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Ionic Equilibrium
Electrolytes and Non-Electrolytes
Substances are classified as electrolytes or non-electrolytes based on their ability to conduct electricity in aqueous solution or molten state. Electrolytes conduct electricity due to the presence of ions, while non-electrolytes do not.
Electrolytes: Substances that conduct electricity in aqueous or molten state (e.g., salts, strong acids, strong bases). Non-Electrolytes: Substances that do not conduct electricity (e.g., glucose, sucrose). Strong Electrolytes: Undergo 100 % ionization in solution (e.g., strong acids, strong bases, most salts). Weak Electrolytes: Undergo partial ionization (e.g., weak acids, weak bases, sparingly soluble salts).
Acid-Base Theories
Three main acid-base theories are used to define acids and bases: Arrhenius, Bronsted-Lowry, and Lewis. Each theory has its own scope and limitations.
Arrhenius Theory: Acid produces in aqueous solution; base produces. Limitation: Only applies to aqueous solutions. Bronsted-Lowry Theory: Acid is a proton ( ) donor; base is a proton acceptor. Conjugate acid-base pairs differ by one proton. The conjugate base of a strong acid is weak, and vice versa. Lewis Theory: Acid is an electron pair acceptor; base is an electron pair donor. Explains acid-base behavior beyond aqueous solutions. Limitation: Does not indicate acid/base strength or always explain reaction speed.
Ionic Product of Water and Scale
Water is a weak electrolyte and undergoes self-ionization, producing and ions. The ionic product ( ) is temperature dependent and forms the basis for the scale.
At C: At C: For neutral water: M at C
at C; at C If , solution is acidic; if , basic; if , neutral (at C) For ,
Ionization Constants of Weak Acids and Bases
Weak acids and bases partially ionize in solution. Their ionization is characterized by equilibrium constants ( for acids, for bases).
For weak acid : For weak base : Degree of ionization ( ) for weak acid:
for weak acids For weak bases, analogous formulas apply with and
Common Ion Effect
The common ion effect refers to the decrease in the degree of ionization of a weak electrolyte upon addition of a strong electrolyte containing a common ion. This is explained by Le Chatelier's principle.
When a strong electrolyte with a common ion is added, the equilibrium shifts backward, decreasing ionization of the weak electrolyte. In calculations, the concentration of the common ion is taken from the strong electrolyte. For a weak acid with added , is dominated by. For a weak base with added , is dominated by.
Buffer Solutions
Buffer solutions resist changes in upon addition of small amounts of acid or base. They are of two main types: acidic and basic buffers.
Acidic Buffer: Mixture of weak acid and its salt with a strong base (e.g., ). Basic Buffer: Mixture of weak base and its salt with a strong acid (e.g., ). of acidic buffer: of basic buffer: Buffer capacity: Number of moles of acid or base required to change by one unit in one liter. Buffer range: range over which buffer action is effective, typically. Maximum buffer action when or.
Solubility Equilibria and Solubility Product ( )
Sparingly soluble salts establish equilibrium between undissolved solid and ions in solution. The solubility product ( ) is the product of ion concentrations at saturation.
For : For $Mg(OH) 2 K{sp} = [Mg^{ 2 +}][OH^-]^ 2 $ Relationship between and solubility ( ): For : For : For : Ionic product (IP): Product of ion concentrations at any moment; compare with to predict precipitation. If : Unsaturated, more salt can dissolve. If : Saturated, equilibrium. If : Precipitation occurs.
Common Ion Effect on Solubility
The presence of a common ion decreases the solubility of a sparingly soluble salt. This is due to the shift in equilibrium according to Le Chatelier's principle.
For in solution, increases, so decreases. Solubility in presence of common ion: , where is the concentration of the common ion from the strong electrolyte. For in solution: , solve for.
Special Cases and Exam Pitfalls
Certain cases require careful attention to avoid mistakes in calculations and conceptual understanding.
For very dilute strong acids/bases ( M), always include water's ionization in calculation. For buffer solutions, remains nearly constant upon small additions of acid/base, but not for large additions. In salt hydrolysis, always identify the nature of the salt to select the correct formula for. For solubility in presence of complex formation (e.g., with ), solubility increases due to complex ion formation. When mixing solutions, always use normality and volume to find final concentrations before calculation.
Table: Salt Type and Nature
Salt Type (^) Nature Formula for Calculation
Strong Acid + Strong Base (^) Neutral ( )
Strong Acid + Weak Base (^) Acidic ( )
Weak Acid + Strong Base Basic ( )
Weak Acid + Weak Base Slightly acidic/basic
