Salt Hydrolysis pH Calculator

Salt Hydrolysis pH Calculator: Find Out If a Salt Solution Is Acidic, Basic, or Neutral

This free salt hydrolysis pH calculator is built for chemistry students, teachers, and anyone studying acid-base chemistry who wants a fast, reliable way to find the pH of a salt solution — without memorizing a different rule for every possible combination of acid and base. In simple words, when a salt dissolves in water, its ions don't always sit there quietly. Sometimes one of them reacts with water itself, in a process called hydrolysis, and that reaction can push the final solution's pH away from neutral 7, either toward acidic or toward basic.

This calculator covers every case you'll run into in a general chemistry course: salts made from a strong acid and a strong base (which stay neutral), salts made from a weak acid and a strong base (which turn basic), salts made from a strong acid and a weak base (which turn acidic), and the trickier case of a salt made from a weak acid and a weak base, where both ions react with water at the same time. Pick a built-in salt preset like sodium acetate or ammonium chloride, or enter your own Ka and Kb values, and get the exact pH along with a full, easy-to-follow written solution.

Beyond the plain pH number, this tool also reports percent hydrolysis (how much of the salt's ion actually reacted with water), the hydrolysis constant used in the calculation, and a visual pH scale so you can instantly see where the result lands compared to everyday substances like milk, vinegar, and baking soda.

What Is Salt Hydrolysis?

Salt hydrolysis is what happens when the ions produced by a dissolved salt react with water molecules, shifting the solution's pH away from 7. It's a natural extension of acid-base chemistry: every salt is technically the product of an acid reacting with a base, and the strength of that original acid and base decides whether the resulting ions are 'happy' staying as they are in water, or whether they grab or release a hydrogen ion.

The key idea to remember is this: the conjugate of a strong acid or strong base is chemically unreactive in water. Chloride ion (from strong HCl) doesn't react with water at all, and neither does sodium ion (from strong NaOH). But the conjugate of a weak acid or weak base is still 'reactive' — it retains some of the acid-base character of its weak parent, and it will pull at or donate a hydrogen ion to water molecules until a new equilibrium is reached.

The Four Types of Salts (and Their pH Behavior)

Every salt falls into one of four categories, based on the strength of the acid and base that originally combined to form it. Salts of a strong acid and a strong base, like sodium chloride (table salt) or potassium nitrate, produce two spectator ions that don't react with water at all — these solutions stay neutral, at pH 7.

Salts of a weak acid and a strong base, like sodium acetate (from acetic acid and sodium hydroxide), leave behind an anion that's the conjugate base of a weak acid. That anion reacts with water — A- + H2O <-> HA + OH- — producing extra hydroxide ions, so these solutions turn basic (pH above 7).

Salts of a strong acid and a weak base, like ammonium chloride (from hydrochloric acid and ammonia), leave behind a cation that's the conjugate acid of a weak base. That cation reacts with water — BH+ + H2O <-> B + H3O+ — producing extra hydrogen ions, so these solutions turn acidic (pH below 7).

Salts of a weak acid and a weak base, like ammonium acetate, are the most interesting case: both ions hydrolyze at the same time, one pulling the pH up and the other pushing it down. Whichever effect is stronger — measured by comparing Ka of the parent acid to Kb of the parent base — decides which way the final pH leans.

The Salt Hydrolysis Formulas Used in This Calculator

For a salt of a weak acid and strong base, this calculator first finds the hydrolysis constant of the anion using Kb = Kw / Ka, where Ka belongs to the parent weak acid. It then solves the same ICE-table quadratic used for any weak base, x^2 + Kb.x - Kb.C = 0, for x = [OH-], rather than relying on the sqrt(Kb.C) shortcut — which keeps the answer accurate even when hydrolysis is larger than the usual 5% rule of thumb allows.

For a salt of a strong acid and weak base, the process mirrors this exactly, but on the acid side: Ka = Kw / Kb, where Kb belongs to the parent weak base, followed by solving x^2 + Ka.x - Ka.C = 0 for x = [H+].

For a salt of a weak acid and a weak base, both ions hydrolyze together, and remarkably, the resulting pH barely depends on the salt's concentration at all — a very useful shortcut applies: pH = 1/2 (pKw + pKa - pKb), where pKa belongs to the parent weak acid and pKb belongs to the parent weak base. This calculator uses that formula directly, since it stays accurate across essentially every concentration you'd encounter in a lab or classroom setting.

Worked Example: Sodium Acetate (a Basic Salt)

Sodium acetate, CH3COONa, comes from acetic acid (Ka = 1.8 x 10^-5) and the strong base sodium hydroxide. At a concentration of 0.1 mol/L, the acetate ion's hydrolysis constant works out to Kb = 1.0 x 10^-14 / 1.8 x 10^-5, or about 5.6 x 10^-10.

