Solubility rules for ionic compounds

Predict whether an ionic compound dissolves in water without opening a Ksp table. The rules on this page are a qualitative shortcut — for anything quantitative (saturation checks, buffer stability, gravimetric yields), reach for a Ksp value instead.

What “soluble” means here

Following Zumdahl 10e §4.5, the working cutoffs used in general-chemistry courses are:

• Soluble. Solubility greater than about 0.1 mol/L in room-temperature water — enough that the compound behaves as a strong electrolyte for stoichiometric purposes.
• Slightly soluble. Between roughly 0.01 and 0.1 mol/L. The compound is written as (aq) in some problems and (s) in others; be explicit when it matters.
• Insoluble. Below about 0.01 mol/L — the compound is treated as a solid precipitate in metathesis reactions.

These cutoffs are conventional, not physical constants: Silberberg 9e Table 4.3 and Petrucci 11e Table 5.1 draw the same boundary lines in slightly different words. Even “insoluble” compounds have a finite Ksp, so “insoluble” here means “precipitates out under typical bench conditions”, not “does not dissolve at all”. Everything below is at 25 °C in pure water; hot water, non-aqueous solvents, and acidic or complexing media all shift the picture.

Always soluble

Compounds containing any of these ions are soluble regardless of the counter-ion:

Mostly soluble — memorize the exceptions

Zumdahl 10e and Silberberg 9e agree on the Ba/Sr/Pb sulfates. They diverge on Ag₂SO₄: Zumdahl calls it slightly soluble (~0.8 g/100 mL cold water, per CRC 97e), while some Silberberg tables list it as insoluble. Treat it as “slightly soluble” and cite Ksp ≈ 1.2×10⁻⁵ when the question is quantitative.

Generally insoluble

Compounds containing these anions are insoluble unless paired with a Rule-1 cation (alkali metal or NH₄⁺):

11 rows total across the three tables above — enough to cover every ionic compound you meet in a first-year metathesis problem.

Order-of-operations tie-breaker

When a compound seems to satisfy two rules — e.g., NaOH is a Rule-1 sodium salt (soluble) and a hydroxide (generally insoluble) — apply this order:

1. Rule 1 (alkali / ammonium) always wins. Anything with Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺, or NH₄⁺ is soluble.
2. “Always soluble” anions win over “generally insoluble” partners. Ba(NO₃)₂, Fe(NO₃)₃, AgNO₃ — all soluble because nitrate is Rule 2.
3. Mostly-soluble exceptions apply only when Rules 1 and 2 don’t. BaSO₄ is insoluble; Na₂SO₄ is soluble.

Rule of thumb: read the cation first. If it’s alkali or ammonium, the compound is soluble — stop reading.

Worked example 1 — is PbSO₄ soluble?

1. Cation Pb²⁺ — not alkali, not ammonium. Rule 1 doesn’t apply.
2. Anion SO₄²⁻ — on the “mostly soluble” list, but Pb²⁺ is a listed exception.
3. PbSO₄ is insoluble (Ksp ≈ 1.6×10⁻⁸ at 25 °C, CRC 97e).

Worked example 2 — precipitation from AgNO₃ + NaCl

Mix aqueous AgNO₃ and aqueous NaCl. Predict the product.

Write the metathesis (double-displacement) and split into ions:

AgNO₃(aq) + NaCl(aq) → AgCl(?) + NaNO₃(?)

Check each product:

• NaNO₃ — Rule 1 (Na⁺) and Rule 2 (NO₃⁻). Soluble → stays as ions.
• AgCl — Cl⁻ is soluble EXCEPT with Ag⁺, Pb²⁺, Hg₂²⁺. Insoluble.

Net ionic equation:

Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

A white curdy precipitate drops out immediately — the classical chloride test in qualitative analysis.

Worked example 3 — Ba(NO₃)₂ + Na₂SO₄

Ba(NO₃)₂(aq) + Na₂SO₄(aq) → BaSO₄(?) + 2 NaNO₃(?)

• NaNO₃ — soluble (Rules 1 & 2).
• BaSO₄ — SO₄²⁻ is normally soluble, but Ba²⁺ is a listed exception. Insoluble.

Net ionic equation:

Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)

BaSO₄ is dense, white, and essentially insoluble (Ksp ≈ 1.1×10⁻¹⁰) — used both as a gravimetric standard and as the X-ray contrast agent for GI studies.

More quick verdicts

Common student trip-ups

1. AgF is soluble, but AgCl, AgBr, and AgI are not. Fluoride is the odd halide out — small, strongly hydrated, and stubbornly happy in water. Never generalize “silver halides are insoluble” without carving out F⁻.
2. Hg₂²⁺ (mercurous, a diatomic cation) and Hg²⁺ (mercuric) behave differently. Hg₂Cl₂ (calomel) is insoluble; HgCl₂ (corrosive sublimate) is a molecular, water-soluble compound. Read the subscript on the mercury.
3. Ba(OH)₂ is soluble even though most Group-2 hydroxides are not. Sr(OH)₂ is also soluble at room T; Ca(OH)₂ is only slightly soluble (limewater). Down Group 2, hydroxide solubility rises — the opposite of the Group-2 sulfate trend.
4. “Slightly soluble” is not the same as “insoluble”. CaSO₄, Ag₂SO₄, PbCl₂, and Ca(OH)₂ dissolve enough to matter in saturated solutions and in Ksp problems, even if a precipitation-prediction question calls them “insoluble”.
5. “Insoluble” here is not zero solubility. Every listed insoluble salt has a finite Ksp. For anything quantitative — buffering, gravimetric yield, environmental modeling — use Ksp, not this table.
6. PbCl₂ is slightly soluble in cold water and freely soluble in hot water. Precipitation of lead(II) chloride is temperature-dependent — a mistake that also derails qualitative-analysis schemes if you filter at the wrong temperature.

See also

• Molarity calculator — convert precipitation predictions to mol/L limits.
• pH calculator — many “insoluble” hydroxides and carbonates dissolve in acid.
• Normality calculator — stoichiometry for acid–base titration with precipitating counter-ions.
• Polyatomic ions cheat sheet — formulas and charges for the anions on this page.
• Common Ka / Kb values — pairs well with the “insoluble in water, soluble in acid” story.

Sources

• Zumdahl & Zumdahl, Chemistry, 10th ed. (Cengage, 2018). §4.5 and Table 4.1 — solubility rules for ionic compounds in water. Primary source for this cheat sheet.
• Silberberg & Amateis, Chemistry: The Molecular Nature of Matter and Change, 9th ed. (McGraw-Hill, 2021). Table 4.3. Used to cross-check exceptions; disagrees with Zumdahl on Ag₂SO₄ (“slightly soluble” vs “insoluble”) — resolved as “slightly soluble” here, in line with CRC quantitative data.
• Petrucci, Herring, Madura & Bissonnette, General Chemistry: Principles and Modern Applications, 11th ed. (Pearson, 2017). Table 5.1 — second cross-reference on halide, sulfate, and hydroxide exceptions.
• CRC Handbook of Chemistry and Physics, 97th ed. (CRC Press, 2016). “Aqueous Solubility of Inorganic Compounds at Various Temperatures” — Ksp and quantitative solubilities (BaSO₄, PbSO₄, CaF₂, Ca(OH)₂, Ag₂SO₄, PbCl₂).