Ionic equilibrium deals with the equilibrium established between ions in a solution. It plays a critical role in various chemical, industrials and biological processes.
What is Ionic Equilibrium?
When ionic compounds dissolve in water, they dissociate into positive and negative ions. Ionic equilibrium is the state where the rate of dissociation of the ions equals the rate of recombination, and stable concentration of ions in the solution. This concept is particularly important for electrolytes, substances that can conduct electricity when dissolved in water.
Ionic Equilibrium in Solution
Electrolytes can be classified into:
1.Strong electrolytes: These completely dissociate into ions in water, e.g., NaCl, HCl, and KOH.
2. Weak electrolytes: These only partially dissociate, e.g., CH₃COOH, NH₄
Key Concepts in Ionic Equilibrium
1.Degree of Dissociation (α): The fraction of the total number of molecules that dissociate into ions is called the degree of dissociation. It is expressed as:
α = Number of dissociated molecules / Total number of molecules
Strong electrolytes have α ≈ 1 while weak electrolytes have α < 1.
2. Equilibrium Constant (K): For weak electrolytes, the equilibrium constant (Kc) calculate the extent of dissociation. For an electrolyte AB dissociating as:
AB ↔ A+ + B− The equilibrium constant is: Kc = [A+] [B−] / [AB]
Here, [A⁺], [B⁻], and [AB] represent the molar concentrations of the respective species.
3. Ionisation of Water (Kw): Pure water undergoes a slight ionisation into H⁺ and OH⁻ ions: H2O ↔ H+ + OH−
The equilibrium constant for this process is called the ionic product of water (Kw): Kw = [H+] [OH−] At 25°C, Kw = 1.0 × 10−14.
4. pH and pOH:
pH: A measure of hydrogen ion concentration, defined as: pH = −log[H+]
pOH: A measure of hydroxide ion concentration:
pOH = −log[OH−]
Relationship: pH + pOH = 14
5. Buffer Solutions (A solution whose pH is not altered to any great extent by the addition of small quantities of either an acid or base is called buffer solution. Buffer is also defined as the solution of reserve acidity or alkalinity which resists change of pH upon the addition of a small amount of acid or alkali):
6 Buffer solutions resist changes in pH upon the addition of small amounts of acids or bases. They are made of:
A weak acid and its salt (e.g., CH₃COOH and CH₃COONa)
A weak base and its salt (e.g., NH₃ and NH₄Cl)
The pH of a buffer can be calculated using the Henderson-Hasselbalch equation:
For acidic buffer: pH = pKa + log[Salt] / [Acid]
For basic buffer: pOH = pKb+log[Salt] / [Base]
7. Common Ion Effect: When a common ion is added to a solution of a weak electrolyte, the degree of dissociation of the weak electrolyte decreases. like, adding NaCl to a solution of HCl suppresses the ionisation of HCl due to the presence of common
Cl.−
Applications of Ionic Equilibrium
1.Acid-Base Reactions: Ionic equilibrium helps to predict the outcome of reactions between acids and bases.
2. Biological Systems: Processes like enzyme activity and nutrient absorption depend on pH, which is governed by ionic equilibrium.
3. Industrial Applications: Buffer solutions are widely used in chemical manufacturing, pharmaceuticals, and food industries to maintain stable pH.
4. Environmental Science: The ionic equilibrium of natural waters affects aquatic life and water quality.
Key points to remembers:–
Substances that conduct the electricity in their aqueous solutions or in molten state are called electrolytes
Strong electrolytes are completely ionised in aqueous solutions and weak electrolytes are partially ionised in aqueous solutions.
In weak electrolytes, an equilibrium is established between ions and unionised molecules, leading to an ionic equilibrium in the aqueous solution. All acids, base and salts may be classified as weak or strong electrolytes.
Note :-
Ionic equilibrium is a connects of the behaviour of ions in solution to deal with phenomena. In this, we deal with acid-base, buffer action, and various industrial and biological processes.
Ionic equilibrium is the dynamic equilibrium established between the ions and the undissociated molecules in a solution of electrolytes. This equilibrium occurs when the rates of dissociation and recombination of ions are equal, resulting in a stable concentration of ions.
Strong electrolytes: It is completely dissociate into ions in water, e.g., NaCl, HCl, and KOH.
Weak electrolytes: It is only partially dissociate into ions, e.g., CH₃COOH and NH₄OH.
The extent of dissociation is calculate by the degree of dissociation (α).
The ionic product of water (Kw) is the equilibrium constant for the ionisation of water: H2O ↔ H+ + OH−
At 25°C, Kw is equal to 1.0×10−14. It is expressed as:
Kw = [H+] [OH−]
The pH of a solution is calculated using the formula:
pH = −log[H+] For basic solutions, pOH is calculated as:
pOH = −log[OH−] The relationship between pH and pOH is:
pH + pOH = 14(at 25°C)
The common ion effect is the suppression of the ionisation of a weak electrolyte due to the presence of a common ion from another source. For example, adding NaCl to a solution of HCl reduces the ionisation of HCl because of the common Cl− ion.
The Henderson-Hasselbalch equation is used to calculate the pH of buffer solutions.
For acidic buffers: pH = pKa + log[Salt] / [Acid]
For basic buffers: pOH = pKb + log[Salt] / [Base]
Buffer solutions resist changes in pH upon the addition of small amounts of acids or bases. They are essential in:
Maintaining pH stability in biological systems, such as blood.
Industrial processes like fermentation.
Chemical reactions requiring controlled pH conditions.