Ionic equilibrium occurs when the rate at which ions are produced equals the rate at which they are consumed, resulting in a constant concentration of ions in solution.
Ionic equilibrium involves the balance between the concentrations of ions in a solution. This concept is essential in understanding how ions interact and maintain a stable environment in aqueous solutions.
Example: In a solution of acetic acid (CH3COOH), an equilibrium is established between the dissociated ions CH3COO- and H+ and the undissociated acetic acid.
2.0Basic Concepts
Electrolytes vs. Nonelectrolytes:
Electrolytes are substances that dissolve in water and separate into ions, allowing the solution to conduct electricity. (e.g., sodium chloride, NaCl).
Nonelectrolytes: Substances that do not dissociate into ions and do not conduct electricity (e.g., sugar).
Degree of Dissociation:
Strong Electrolytes: Completely dissociate into ions in solution (e.g., NaCl, HCl).
Characteristics:
Degree of dissociation (α): Approximately 1 (100% ionization).
Irreversible ionization.
High electrical conductivity.
Examples:
Strong Acids: HCl, HNO₃, H₂SO₄, HClO₄, HI, HBr.
Strong Bases: NaOH, KOH, Ca(OH)₂, Ba(OH)₂.
Salts: NaCl, KBr, CaCl₂, Na₂SO₄.
Dissociation Reactions:
HCl: HCl → H⁺ + Cl⁻.
NaCl: NaCl → Na⁺ + Cl⁻.
Weak Electrolytes: Partially dissociate in solution (e.g., acetic acid,CH3COOH).
Characteristics:
Degree of dissociation (α): Less than 1 (partial ionization).
Weak electrolytes establish equilibrium between undissociated molecules and ions.
Ionization constant (Kₐ for acids and Kₑ for bases) quantifies strength:
The degree of ionisation (or dissociation) is the fraction of the total number of molecules that ionise (dissociate) into constituent ions in a solution.
α= number of molecules ionised or dissociated total number of molecules taken
For strong electrolytes,𝛼 =1
For weak electrolytes,α<1.
The value of the degree of dissociation (α) depends on the following factors:
Nature of the solute: Different solutes have varying tendencies to dissociate into ions. Strong acids, bases, and salts usually dissociate completely, whereas weak acids and bases only partially dissociate.
Nature of the solvent: The solvent's ability to stabilize ions affects the degree of dissociation. Solvents with high dielectric constants, such as water, can better stabilize ions, leading to higher degrees of dissociation.
Concentration of the solute: The degree of dissociation often decreases with increasing concentration due to the common ion effect and other inter-ionic interactions in the solution.
Temperature: Increasing temperature generally increases the degree of dissociation for most solutes, as higher temperatures provide more energy to overcome the forces holding the ions together.
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3.0Ostwald’s Dilution Law
According to Ostwald, the degree of dissociation (α) of a weak electrolyte is inversely related to the square root of the solution's molar concentration.
Let K be the dissociation constant and C be the molar concentration of the solution.
The relationship is given by Ostwald's Dilution Law:
α=CK
This law indicates that as the concentration C decreases (i.e., the solution is diluted), the degree of dissociation 𝛼 increases. This is particularly applicable to weak electrolytes, which do not fully dissociate in solution.
For a weak electrolyte, the dissociation constant KC can be expressed in terms of the degree of dissociation 𝛼 and the concentration 𝐶:
Kc=(1−α)Cα2
However, for very dilute solutions (where α is much less than 1), this simplifies to:
𝐾≈𝐶𝛼2
Thus, rearranging gives:
α≈CK
This simplified version of Ostwald's Dilution Law demonstrates the inverse relationship between the degree of dissociation and the square root of the molar concentration.
Ka represents the dissociation constant of an acid and Kb represents the dissociation constant of a base. The relationship between Ka and Kb is:
Ka × Kb = Kw = 1.0×10−14 ,where Kw is the ion-product constant of water.
