2. Ionic equilibrium

Ionic equilibrium is the study of equilibrium in electrolyte solutions — the balance between ions and undissociated molecules.

Ionic Reactions

When electrolytes dissolve in water, they dissociate into charged particles called ions. Reactions that occur between ions in solution are called ionic reactions. Ionic reactions form the basis of all acid-base chemistry, precipitation, and complexation phenomena.

When silver nitrate (AgNO₃) and sodium chloride (NaCl) solutions are mixed, a white precipitate of silver chloride (AgCl) forms. We can write this in three ways:

Acid-Base Concepts — Overview

The three important acid-base concepts developed in this order:

1. Arrhenius concept  →  2. Brønsted–Lowry concept  →  3. Lewis concept

1. Arrhenius Concept

An acid is a substance that gives hydrogen ions (H⁺) when dissolved in water.

A base is a substance that gives hydroxyl ions (OH⁻) when dissolved in water.

A salt is an ionic compound formed from the positive ion of a base and the negative ion of an acid.

Neutralisation according to Arrhenius: An acid and a base react to form water and a salt:

Limitations of Arrhenius’s Concept

The main limitation is that Arrhenius cannot adequately describe substances that can behave as acids and bases outside the narrow H⁺/OH⁻ aqueous framework (cannot explain acidic/basic behaviour in non-aqueous media). For example:

• Cannot explain the basicity of NH₃ and its organic derivatives, as they do not contain OH⁻.
• Metal oxides (CaO, Na₂O) act as bases but do not contain OH⁻.
• Certain compounds (AlCl₃, BF₃) show acidic nature without having H⁺ ions.
• Fails to account for the amphoteric nature of substances (e.g., Al₂O₃, ZnO).

2. Brønsted–Lowry Concept (Proton Transfer Theory)

• Acid: A species that donates one or more protons (H⁺) — proton donor.
• Base: A species that accepts one or more protons (H⁺) — proton acceptor.

Species that tend to give as well as accept a proton are termed amphoteric species or substances. Example: H₂O, HCO₃⁻, and HSO₄⁻ are amphoteric or amphiprotic substances.

Amphoteric is a general term for a substance showing the dual nature of acid and base. While the amphiprotic term is used when acid-base behaviour is restricted to proton transfer.

Conjugate acid-base pairs

When an acid donates a proton, the remaining species is its conjugate base. When a base accepts a proton, the new species is its conjugate acid.

From the above reaction:

• HCl (acid) and Cl⁻ (conjugate base) form one pair.
• H₂O (base) and H₃O⁺ (conjugate acid) form the other pair.

The Relative Strength of Acids and Bases

If an acid is strong, its conjugate base will be weak, and vice versa.

↑ Acid strength increases upward   |   Base strength increases downward ↓

• Strong acid → weak conjugate base (stable, not eager to accept a proton back).
• Weak acid → strong conjugate base (unstable, eager to accept a proton).
• Strong base → weak conjugate acid.
• Weak base → strong conjugate acid.

e.g., HCl (a strong acid) has Cl⁻ as the conjugate base (a very weak base that does not accept H⁺ easily).
CH₃COOH (weak acid) has CH₃COO⁻ as its conjugate base (a moderately strong base).

Advantages over Arrhenius

1. Works in any solvent (water, alcohol, liquid NH₃) and even in the gas phase.
2. Explains NH₃ as a base (it accepts a proton from water: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻).
3. Explains H₂O as amphiprotic (can act as an acid or base).
4. Explains metal oxides as bases (O²⁻ + H₂O → 2 OH⁻, then OH⁻ accepts protons).

Limitation of Brønsted–Lowry

It still requires protons. It cannot explain acid-base reactions that do not involve proton transfer (e.g., BF₃ + NH₃ → F₃B–NH₃, or metal ions acting as acids).

3. Lewis Concept (Electronic Concept)

Acid: A species that accepts a pair of electrons from a base to form a coordinate covalent bond, i.e., an electron pair acceptor. An electrophile (electron-loving species) acts as a Lewis acid.

Base: A species (atoms, ions, molecules, etc.) that donates a pair of electrons to an acid to form a coordinate covalent bond. Nucleophiles (nucleus-loving species) act as Lewis bases.

