What Salt Hydrolysis Means
When a salt dissolves in water it ionises into its constituent cation and anion. The fate of those ions is not always passive. The cations of strong bases, such as $\ce{Na+}$, $\ce{K+}$, $\ce{Ca^2+}$ and $\ce{Ba^2+}$, and the anions of strong acids, such as $\ce{Cl-}$, $\ce{Br-}$, $\ce{NO3-}$ and $\ce{ClO4-}$, merely become hydrated and do not react further. Other ions, however, interact with water and reform the corresponding weak acid or weak base. This process of interaction between water and the cations or anions, or both, of a salt is called hydrolysis.
Hydrolysis disturbs the self-ionisation balance of water. If the reaction releases $\ce{OH-}$ ions the solution turns basic; if it releases $\ce{H+}$ ions the solution turns acidic. The key insight, taken straight from NCERT, is that a conjugate of a strong parent has no measurable tendency to hydrolyse, while the conjugate of a weak parent does. Hydrolysis is therefore the mirror image of the ionisation you studied for weak acids and bases: the conjugate base of a weak acid is itself a moderately effective base, and the conjugate acid of a weak base is a moderately effective acid.
It helps to see hydrolysis as a Bronsted proton transfer. The acetate anion is a proton acceptor, so when it meets water it pulls a proton off and leaves $\ce{OH-}$ behind. The ammonium cation is a proton donor, so it hands a proton to water and leaves $\ce{H+}$ behind. Ions that are conjugates of strong electrolytes have effectively zero proton-transfer affinity, which is exactly why they refuse to hydrolyse. The strength of the parent and the extent of hydrolysis are thus inversely linked.
Four combinations of parent acid and base are possible, and each gives a characteristic behaviour. We examine them in turn, then unify them with the hydrolysis constant $K_h$ and the degree of hydrolysis. The four NIOS-listed archetypes are $\ce{NaCl}$ (strong+strong), $\ce{CH3COONa}$ (weak acid+strong base), $\ce{NH4Cl}$ (strong acid+weak base) and $\ce{CH3COONH4}$ (weak acid+weak base), and every salt you meet in NEET reduces to one of these patterns.
Each salt category lands on a different region of the pH axis. The conjugate of the weaker parent is the ion that hydrolyses and pushes the pH away from 7.
Strong Acid + Strong Base: Neutral
Sodium chloride is the textbook example, formed from the strong acid $\ce{HCl}$ and the strong base $\ce{NaOH}$. In water it dissociates completely:
$$\ce{NaCl(aq) -> Na+(aq) + Cl-(aq)}$$
Neither $\ce{Na+}$ nor $\ce{Cl-}$ reacts with water. $\ce{Na+}$ is the conjugate of a strong base and $\ce{Cl-}$ is the conjugate of a strong acid; both are too weak to abstract or donate a proton. The ions simply hydrate, the water self-ionisation equilibrium stays undisturbed, and the solution is neutral with $\text{pH} = 7$ at 298 K. This is the only salt category in which no hydrolysis occurs.
The same conclusion holds for $\ce{KNO3}$, $\ce{NaBr}$, $\ce{KClO4}$ and $\ce{CaCl2}$, all of which combine a strong-base cation with a strong-acid anion. Because the hydration energy of these ions is fully satisfied by the surrounding water dipoles, there is no thermodynamic driving force to regenerate the parent acid or base. The neutrality is a direct consequence of both parents being completely ionised electrolytes whose conjugates are non-reactive.
Weak Acid + Strong Base: Basic
Sodium acetate, $\ce{CH3COONa}$, is the salt of the weak acid acetic acid and the strong base $\ce{NaOH}$. It ionises completely in water:
$$\ce{CH3COONa(aq) -> CH3COO-(aq) + Na+(aq)}$$
The sodium ion does not hydrolyse, but the acetate ion is the conjugate base of a weak acid and is therefore appreciably basic. It abstracts a proton from water:
$$\ce{CH3COO-(aq) + H2O(l) <=> CH3COOH(aq) + OH-(aq)}$$
Because acetic acid is weak ($K_a = 1.8 \times 10^{-5}$) it remains largely unionised, so the equilibrium lies far enough to the right to raise the $\ce{OH-}$ concentration. The solution becomes alkaline and its pH exceeds 7. This anionic hydrolysis is the signature of every weak-acid, strong-base salt: sodium carbonate, sodium cyanide and potassium nitrite all behave the same way.