Related Video:
4.0Dissociation Constants and pKa Values for Acids at 25°C
Name
Formula
K a1
pKa1
K a2
pKa2
Acetic acid
CH3COOH
1.75 × 10−5
4.756
Arsenic acid
H3AsO4
5.5 × 10−3
2.26
1.7 × 10−7
6.76
Benzoic acid
C6H5COOH
6.25 × 10−5
4.204
Bromoacetic acid
CH2BrCOOH
1.3 × 10−3
2.9
Carbonic acid
H2CO3
4.5 × 10−7
6.35
4.7 × 10−11
10.33
Chloroacetic acid
CH2ClCOOH
1.3 × 10−3
2.87
Chlorous acid
HClO2
1.1 × 10−2
1.94
Chromic acid
H2CrO4
1.8 × 10−1
0.74
3.2 × 10−7
6.49
Citric acid
C6H8O7
7.4 × 10−4
3.13
1.7 × 10−5
4.76
Cyanic acid
HCNO
3.5 × 10−4
3.46
Dichloroacetic acid
CHCl2CO2H
4.5 × 10−2
1.35
Fluoroacetic acid
CH2FCOOH
2.6 × 10−3
2.59
Formic acid
CH2O2
1.8 × 10−4
3.75
Hydrazoic acid
HN3
2.5 × 10−5
4.6
Hydrocyanic acid
HCN
6.2 × 10−10
9.21
Hydrofluoric acid
HF
6.3 × 10−4
3.2
Hydrogen selenide
H2Se
1.3 × 10−4
3.89
1.0× 10−11
11
Hydrogen sulfide
H2S
8.9 × 10−8
7.05
1 × 10−19
19
Hypobromous acid
HBrO
2.8 × 10−9
8.55
Hypochlorous acid
HClO
4.0 × 10−8
7.4
Hypoiodous acid
HIO
3.2 × 10−11
10.5
Iodic acid
HIO3
1.7 × 10−1
0.78
Iodoacetic acid
CH2ICOOH
6.6 × 10−4
3.18
Nitrous acid
HNO2
5.6 × 10−4
3.25
Oxalic acid
C2H2O4
5.6 × 10−2
1.25
1.5 × 10−4
3.81
Periodic acid
HIO4
2.3 × 10−2
1.64
Phenol
C6H5OH
1.0 × 10−10
9.99
Phosphoric acid
H3PO4
6.9 × 10−3
2.16
6.2 × 10−8
7.21
Pyrophosphoric acid
H4P2O7
1.2 × 10−1
0.91
7.9 × 10−3
2.1
Resorcinol
C6H4(OH)2
4.8 × 10−10
9.32
7.9 × 10−12
11.1
Sulfuric acid
H2SO4
Strong
Strong
1.0 × 10−2
1.99
Sulfurous acid
H2SO3
1.4 × 10−2
1.85
6.3 × 10−8
7.2
5.0Acids and Bases
Arrhenius Concept
Arrhenius Acid: A substance that furnishes hydrogen ions (H⁺) in an aqueous solution.
Example:
HCl → H⁺ + Cl⁻ (strong acid)
CH₃COOH ⇌ CH₃COO⁻ + H⁺ (weak acid)
Arrhenius Base: A substance that furnishes hydroxide ions (OH⁻) in an aqueous solution.
Example:
NaOH → Na⁺ + OH⁻ (strong base)
NH₄OH ⇌ NH₄⁺ + OH⁻ (weak base)
The strength of an acid or base depends on its tendency to furnish H⁺ or OH⁻ ions in solution, respectively.
Water is amphoteric because it can furnish both H⁺ and OH⁻ ions.
H₂O ⇌ H⁺ + OH⁻
The neutralization of an acid and base is a reaction between H⁺ and OH⁻ ions to form water.
H⁺(aq) + OH⁻(aq) ⇌ H₂O(l)
Bronsted-Lowry Concept
A Bronsted-Lowry acid is a proton donor.
A Bronsted-Lowry base is a proton acceptor.
The strength of an acid depends on its tendency to donate protons.
The strength of a base depends on its tendency to accept protons.
Water is amphoteric because it can both donate and accept protons
H2O + H2O ↔ OH− + H3O+
Acid1 Base2 Base1 Acid2
The proton donated by an acid does not exist freely and is always solvated.
CH3COOH + H2O ↔ CH3COO−+H3O+
Each cation behaves as an acid, and each anion behaves as a base. Some can be amphoteric.
Na++2H2O ↔ NaOH+H3O+
Cl−+H2O ↔ HCl + OH−
In all acid-base reactions, two conjugate acid-base pairs are involved.
NH3 + H2O ↔ NH4+ + OH−
Acid1 Base2 Base1 Acid2
More examples of Conjugate Acid-Base Pairs:
HCl ↔ H++Cl−
C6H5NH2+H+↔C6H5NH3 +
H2O+Al(OH)3↔Al+3(𝑎𝑞)+4H2O
According to Bronsted-Lowry, all acid-base reactions involve the transfer of protons between two conjugate acid-base pairs.