Merits of the Lewis concept

1. Includes non-proton reactions: Covers coordination chemistry, metal complexes, and many organic reactions (e.g., Friedel-Crafts alkylation using AlCl₃ as a Lewis acid).
2. It also includes the basic properties of metallic oxides and the acidic properties of nonmetallic oxides.
3. It can also explain that H⁺ and OH⁻ are Lewis acids and Lewis bases, respectively.
4. Explains acidity of salts like AlCl₃, BF₃, FeCl₃: They are electron-deficient and accept electron pairs.

Demerits of Lewis Concept

1. Fails to arrange the acids and bases according to relative strength.
2. Says a coordinate covalent bond is formed between an acid and a base. However, strong mineral acids such as H₂SO₄, HCl, and HNO₃ do not form any coordinate covalent bonds.
3. Most acid-base reactions (e.g., HCl + NaOH in water) are very fast, but forming a coordinate covalent compound is slow. Such a rapid acid-base reaction is not accounted for.

Comparison Table

Summary of Progression

1. Arrhenius started the idea: acids give H⁺, bases give OH⁻ in water. Problem: Too narrow.
2. Brønsted–Lowry expanded: acids donate H⁺, bases accept H⁺ in any medium. Problem: Still requires protons.
3. Lewis gave the most general definition: acids accept electron pairs, bases donate electron pairs. Problem: Loses quantitative strength ordering.

• Electrolytes are substances like acids, bases, and salts that can conduct electricity in their aqueous solution due to ionisation.
• Some electrolytes, like potassium chloride and sodium hydroxide, completely dissociate into their constituent ions in solution. These are known as strong electrolytes.

• Weak electrolytes ionise only partially in water.
• In the aqueous solution of weak electrolytes, the constituent ions are in equilibrium with undissociated molecules of the electrolyte.

• An equilibrium exists between ions and unionised molecules, called ionic equilibrium.
• Weak acids and weak bases are good examples of weak electrolytes.

Degree of ionisation (α): The fraction of the initial number of molecules (or formula units) that ionise into ions is called the degree of ionisation or dissociation. The degree of ionisation increases with dilution.

Ostwald’s dilution law describes how the degree of ionisation (α) of a weak electrolyte changes with dilution (i.e., when you add more solvent, usually water).

Consider a weak binary electrolyte AB with an initial concentration of the electrolyte, C (in mol/L), and degree of ionisation α (fraction of AB that dissociates) that dissociates as:

The equilibrium constant or dissociation constant K is given by the law of mass action as:

For a weak electrolyte, α is very small (α << 1). Therefore, 1−α ≈ 1. The equation simplifies to:

Solving for α:

Therefore, when concentration C decreases, α increases.

If one mole of solute is dissolved in V litres of solution, the concentration is given as:

Therefore:

• For weak electrolytes, the ions in solution are far apart at equilibrium, so the approximation holds.
• For strong electrolytes, they are almost fully dissociated even at high concentration. Dilution does not increase α (it’s already ~1), and the simple equilibrium constant expression does not correctly describe their behaviour due to interionic attractions.

Ionisation of Water

Water is a weak electrolyte and undergoes very slight ionisation:

• This is a reversible (equilibrium) reaction.
• Only a very small fraction of water molecules ionise at any time.
• Most of the water remains as unionised H₂O molecules.

👉 More accurately: H⁺ exists as H₃O⁺ (hydronium ion), but we write H⁺ for simplicity.

Ionic Product of Water (K_w)

The product of hydrogen ion and hydroxide ion concentrations is constant:

At 25°C (room temperature):

• [H⁺] = 1×10⁻⁷ mol L⁻¹
• [OH⁻] = 1×10⁻⁷ mol L⁻¹

Important Implications

• If [H⁺] increases → [OH⁻] must decrease
• If [OH⁻] increases → [H⁺] must decrease

👉 Because: [H⁺] × [OH⁻] = 10⁻¹⁴

Effect in Acidic Solutions

• Acids increase [H⁺]
• So [OH⁻] decreases automatically

e.g., If [H⁺] = 10⁻³, then [OH⁻] = 10⁻¹¹ ✔ Solution is acidic

Effect in Basic Solutions

• Bases increase [OH⁻]
• So [H⁺] decreases

e.g., If [OH⁻] = 10⁻², then [H⁺] = 10⁻¹² ✔ Solution is basic

Neutral Solution Condition

• [H⁺] = [OH⁻]
• Both are 10⁻⁷ ✔ Solution is neutral at 25°C (pH = 7)
• K_w generally increases with temperature.