The weaker the parent acid, the more strongly its conjugate base hydrolyses. Sodium cyanide, derived from the very weak $\ce{HCN}$, gives a markedly more basic solution than sodium acetate at the same concentration, because the cyanide ion is a far better proton acceptor than acetate. This trend is captured quantitatively by $K_h = K_w/K_a$, which grows as $K_a$ shrinks. Carbonate salts behave even more strongly because $\ce{CO3^2-}$ hydrolyses in two steps, which is why washing soda solutions feel slippery and are distinctly alkaline.
Acetate, the conjugate of a weak acid, pulls a proton off water. Acetic acid stays largely unionised, so the freed $\ce{OH-}$ accumulates and the solution turns basic.
Strong Acid + Weak Base: Acidic
Ammonium chloride, $\ce{NH4Cl}$, is the salt of the strong acid $\ce{HCl}$ and the weak base $\ce{NH4OH}$. It dissociates fully:
$$\ce{NH4Cl(aq) -> NH4+(aq) + Cl-(aq)}$$
The chloride ion does not hydrolyse, but the ammonium ion is the conjugate acid of a weak base and donates a proton to water:
$$\ce{NH4+(aq) + H2O(l) <=> NH4OH(aq) + H+(aq)}$$
Ammonium hydroxide is a weak base ($K_b = 1.77 \times 10^{-5}$) and stays almost unionised, so the released $\ce{H+}$ ions accumulate. The $\ce{H+}$ concentration rises, and the pH of an $\ce{NH4Cl}$ solution is less than 7. This cationic hydrolysis governs all strong-acid, weak-base salts, including ammonium nitrate and aniline hydrochloride.
A useful symmetry now appears between the second and third categories. In the basic case the anion is the active species and the cation is a spectator; in the acidic case the cation is active and the anion is a spectator. In both, exactly one ion of the salt hydrolyses, and that ion is always the conjugate of the weaker parent. Hydrated metal cations of small, highly charged ions such as $\ce{Al^3+}$ and $\ce{Fe^3+}$ also behave as weak-base conjugates and make their salt solutions acidic, which is why $\ce{AlCl3}$ and $\ce{FeCl3}$ solutions test acidic to litmus despite chloride being inert.
Salt hydrolysis rests entirely on the $K_a$, $K_b$ and $K_w$ relations. Revise them in pH, Ionisation & Buffer Solutions before tackling the formulae below.
Weak Acid + Weak Base
Ammonium acetate, $\ce{CH3COONH4}$, is formed from a weak acid and a weak base, so both ions hydrolyse simultaneously:
$$\ce{CH3COO- + NH4+ + H2O <=> CH3COOH + NH4OH}$$
The two weak products themselves remain only partially dissociated, alongside the water equilibrium:
$$\ce{CH3COOH <=> CH3COO- + H+} \qquad \ce{NH4OH <=> NH4+ + OH-} \qquad \ce{H2O <=> H+ + OH-}$$
The nature of the solution now depends on a tug of war between the two hydrolyses. NCERT states that, without detailed calculation, the degree of hydrolysis here is independent of the concentration of the solution, and the pH is fixed by the $pK$ values alone.
$$\text{pH} = 7 + \tfrac{1}{2}\,(pK_a - pK_b)$$
The solution is basic if $pK_a > pK_b$ (the acid is the weaker partner), acidic if $pK_a < pK_b$, and almost exactly neutral when $pK_a = pK_b$, as in ammonium acetate where the two are nearly equal.
