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IONIC Equilibrium In One Shot - JEE/NEET/Class 11th Boards | Victory Batch

Introduction

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Equilibrium is the state where measurable properties like pressure, volume, density, and color remain unchanged over time. Ionic equilibrium in weak electrolytes is established by the balance between unionized molecules and the ions produced through partial ionization. The reversible nature of this process sets it apart from the complete ionization seen in strong electrolytes, emphasizing a dynamic yet stable chemical balance.

Electrolyte

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Electrolytes are categorized by how completely they dissociate into ions, with strong electrolytes fully ionizing and weak electrolytes only partially doing so. The degree of dissociation measures the extent of ion formation, reaching 100% for strong electrolytes while remaining incomplete for weak ones. Strong electrolytes include acids like HCl and H2SO4, bases such as NaOH and KOH, and salts like KCl, whereas weak examples include acids like CH3COOH and corresponding weak bases.

Arrhenius Theory of Electrolytic Dissociation

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Acid and Base Definitions via Ion Donation Arrhenius theory defines acids as substances that donate H+ ions in water and bases as those that release OH- ions. When HNO3 dissolves in water, it dissociates into H+ and NO3- ions, exemplifying an acid’s behavior. Similarly, NaOH splits into Na+ and OH- ions in solution, confirming its basic characteristic.

Limited Scope for Ion-Based Definitions The theory is restricted to aqueous solutions, limiting its broader applicability. Compounds without inherent H+ ions, like CuSO4, or those lacking OH- ions, such as NH3, do not fit neatly into this ion donation model. These exceptions underscore the constraints of defining acid and base properties solely through electrolytic dissociation.

Bronsted Lowry Concept

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Acids Donate; Bases Accept This theory defines acids as substances that donate H+ ions and bases as those that accept them. HCl loses an H+ ion to convert into Cl– while water converts into H3O+ by accepting the proton. A similar process occurs when HCl reacts with NH3, forming NH4+ through proton donation.

Conjugate Pairs and Amphoteric Water Strong acids yield weak conjugate bases, and strong bases produce weak conjugate acids, with the reverse relationship holding for weaker substances. Water, being amphoteric, adapts to act as an acid or a base depending on the reaction context. Examples include water acting as a base with acetic acid and as an acid with ammonia.

Limitations in Non-Protonic Behavior The theory cannot explain acidic behavior in compounds lacking H+ ions, such as CuSO4 and AlCl3. Similarly, substances that cannot accept protons, like calcium oxide or magnesium oxide, illustrate its inability to account for basic characteristics. These limitations highlight scenarios where the H+ transfer model is insufficient.

Lewis Acids

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Lewis acids are defined by their ability to accept electrons, which is demonstrated through their positive charge that attracts electron pairs. Simple cations like H⁺, Ag⁺, and Fe²⁺ naturally behave as Lewis acids because they seek electrons to achieve neutrality. Electron-deficient molecules, such as BF₃ and AlCl₃, also act as Lewis acids by completing their octet with incoming electron pairs. Additionally, molecules featuring central atoms capable of expanding their octet via available d orbitals broaden the scope of Lewis acidity through increased bonding capacity.

Lewis Bases

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A Lewis base is defined as a species that donates electrons, marking a shift from the traditional proton-focused acid-base theory. Species with available lone pairs, negatively charged ions, or multiple bonds between atoms of varying electronegativities have the capacity to donate electrons. In contrast, an electron-deficient species like BF3 accepts electrons from an electron-rich counterpart such as NH3, forming a coordinate covalent bond. This mechanism of electron donation and acceptance stabilizes both participants and illustrates the fundamental nature of Lewis acid-base interactions.

Dissociation constant for weak acid and weak base

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Weak acid HA dissociates into H+ and A-, with its dissociation constant Ka defined as the ratio of the concentrations of the products to the undissociated acid. Similarly, a weak base like BOH splits into B+ and OH-, and its dissociation constant Kb is determined by dividing the concentration of products by the concentration of the base. This clear method of representing dissociation provides a consistent framework to understand the behavior of both acids and bases, paving the way for related concepts such as the ionic product of water.