What is pH?

pH is a measure of the acidity or alkalinity of a solution. It depends on the concentration of hydrogen ions (H⁺) present in the solution.

Higher [H⁺] → lower pH → more acidic.   Lower [H⁺] → higher pH → more basic.

pH Scale

The pH scale generally ranges from 0 to 14:

• pH < 7 → Acidic solution
• pH = 7 → Neutral solution (pure water) at 25°C
• pH > 7 → Basic (alkaline) solution

Real-Life Importance of pH

• Biological systems: Blood pH ≈ 7.4 (slight change is dangerous)
• Agriculture: Soil pH affects crop yield
• Industry: Many reactions require controlled pH
• Daily life: Tooth decay occurs when the mouth pH < 5.5

Weak acids like ethanoic acid dissociate partially as:

An equilibrium is established between the ions and undissociated molecules. The equilibrium constant characterises the equilibrium as:

For dilute solutions, the concentration of water is relatively large and constant. Therefore:

A higher value of K_a indicates that the acid is ionised to a larger extent; hence it’s a strong acid, and vice versa.

Acid strength is often expressed as pK_a, where pK_a = −log K_a.

The larger the K_a, the smaller the pK_a. Hence, a smaller pK_a value indicates a relatively stronger acid.

When a sparingly soluble salt like AgCl is dissolved in water, it dissolves only slightly and forms a saturated solution at a given temperature. At this stage, a dynamic equilibrium is established between the dissolved ions and the undissolved solid.

This equilibrium can be represented as:

Applying the law of mass action to the above reversible equation at the equilibrium condition:

The concentration of pure solid AgCl is treated as constant, so it is not included in the K_sp expression. Hence, the above equation can be written as:

Where K_sp is known as the solubility product constant or solubility product; its value for a sparingly soluble salt is constant at a given temperature.

Application of Solubility Product Principle

Precipitation of chlorides of Group I basic radicals (Ag⁺, Pb²⁺, Hg₂²⁺)

When HCl is added to the solution containing various metal ions, the solubility product of only AgCl, PbCl₂, and Hg₂Cl₂ is less than the ionic product. On the other hand, the solubility product of the chlorides of other cations is higher than the ionic product. Hence, they remain in the solution without precipitation.

The Common Ion Effect is the suppression of the degree of ionisation of a weak electrolyte by adding a strong electrolyte with an ion common to the weak electrolyte.

Due to the addition of NH₄Cl, the concentration of NH₄⁺ ions in solution increases. The increased concentration of NH₄⁺ in the solution shifts the above equilibrium reaction towards the left. Thus, some NH₄⁺ and OH⁻ ions recombine to form unionised NH₄OH — i.e., the degree of ionisation of the weak electrolyte is suppressed, thereby decreasing the concentration of OH⁻ in the solution.

Application of the Common Ion Effect

A. In the selective precipitation of basic radicals in qualitative analysis:

1. Precipitation of sulphides of Group II basic radicals (Hg²⁺, Pb²⁺, Bi³⁺, Cu²⁺)

The group reagent for the Group II basic radical is H₂S in the presence of HCl. When H₂S gas is passed through the solution containing HCl under hot conditions, the dissociation takes place as:

Addition of HCl increases [H⁺], which suppresses the ionisation of H₂S and therefore decreases [S²⁻].

The concentration of metal ions is already present. For sufficiently insoluble Group II sulphides, the ionic product [M²⁺][S²⁻] can still exceed K_sp, so they precipitate. Therefore, only the sulphides of the Group II metal ions get precipitated out.

B. Purification of Common Salt:

The impure salt (sea salt) can be made pure by passing HCl gas through the saturated sea salt solution.

Here, the increase in the concentration of Cl⁻ ions results in an increase in the ionic product of Na⁺ and Cl⁻ over the solubility product of NaCl. Hence, pure NaCl gets precipitated while all impurities remain in the solution.

