Ionic Equilibrium: One Mark Questions, Book Back Answers and Notes

About Ionic Equilibrium

Medium ~120 min study

When acetic acid dissolves in water, only a small part of it splits into H⁺ and CH₃COO⁻ ions. The rest stays as molecules, and the two forms sit in equilibrium. This chapter applies the equilibrium ideas from Class 11 to such solutions of acids, bases and salts, where the concentrations of ions decide how the solution behaves.

It begins with three definitions of acids and bases, each wider than the last. Then it turns to water itself, which ionises very slightly, and uses that to define pH. Weak acids and bases need a dissociation constant, K_a or K_b, because they do not ionise completely, and Ostwald's dilution law connects that constant to the degree of dissociation.

This unit is heavy on calculation. You will be asked for the pH of a weak acid, the pH of a buffer from the Henderson equation, the pH of a hydrolysed salt, and whether a precipitate forms when two solutions are mixed. Get comfortable with logarithms and with writing K_sp for salts like Ag₂CrO₄, and most of these become routine.

What you'll learn

• Explain the three acid-base theories and identify conjugate acid-base pairs.
• Calculate pH and pOH of strong and weak acids and bases.
• Derive Ostwald's dilution law and use it for weak electrolytes.
• Explain buffer action in acidic and basic buffers and calculate buffer pH.
• Find the hydrolysis constant and pH of different types of salt solutions.
• Predict precipitation by comparing ionic product with the solubility product.

Before you start

• Chemical equilibrium, equilibrium constant and Le Chatelier's principle.
• Molarity and other concentration terms.
• Logarithms and scientific notation.

Ionic Equilibrium explained

Acids, Bases, Buffers and Solubility Product

Theories of Acids and Bases

Arrhenius defined an acid as a substance that gives H⁺ ions in water and a base as one that gives OH⁻ ions. This works only in aqueous solution. Lowry and Bronsted called an acid a proton donor and a base a proton acceptor. When HCl gives a proton to water, Cl⁻ is the conjugate base of HCl and H₃O⁺ is the conjugate acid of water. Lewis went further: an acid accepts an electron pair and a base donates one. BF₃ and AlCl₃ are Lewis acids even though they have no proton to give.

Ionic Product of Water and pH

Water ionises slightly: H₂O ⇌ H⁺ + OH⁻. The product of the two ion concentrations is the ionic product of water, \(K_{w}\), equal to 1.0 × 10^-14 at 298 K. Since these concentrations are so small, we use pH = −log[H₃O⁺]. At 298 K, pH + pOH = 14 and a neutral solution has pH 7. Because \(K_{w}\) increases with temperature, the pH of pure water at a higher temperature is below 7 even though it is still neutral.

Weak Electrolytes and Ostwald's Dilution Law

A weak acid HA with concentration C and degree of dissociation α has K_a = α²C/(1 − α). Since α is small, K_a ≈ α²C, so α = √(K_a/C). This is Ostwald's dilution law: as C decreases on dilution, α increases. It applies to weak acids like acetic acid and weak bases like ammonia, not to strong electrolytes.

Common Ion Effect and Buffers

Adding sodium acetate to acetic acid increases the acetate ion concentration and pushes the dissociation of acetic acid backwards. This suppression is the common ion effect. It is the reason a buffer works. A buffer resists change in pH when a small amount of acid or base is added. An acidic buffer is a weak acid with its salt (CH₃COOH + CH₃COONa); a basic buffer is a weak base with its salt (NH₄OH + NH₄Cl). Its pH is found from the Henderson–Hasselbalch equation, pH = pK_a + log([salt]/[acid]). Blood is kept near pH 7.4 by buffer systems.

Salt Hydrolysis and Solubility Product

In salt hydrolysis the ions of a salt react with water. A salt of a strong acid and weak base (NH₄Cl) gives an acidic solution; a salt of a weak acid and strong base (CH₃COONa) gives a basic one; a salt of a strong acid and strong base (NaCl) does not hydrolyse. For a sparingly soluble salt such as AgCl, the solubility product \(K_{sp}\) = [Ag⁺][Cl⁻] at saturation. If the ionic product in a solution exceeds \(K_{sp}\), a precipitate forms. This rule is used in salt analysis to separate groups of cations.

Common mistakes to avoid

• Mixing up the members of a conjugate pair. The conjugate acid always has one proton more than its conjugate base, as in H2O and OH−.
• Applying Ostwald's dilution law to strong electrolytes like HCl, which are fully ionised at every concentration.
• Assuming pH 7 is neutral at every temperature. That holds only at 298 K, where \(K_{w}\) = 1.0 × 10-14.
• Assuming every salt solution is neutral. NH4Cl is acidic and CH3COONa is basic because of hydrolysis.
• Taking Ksp as the solubility itself. For AgCl, Ksp = s2, but for a salt like Ag2CrO4, Ksp = 4s3.

Test yourself on these with the practice test, then check the worked reasoning in the solved MCQs.

Frequently asked questions

How does a buffer solution keep its pH nearly constant?

A buffer holds both an acid part and a base part. In the acetic acid/sodium acetate buffer, added H+ is taken up by acetate ions to form acetic acid, and added OH− is neutralised by acetic acid. The added ions are removed before they can change pH much. An ammonia/ammonium chloride buffer works the same way on the basic side.

What is the difference between a Lewis acid and a Bronsted acid?

A Bronsted acid is a proton donor, such as HCl. A Lewis acid is an electron-pair acceptor, such as BF3 or AlCl3. The two sets overlap but are not the same. HCl itself does not accept an electron pair; the proton it releases is the Lewis acid. On the other side, every Bronsted base is also a Lewis base, because accepting a proton means donating an electron pair to it.

Why does the degree of dissociation increase on dilution?

Dilution lowers the concentration of all species. By Le Chatelier's principle the equilibrium HA ⇌ H+ + A− shifts to the side with more particles, which is the ionised side. Mathematically, α = √(Ka/C), so a smaller C gives a larger α.

Can a neutral solution have a pH other than 7?

Yes. Neutral means [H3O+] = [OH−], not pH = 7. Ionisation of water is endothermic, so Kw increases with temperature. Above 298 K both ion concentrations rise, and the pH of pure water drops below 7 while the solution stays neutral.

When does the common ion effect occur?

It occurs when a strong electrolyte that shares an ion with a weak electrolyte is added to its solution. Adding NH4Cl to NH4OH raises [NH4+], the equilibrium shifts back, and the dissociation of NH4OH falls. This is why [OH−] is kept low in group III of salt analysis.

Which salts undergo hydrolysis?

Salts in which at least one ion comes from a weak acid or a weak base. That ion reacts with water to re-form the weak parent, leaving extra H3O+ or OH−. Salts of a strong acid and a strong base, like NaCl, do not hydrolyse and give a neutral solution.