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Ionic Equilibrium is a crucial topic in chemistry that deals with the behavior of ions in solutions and how they establish equilibrium under different conditions. This chapter explains concepts such as acids, bases, salts, pH, buffer solutions, solubility equilibrium, and common ion effect. Understanding ionic equilibrium is essential for solving numerical problems and conceptual questions in NEET.
According to Arrhenius, acids are substances that increase the concentration of hydrogen ions (H⁺) in an aqueous solution, while bases increase the concentration of hydroxide ions (OH⁻). This definition applies only to aqueous solutions and does not consider acid-base behavior in non-aqueous solvents.
The Bronsted-Lowry concept defines acids as proton donors and bases as proton acceptors. This definition is more general and applies to both aqueous and non-aqueous solutions. It also introduces the idea of conjugate acid-base pairs, where an acid donates a proton to form its conjugate base, and a base accepts a proton to form its conjugate acid.
Lewis defined acids as substances that can accept an electron pair, while bases are substances that can donate an electron pair. This concept is the broadest among the three and includes acid-base reactions that do not involve hydrogen or hydroxide ions.
Strong acids and bases ionize completely in water, meaning they dissociate fully into their respective ions.
Weak acids and bases ionize partially in water, establishing an equilibrium between the ionized and non-ionized forms.
The degree of ionization refers to the fraction of an acid or base that dissociates in solution. Strong acids and bases have a high degree of ionization, while weak acids and bases have a lower degree.
Dilution increases the ionization of weak acids and bases because the equilibrium shifts to produce more ions in solution. This principle is explained by Le Chatelier’s principle.
The pH scale is used to measure the acidity or alkalinity of a solution. A lower pH indicates a more acidic solution, while a higher pH indicates a more basic solution.
Acidic solutions have a pH less than 7.
Neutral solutions have a pH of 7.
Basic solutions have a pH greater than 7.
The pOH scale is used to measure the concentration of hydroxide ions in a solution. It is related to pH, and both values help determine the nature of a solution.
A buffer solution is a solution that resists changes in pH when small amounts of acid or base are added. Buffers are essential in maintaining the stability of pH in biological and chemical systems.
Acidic buffer: Contains a weak acid and its salt with a strong base (e.g., acetic acid and sodium acetate).
Basic buffer: Contains a weak base and its salt with a strong acid (e.g., ammonia and ammonium chloride).
Used in biological systems, such as maintaining blood pH.
Employed in chemical laboratories for experiments requiring constant pH.
Found in industrial processes like fermentation and drug formulation.
Salt hydrolysis refers to the reaction of salts with water to produce acidic or basic solutions. It depends on the nature of the acid and base from which the salt is formed.
Salt of a strong acid and a strong base: Neutral (e.g., NaCl in water).
Salt of a strong acid and a weak base: Acidic (e.g., NH₄Cl in water).
Salt of a weak acid and a strong base: Basic (e.g., CH₃COONa in water).
Salt of a weak acid and a weak base: pH depends on the relative strengths of the acid and base.
Used in water treatment processes.
Important in biological reactions and enzyme activities.
Helps in the formulation of pharmaceuticals.
The common ion effect occurs when the addition of an ion, already present in solution, suppresses the ionization of a weak acid or base. This effect is crucial in maintaining equilibrium in buffer solutions and solubility control.
The solubility of AgCl decreases when NaCl is added due to the presence of common chloride ions.
The ionization of acetic acid is suppressed in the presence of sodium acetate.
Used in precipitation reactions to control solubility.
Helps in pH regulation in buffer solutions.
Used in qualitative chemical analysis to separate ions.
Solubility refers to the ability of a substance to dissolve in a solvent. In ionic equilibrium, solubility is influenced by temperature, pressure, and the presence of common ions.
Nature of solute and solvent: Polar solvents dissolve polar solutes, while non-polar solvents dissolve non-polar solutes.
Temperature: Solubility of most solids increases with temperature, while solubility of gases decreases.
Pressure: Affects the solubility of gases but has little effect on solids and liquids.
Essential in pharmaceutical formulations.
Used in water purification processes.
Helps in understanding geological formations like limestone dissolution.
Blood maintains a constant pH through buffer systems.
Enzymatic reactions depend on specific pH levels.
Acid-base balance is crucial for metabolic functions.
Food industry: pH control in fermentation processes.
Cosmetic industry: pH regulation in skincare products.
Agriculture: Soil pH management for better crop growth.
Acid rain formation and its neutralization.
Water treatment and purification processes.
Marine life depends on stable pH conditions in oceans.
Ionic Equilibrium is a fundamental concept in chemistry that explains how ions behave in solutions and establish equilibrium. Understanding acids, bases, pH, buffers, salt hydrolysis, common ion effect, and solubility equilibrium is essential for NEET aspirants. Mastering these topics helps in solving numerical problems and understanding real-life applications in medicine, industry, and environmental science.