A buffer solution is a solution that resists significant change in pH when small amounts of a strong acid or a strong base are added to it. Buffer solutions are vital in biological systems, industrial processes, and pharmaceutical preparations.

A. Basic Buffer

Composed of a weak base and its conjugate salt (formed with a strong acid).

Example: NH₄OH (ammonium hydroxide) + NH₄Cl (ammonium chloride)

B. Acidic Buffer

Composed of a weak acid and its conjugate salt (formed with a strong base).

Example: CH₃COOH (acetic acid) + CH₃COONa (sodium acetate)

The weak acid provides an H⁺ reserve to neutralise added base; the conjugate base from the salt neutralises added acid.

When acid is added:   CH₃COO⁻ + H⁺ → CH₃COOH
When base is added:   CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O

Applications of Buffer Solutions

• Blood buffer: Blood is maintained at pH 7.35–7.45 by the H₂CO₃/HCO₃⁻ buffer system. Blood pH must remain within a narrow range for normal body function.
• Industry: Electroplating, fermentation, and dyeing industries require precise pH control using buffers.
• Agriculture: Soil buffers maintain optimal pH for plant growth.
• Pharmaceuticals: Medicines (especially injections) are buffered to prevent irritation or degradation.
• Laboratory: Buffer solutions are used to calibrate pH meters.

Indicators are generally weak organic acids or bases that indicate the endpoint of a reaction by changing their colours. The indicators have different colours in acidic and alkaline solutions. Some common indicators for acid-base titrations include methyl orange, phenolphthalein, and methyl red.

The choice of a suitable indicator for a particular acid-base titration depends upon the nature of the acid and base used.

During acid-base titration, neutralisation of an acid by a base and vice versa occurs. During neutralisation, H⁺ ions and OH⁻ ions combine, forming water molecules. Hence, the pH of the resulting solution changes.

A plot of the pH of the solution against the volume of acid or base added from the burette is called a titration curve or pH curve. Such a titration curve graphically shows the equivalence point and helps select the proper indicator.

Acid-Base Titrations are of the Following Types

1. Strong acid–strong base (HCl vs NaOH) titration:
When a strong acid is titrated against a strong base, a pH curve with a sharp pH change from approx. 3–11 is obtained. Therefore, the indicators methyl orange (pH range 3.1–4.4) and phenolphthalein (pH range 8.2–10) can both be used as suitable indicators.

2. Strong acid–weak base (HCl vs NH₄OH) titration:
When a strong acid is titrated against a weak base, a pH curve with a sharp change in pH from approx. 3–8 is obtained. Therefore, the indicator methyl orange, having a pH range of 3.1–4.4, can be used as a suitable indicator.

3. Weak acid–strong base (CH₃COOH vs NaOH) titration:
When a weak acid is titrated against a strong base, a pH curve with a sharp pH change from 6–11 (approx.) is obtained. Therefore, the indicator phenolphthalein, having a pH range of 8.2–10, can be used as a suitable indicator.

4. Weak acid–weak base (CH₃COOH vs NH₄OH) titration:
When a weak acid is titrated against a weak base, a pH curve with no sharp (narrow) change in pH (6–7.5) is obtained. Therefore, there is no suitable indicator available for this type of titration.

When you hear the word salt, the first thing that probably comes to mind is the white crystals you sprinkle on food. But chemistry reveals a far richer story. Baking soda that makes your cake rise, washing soda that cleans your clothes, alum that purifies drinking water, blue vitriol used in laboratories — all of these are salts!

Salts are everywhere — in industry, agriculture, medicine, water treatment, and in every biological cell of your body.

A salt is an ionic compound formed when the hydrogen ion (H⁺) of an acid is wholly or partially replaced by a metal ion or an ammonium ion (NH₄⁺).

