 2s, not s — forgetting the factor of 2 gives wrong answers
"Insoluble""Does not dissolve at all""Insoluble" salts still have tiny, measurable solubility (their Ksp is just small)
Solubility rulesAbsolute lawsRules are empirical summaries with genuine exceptions (e.g., some sulfates, Ag2SO4) — use them as guidelines
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

Solubility is how much sugar your tea can hold before extra sugar just sits at the bottom. Hot tea holds more sugar than cold tea — most solids are like that. But bubbles (gases) are the opposite: cold soda keeps its fizz, and warm soda goes flat. And if you pressurize a bottle, you force more bubble-gas into the drink — that's why soda is fizzy until you open it.

Key takeaways

  • Solubility = maximum solute in a given solvent at a given temperature; saturated = dynamic equilibrium.
  • Most solids: solubility ↑ with temperature; all gases: solubility ↓ with temperature.
  • Pressure affects only gases: Henry's law C = kP — gas concentration is proportional to partial pressure.
  • Like dissolves like still rules the qualitative prediction: polar/ionic in water, nonpolar in nonpolar solvents.
  • Solubility rules (general): nitrates and group-1/NH4+ salts always soluble; chlorides except Ag+, Pb2+, Hg22+; sulfates except Ba2+, Sr2+, Ca2+, Pb2+, Ag+; carbonates/phosphates/sulfides/hydroxides mostly insoluble except group-1/NH4+.
  • Ksp = equilibrium constant for dissolution; relate ion concentrations to molar solubility s via stoichiometry: AgCl: Ksp = s2; PbCl2: Ksp = 4s3; CaF2: Ksp = 4s3.
  • Small Ksp ↔ low molar solubility; comparing Ksp values directly is only valid for salts with the same ion stoichiometry.
  • Applications: fizz (pressure), thermal pollution (warm water → less O2), kidney stones (supersaturated urine), decompression sickness.

Check yourself

6 review questions from the chapter. Try each one, then open the answer.

  1. Define in equilibrium terms.

    Show answer

    A saturated solution holds the maximum solute at that temperature, with dissolved and undissolved solute in dynamic equilibrium (dissolving rate = precipitating rate).

  2. How does temperature affect gas solubility, and what law quantifies the pressure effect?

    Show answer

    Gas solubility decreases as temperature rises; pressure effects follow Henry's law, C = kP.

  3. Write Henry's law and state what each symbol means.

    Show answer

    C = kP: C = gas concentration in solution (mol/L), P = partial pressure of the gas above the solution (atm), k = Henry's law constant for that gas–solvent pair at that temperature.

  4. Ksp(CaF2) = 3.9 × 10-11. Write the relationship between Ksp and molar solubility s for this salt.

    Show answer

    CaF2(s) ⇌ Ca2+ + 2F−, so [Ca2+] = s, [F−] = 2s, and Ksp = (s)(2s)2 = 4s3.

  5. Give two real-world phenomena explained by Henry's law.

    Show answer

    Fizzing of carbonated drinks when opened (pressure drop) and reduced dissolved oxygen in warm water (thermal pollution); also decompression sickness in diving.

  6. Which is more soluble: a salt with Ksp = 10-5 or one with Ksp = 10-12 (same stoichiometry)?

    Show answer

    The 10-5 salt — larger Ksp (same stoichiometry) means larger molar solubility.

Keep learning

Ready to build on this? Continue to the next lesson.

Study tools & related lessonsKey vocabulary · Related

Key vocabulary

solubility
Max solute that dissolves in a given solvent at a given temperature
saturated solution
At the solubility limit, in dynamic equilibrium with undissolved solute
supersaturated solution
Temporarily holds more solute than equilibrium allows
Henry's law
C = kP: gas concentration ∝ partial pressure above the solution
solubility rules
General guidelines predicting which ionic compounds dissolve
solubility product (Ksp)
Equilibrium constant for a dissolving salt
molar solubility
Solubility expressed in mol/L (s)

Sources & references

  1. openstax.org — Chemistry 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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