Lewis Concepts
Lewis extended the Bronsted concept to include the following:
Lewis Acid: An electron pair acceptor.
Lewis Base: An electron-pair donor.
Acid-base reactions involve sharing an electron pair from the base to the acid, forming a coordinate complex.
Example:
BF3+NH3→(H3N→BF3)
BF3+F−→[BF4]−
Ag+ + 2CN−→[Ag(CN)2]−
Lewis acids accept electron pairs, while Lewis bases donate them, forming coordinate complexes in acid-base reactions. The strength of these acids and bases depends on their charge, size, and electronegativity.
Acid Strength and pH
The acid strength is defined by its pH:
pH = −log[H+]
Base Strength and pOH
The strength of a base is defined by its pOH:
pOH = −log[OH−]
The relationship between pH and pOH is:
pH + pOH =14
6.0Buffer Solution
A buffer solution maintains its pH relatively stable when small quantities of strong acid or base are introduced. It is also called reserve acidity or basicity, demonstrating buffer action.
Features:
The solution maintains a specific pH.
The pH remains stable over time.
pH does not change with dilution.
Adding small amounts of strong acid or base causes only slight, unnoticeable pH changes.
7.0Henderson-Hasselbalch Equation
By using the Henderson-Hasselbalch equation, the pH of a buffer solution can be determined
pH=pKa+log10[HA][A−]
Where: A- is the concentration of the conjugate base. HA is the concentration of the acid.
For a primary buffer (weak base + conjugate acid), the pOH is calculated as:
pOH=pKb+log([B−][BH])
Where: B− is the concentration of the base. BH is the concentration of the conjugate acid.
8.0Salts and Their Ionization
When acids react with bases, salts are generated. Upon ionization in solution, the ions of these salts may subsequently interact with water, resulting in the formation of acidic, basic, or neutral solutions depending on the properties of the ions involved.
Type of Solution
Cations
Anions
pH
Acidic
From weak bases
From strong acids
< 7
Basic
From strong bases
From weak acids
> 7
Neutral
From strong bases
From strong acids
= 7
Common Ion Effect
When equilibrium establishes between an ionic compound and its ions, adding a common ion will shift the equilibrium to consume that ion, according to Le Chatelier’s principle.
The common ion effect is used in:
Purification of Common Salt: Adding a common ion reduces solubility, helping to purify salt.
Salting Out of Soap: Enhances precipitation of soap from solution.
Qualitative Analysis:
Group II Radicals: Precipitated in the presence of HCl, which suppresses S2− ion concentration, sufficient to precipitate only Group II radicals.
Group III Radicals:NH4OH is added in the presence of NH4Cl to avoid the precipitation of Group V radicals.
Isohydric Solutions
Isohydric solutions are those where the concentration of common ions(e.g.,OH−−ions in Ca(OH)2 and Ba(OH)2 solutions) is the same. When mixed, the degree of dissociation of either electrolyte does not change.
9.0Solubility and Solubility Product
Solubility: The maximum amount of a substance that dissolves in a fixed amount of a solvent at a specific temperature.
For an ionic compound XY that dissociates in water:
𝑋𝑌↔𝑋++𝑌−
The solubility product (Ksp) is expressed as:
Ksp= [X+][Y−]
Ionic Product (I.P.): It is the product of the ions' concentrations in the solution.
Unsaturated Solution: I.P.< Ksp
Supersaturated Solution: I.P.> Ksp (precipitate will form)
Saturated Solution: I.P.= Ksp (no precipitate formation)
10.0Solved Examples
Ques 1.In the reaction equilibrium P + Q ↔ R + S, what will happen to the concentrations of P, Q, and S if the concentration of R is increased?
Ans When the concentration of R is increased in the reaction equilibrium P + Q ↔ R + S, the system will respond according to Le Chatelier's Principle. This principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change. Equilibrium constant of the reaction can be shown as:
α≈CK
If the concentration of product is increased then the concentration of other components changes which decreases the concentration of R and vice – versa. Thus, increasing the concentration of R shifts the equilibrium to the left, leading to an increase in the concentrations of P and Q and a decrease in the concentration of S.
Ques 2. Write the relationship between molar Solubility (S) and solubility product Ksp for PbI2
Ans The relationship between molar solubility (𝑆) and the solubility product (Ksp) for lead(II) iodide(PbI2):
Dissociation Equation can be written as: PbI2(s)↔Pb2+(aq)+2I−(aq)