In practical terms, salts are most commonly produced during a neutralisation reaction between an acid and a base:

Every salt is made up of two essential components:

Salts can be classified as:

• A. Based on composition: Simple salts, Double salts, Complex salts, Acid salts
• B. Based on behaviour in water: Neutral salts, Acidic salts, Basic salts

Simple Salts

Simple salts contain only one type of cation and one type of anion. When soluble simple salts dissolve in water, they generally dissociate into their constituent ions. For example:

Simple salts are further classified based on the nature of the solution they produce when dissolved in water:

Acidic Salts (pH < 7)

These salts produce an acidic solution when dissolved in water. They are typically formed from a strong acid + weak base combination. When dissolved in water, they undergo hydrolysis and increase the concentration of H⁺ ions:

Basic Salts (pH > 7)

These salts produce a basic solution when dissolved in water. They are formed from a weak acid + strong base combination. Upon hydrolysis, they increase the concentration of OH⁻ ions:

Neutral Salts (pH ≈ 7)

These salts produce a neutral solution in water with a pH close to 7. They are formed from a strong acid + strong base combination.

Why neutral? Neither the cation nor the anion undergoes significant hydrolysis, so there is no net change in H⁺ or OH⁻ concentration.

Double Salts

Compounds formed by the combination of two simple salts in a definite ratio, which dissociate completely into their constituent ions in solution.

Example: Mohr’s salt, FeSO₄(NH₄)₂SO₄·6H₂O.

Acid Salts

Acid salts are formed by the partial neutralisation of a polybasic acid. They still retain one or more replaceable hydrogen atoms in their formula.

Compare what happens when sulphuric acid (a dibasic acid) reacts with NaOH:

Complex Salts

Complex salts contain a complex ion along with a counter ion. A complex ion consists of a central metal ion surrounded by ligands — atoms, molecules, or ions that donate a lone pair of electrons to the metal. Common ligands include:

Formation of a Complex Salt — A Striking Example

Consider what happens when ammonium hydroxide (NH₄OH) is added to a copper sulphate (CuSO₄) solution step by step:

Step 1: Small amount of NH₄OH added
A pale blue precipitate of copper hydroxide forms:

Step 2: Excess NH₄OH added
The pale blue precipitate dissolves, and a stunning deep blue complex ion forms!

In a solution containing sulfate ions, this corresponds to the complex salt [Cu(NH₃)₄]SO₄ (Tetraamminecopper(II) sulphate).

Understanding the structure of this complex salt:

Complex salts:

• Contain one or more complex ions enclosed in square brackets [ ].
• Usually exhibit characteristic, vivid colours (e.g., deep blue, red, yellow).
• Show special chemical properties not seen in simple salts.
• Are widely used in analytical chemistry, medicine, and industrial processes.

The reaction of the cation or anion (or both) of a salt with water to produce an acidic or basic solution is called hydrolysis of salts.

Based on the relative strength of the acid and base produced (parent acid and base), the solution may be acidic, basic, or neutral, and the salts are divided into 4 classes.

A. Salt of a Strong Acid and a Strong Base

Halides, nitrates, and sulphates of Na and K, like NaCl, NaNO₃, KNO₃, and K₂SO₄, are strong electrolytes which give cations and anions when dissolved in water. These ions do not interact with water (they do not get hydrolysed). Therefore, the concentration of H⁺ and OH⁻ ions will not change.

Hence, in general, the salt of a strong acid and a strong base does not undergo hydrolysis. Hence, the resulting solution is neutral.

B. Salt of a Weak Acid and a Strong Base

Salts like Na₂CO₃, CH₃COONa, and K₂CO₃ are examples of this type. When these salts are dissolved in water, they produce a strong base and a weak acid.

For example, when Na₂CO₃ is dissolved in water, it dissociates into sodium and carbonate ions:

Thus, OH⁻ is produced → the solution is basic.

C. Salt of Strong Acid and Weak Base

Salts like CaCl₂, NH₄Cl, and FeCl₃ are examples of this category. When such salts are dissolved in water, they produce a strong acid and a weak base.

e.g., NH₄Cl, when dissolved in water, dissociates as:

Thus, H⁺ is produced → the solution is acidic.

D. Salt of a Weak Acid and a Weak Base

Salts like CH₃COONH₄, (NH₄)₂CO₃, (CH₃COO)₂Ca, and HCOONH₄ are examples of this category. When such salts are dissolved in water, they produce weak acids and weak bases.

Therefore, the resulting solution may be neutral, acidic, or alkaline depending on the relative strength of the weak acids and bases produced.

Summary: Hydrolysis of Salts