The Hydrolysis Constant Kh
Each hydrolysis equilibrium has its own equilibrium constant, the hydrolysis constant $K_h$. For the acetate-ion case, writing the constant and then multiplying numerator and denominator by $[\ce{H+}]$ reveals how it connects to $K_a$ and $K_w$:
$$K_h = \frac{[\ce{CH3COOH}][\ce{OH-}]}{[\ce{CH3COO-}]} = \frac{K_w}{K_a}$$
The same algebra applied to the three salt categories that hydrolyse gives a compact set of relations. The weaker the parent, the smaller its $K_a$ or $K_b$, and therefore the larger $K_h$ and the greater the extent of hydrolysis.
| Salt type | Ion that hydrolyses | Hydrolysis constant |
|---|---|---|
| Weak acid + strong base | anion | $K_h = K_w / K_a$ |
| Strong acid + weak base | cation | $K_h = K_w / K_b$ |
| Weak acid + weak base | both ions | $K_h = K_w / (K_a K_b)$ |
Degree of Hydrolysis & pH Formulae
The degree of hydrolysis, $h$, is the fraction of the salt ion that has reacted with water at equilibrium. For a salt of a weak acid and strong base at concentration $c$, substituting equilibrium amounts into the $K_h$ expression and assuming $h \ll 1$ gives a square-root law, from which the hydroxide concentration and the pH follow.
$$h = \sqrt{\frac{K_h}{c}} = \sqrt{\frac{K_w}{K_a\,c}} \qquad \text{pH} = 7 + \tfrac{1}{2}\,(pK_a + \log c)$$
For a salt of a strong acid and weak base the parallel expressions involve $K_b$, and the solution is acidic:
$$h = \sqrt{\frac{K_w}{K_b\,c}} \qquad \text{pH} = 7 - \tfrac{1}{2}\,(pK_b + \log c)$$
Note that $c$ appears under the root, so for these two categories dilution increases the degree of hydrolysis. For the weak-acid, weak-base salt, by contrast, $h = \sqrt{K_h} = \sqrt{K_w/(K_a K_b)}$ and the pH formula carry no concentration term, which is why ammonium acetate gives essentially the same pH at any dilution.
The dilution result deserves a moment's thought because it seems counter-intuitive. Adding water lowers $c$, which by Le Chatelier's principle shifts the hydrolysis equilibrium toward more products, so a larger fraction of the salt hydrolyses. Yet the absolute concentration of $\ce{OH-}$ or $\ce{H+}$ produced still falls, so a dilute sodium acetate solution is closer to neutral than a concentrated one, even though its degree of hydrolysis is higher. The two statements are consistent: a bigger fraction of a smaller amount can still be a smaller quantity.
For the weak-acid, weak-base salt the cancellation of the concentration term is the reason its pH is decided purely by the relative weakness of the two parents. When $pK_a$ exceeds $pK_b$ the conjugate base hydrolyses more than the conjugate acid and the medium is basic; when $pK_b$ exceeds $pK_a$ the reverse holds and the medium is acidic; and when the two are equal the hydrolyses are balanced and the solution is neutral. This is the analytical content behind the one-line formula $\text{pH} = 7 + \tfrac{1}{2}(pK_a - pK_b)$.
The $pK_a$ of acetic acid and the $pK_b$ of ammonium hydroxide are 4.76 and 4.75 respectively. Calculate the pH of an ammonium acetate solution.
Ammonium acetate is the salt of a weak acid and a weak base, so the concentration-free formula applies:
$$\text{pH} = 7 + \tfrac{1}{2}(pK_a - pK_b) = 7 + \tfrac{1}{2}(4.76 - 4.75) = 7 + 0.005 = 7.005$$
The two parents are almost equally weak, so the solution is essentially neutral, just a hair above 7. This is Problem 6.25 of NCERT Equilibrium.
Estimate the pH of a 0.1 M sodium acetate solution, given $K_a$ of acetic acid $= 1.8 \times 10^{-5}$ ($pK_a = 4.74$) and $K_w = 1.0 \times 10^{-14}$.