Solving the ICE-table quadratic for this Kb and concentration gives [OH-] of roughly 7.5 x 10^-6 mol/L, which corresponds to a pOH of about 5.13 and a pH of about 8.87. That confirms sodium acetate solutions are mildly basic — exactly the kind of result this calculator produces instantly for any weak-acid-derived salt.

Worked Example: Ammonium Chloride (an Acidic Salt)

Ammonium chloride, NH4Cl, comes from the strong acid hydrochloric acid and the weak base ammonia (Kb = 1.8 x 10^-5). At 0.1 mol/L, the ammonium ion's Ka works out to 1.0 x 10^-14 / 1.8 x 10^-5, also about 5.6 x 10^-10.

Solving the equivalent quadratic for [H+] gives roughly 7.5 x 10^-6 mol/L, which corresponds to a pH of about 5.13 — a mildly acidic solution. Interestingly, this mirrors the sodium acetate example almost exactly, just flipped from base to acid, since acetic acid's Ka and ammonia's Kb happen to be numerically very close.

Worked Example: Ammonium Acetate (Weak Acid + Weak Base)

Ammonium acetate, CH3COONH4, is the classic textbook example of a weak-acid-weak-base salt, since both its parent acid (acetic acid, Ka = 1.8 x 10^-5) and its parent base (ammonia, Kb = 1.8 x 10^-5) are weak, and happen to have almost identical strength.

Using pH = 1/2 (pKw + pKa - pKb), with pKa and pKb both equal to about 4.74, the two terms cancel out almost perfectly, giving a pH extremely close to 7 — essentially neutral, even though both ions are actively hydrolyzing. This is a good reminder that 'both ions react with water' does not automatically mean the solution ends up far from neutral; what matters is the balance between the two competing hydrolysis reactions.

Reading Percent Hydrolysis

Percent hydrolysis tells you what fraction of the dissolved salt's ion actually went on to react with water, calculated the same way percent dissociation is calculated for a weak acid or weak base: (concentration of the reacted species / starting concentration) x 100. A low percent hydrolysis, typically well under 5% for most common salts at normal lab concentrations, confirms that only a small fraction of the ion actually hydrolyzes — most of it stays in its original, unreacted form.

This number is a useful sanity check: an unusually high percent hydrolysis for a salt believed to have a very small hydrolysis constant can be a sign of an incorrect Ka or Kb value, or an unusually dilute solution where the standard approximations start to break down.

Common Mistakes When Calculating Salt Hydrolysis pH

The most common mistake is forgetting to convert Ka (or Kb) of the parent weak acid or base into the actual hydrolysis constant of the ion in question, using Kb = Kw / Ka or Ka = Kw / Kb. Plugging the parent's Ka directly into a pH formula meant for the conjugate ion gives a badly wrong answer.

A second common mistake is treating every salt of a weak acid and weak base as automatically 'strongly' acidic or basic. In reality, as the ammonium acetate example shows, these salts can land very close to neutral if the parent acid and base have similar strength — the direction depends entirely on comparing Ka to Kb, not on the fact that both are weak.

A third mistake is applying the strong-strong 'always neutral' rule to a salt that only looks simple, like ammonium nitrate — the nitrate ion is indeed unreactive (from strong nitric acid), but the ammonium ion is the conjugate acid of a weak base and does hydrolyze, so the solution is actually mildly acidic, not neutral.

Real-World Uses of Salt Hydrolysis

Salt hydrolysis isn't just an exam topic — it explains everyday chemistry all around us. Baking soda (sodium bicarbonate) is basic in water because the bicarbonate ion is the conjugate base of the weak acid carbonic acid, which is exactly why it's used to neutralize acidic odors and stomach acid. Many buffer systems used in biology, medicine, and food preservation rely on carefully chosen salts of weak acids or weak bases to hold a solution's pH steady within a target range.

In agriculture and soil science, the hydrolysis behavior of fertilizer salts like ammonium nitrate or ammonium sulfate directly affects soil pH over time, since the ammonium ion slowly acidifies the soil as it hydrolyzes. In water treatment, understanding which dissolved salts will shift a water supply's pH toward acidic or basic is essential for keeping drinking water and industrial process water within a safe, stable range.

Salt Hydrolysis pH Calculator: Quick Reference Summary

Salts of a strong acid + strong base stay neutral (pH 7), since neither ion reacts with water. Salts of a weak acid + strong base turn basic: find Kb = Kw / Ka, then solve x^2 + Kb.x - Kb.C = 0 for [OH-]. Salts of a strong acid + weak base turn acidic: find Ka = Kw / Kb, then solve x^2 + Ka.x - Ka.C = 0 for [H+]. Salts of a weak acid + weak base use pH = 1/2 (pKw + pKa - pKb), a formula that barely depends on concentration at all.

This free calculator is meant to support learning, homework checking, and everyday chemistry questions. Ka and Kb values can vary slightly between textbooks and reference sources and change with temperature, so for lab reports, research, or any safety-critical application, always confirm the specific values with your course materials or a peer-reviewed reference source.