Acetate is the conjugate base of a weak acid, so the anion hydrolyses. Using the basic-salt pH formula with $c = 0.1$ M, so $\log c = -1$:
$$\text{pH} = 7 + \tfrac{1}{2}(pK_a + \log c) = 7 + \tfrac{1}{2}(4.74 - 1) = 7 + 1.87 = 8.87$$
The solution is basic, consistent with the qualitative result that a weak-acid, strong-base salt gives pH greater than 7. The hydrolysis constant here is $K_h = K_w/K_a = 10^{-14}/(1.8\times10^{-5}) = 5.6 \times 10^{-10}$.
Master Table of Salt Types
The four categories can be collapsed into a single reference. Read it by first identifying the parent acid and base, then locating the ion that hydrolyses.
| Salt type | Example | Hydrolysing ion | Nature | Approx. pH (0.1 M) |
|---|---|---|---|---|
| Strong acid + strong base (SA+SB) | $\ce{NaCl}$, $\ce{KNO3}$ | none | neutral | = 7 |
| Weak acid + strong base (WA+SB) | $\ce{CH3COONa}$, $\ce{Na2CO3}$ | anion | basic | > 7 (about 8.9) |
| Strong acid + weak base (SA+WB) | $\ce{NH4Cl}$, $\ce{NH4NO3}$ | cation | acidic | < 7 (about 5.1) |
| Weak acid + weak base (WA+WB) | $\ce{CH3COONH4}$ | both | depends on $pK_a$ vs $pK_b$ | near 7 (set by $pK$) |
Predicting nature from parent strength
Students often guess the nature of a salt solution from the salt's own appearance instead of tracing its parents. The rule is mechanical: split the salt into the acid and base that would form it, then ask which one is weak. The conjugate of the weaker parent is the ion that hydrolyses and dictates the pH. A strong-strong pair leaves the water balance untouched, so the solution is neutral.
For a WA+WB salt, do not call it neutral by default. Compare $pK_a$ and $pK_b$: basic if $pK_a > pK_b$, acidic if $pK_a < pK_b$, neutral only when they are equal.
Predicting Acidic, Basic or Neutral
The fastest exam strategy is to memorise the parent strengths of the common acids and bases, then apply the master table. $\ce{HCl}$, $\ce{HNO3}$, $\ce{H2SO4}$, $\ce{HClO4}$ and $\ce{HBr}$ are the standard strong acids; $\ce{NaOH}$, $\ce{KOH}$, $\ce{Ca(OH)2}$ and $\ce{Ba(OH)2}$ are the standard strong bases. Anything outside these lists, such as acetic acid, carbonic acid, $\ce{HCN}$ or ammonia, is weak and its conjugate will hydrolyse.
Once the hydrolysing ion is known, the direction of the pH shift follows immediately, and the formulae of the previous section convert that into a number. The recurring NEET demand is qualitative, so the prediction step alone is usually worth the mark.
A second layer of questions tests the litmus or indicator response of a salt solution, which is just the nature dressed differently: a basic salt turns red litmus blue, an acidic salt turns blue litmus red, and a neutral salt leaves both unchanged. Acid-base titration end-points draw on the same idea, because the salt formed at equivalence sets the pH there. A strong-acid, strong-base titration ends near pH 7; a weak-acid, strong-base titration ends above 7 because the product salt hydrolyses to a basic solution; and a strong-acid, weak-base titration ends below 7 for the mirror reason.
Hydrolysis of salts in one screen
- Hydrolysis is the reaction of a salt's cation or anion with water; conjugates of strong parents do not hydrolyse.
- SA+SB salts (NaCl) are neutral, pH 7; no ion hydrolyses.
- WA+SB salts (CH3COONa) are basic, pH > 7; the anion hydrolyses giving $\ce{OH-}$.
- SA+WB salts (NH4Cl) are acidic, pH < 7; the cation hydrolyses giving $\ce{H+}$.
- WA+WB salts (CH3COONH4): pH $= 7 + \tfrac{1}{2}(pK_a - pK_b)$, independent of concentration.
- Hydrolysis constants: $K_h = K_w/K_a$, $K_w/K_b$ or $K_w/(K_a K_b)$; degree of hydrolysis $h = \sqrt{K_h/c}$ for the first two, so dilution increases hydrolysis.