Ionic Product of Water

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Water naturally dissociates into hydrogen and hydroxide ions, leading to an equilibrium expression where the product of their concentrations defines the ionic product, Kw. At 25°C, both [H⁺] and [OH⁻] are 10⁻⁷ M, making Kw equal to 10⁻¹⁴. Increasing temperature elevates Kw, reflecting the sensitivity of this equilibrium to thermal changes.

pH scale

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The pH scale quantifies a solution's acidity by expressing the hydronium ion concentration as a negative logarithm, encapsulated by the formula pH = -log[H+]. This relationship enables the direct calculation of ion concentration using 10^(-pH) without ambiguity in the formula's application. A parallel definition applies to pOH, which is determined from the negative logarithm of the hydroxide ion concentration. At 25°C, the scale ranges from 0 to 14, where a value of 7 signifies neutrality, values below indicate acidity, and values above indicate basicity.

Common Ion Effect

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Common Ion Effect Drives Equilibrium Shifts Acetic acid dissociates into acetate and hydrogen ions, and the addition of a strong acid introduces extra hydrogen ions into the reaction. The increase in hydrogen ion concentration pushes the equilibrium backward according to Le Chatelier’s principle. This dynamic illustrates how a common ion can force an equilibrium to readjust and restore balance.

Buffer Creation Modulated by Common Ions Mixing a weak acid with its salt forms an acidic buffer, as the extra acetate ion lowers the free hydrogen ion concentration and raises the pH. In contrast, combining a weak base with its salt produces a basic buffer where the common ammonium ion diminishes the hydroxide ions, resulting in a lowered pH. These examples demonstrate how the common ion effect is harnessed to fine-tune pH by shifting equilibria, linked by the relationship between acid and base dissociation constants and the ionic product of water.

Factors affecting Acidic Strength

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A weak hydrogen-acid bond facilitates the release of hydrogen ions, making the acid stronger. The bond’s polarity, enhanced by a greater electronegativity difference, further promotes ion dissociation. Comparative examples among halogen acids and other elements show that increased electronegativity differences augment bond polarity and, consequently, acid strength.

Buffer Solution

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Buffer Solutions Stabilize pH in Water When acids or bases are added to water, the pH naturally shifts, and buffer solutions are formulated to counteract these changes. The technique involves preparing either an acidic buffer using a weak acid and its conjugate salt or a basic buffer employing a weak base with its corresponding salt. This balancing act ensures the solution maintains stability even when external substances alter the pH.

Strategic pH Calculations in Varied Acid-Base Systems The pH of an acidic buffer is determined using the equation pH = pKa + log ([salt]/[acid]), while a basic buffer is addressed through its pOH form, pOH = pKb + log ([salt]/[base]). Four specific scenarios emerge: combining strong acid and strong base yields a neutral pH of 7; a mix of strong acid with weak base gives a pH below 7 using a subtractive adjustment; weak acid paired with strong base results in a pH above 7 via an additive process; and weak acid with weak base approximates pH as 7 plus half the difference between pKa and pKb. These formulas provide critical quantitative insights for tackling numerical acid-base challenges.

Solubility and Solubility Product

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Defining Solubility and Salt Composition Solubility is introduced as the measure of the number of moles of a solute present in one liter of a saturated solution. A salt represented as AxBy splits into its constituent ions, with the cation noted first and the anion second, ensuring charge balance. Balancing the charges within the formula is critical as it directly influences how the salt dissociates in aqueous solution.

Establishing the Solubility Product Constant The equilibrium in a saturated solution is described by the solubility product constant, Ksp, which is derived from the product of the concentrations of the dissolved ions. For a simple salt like AgCl, each ion's concentration is denoted by s, making Ksp equal to s multiplied by s, or s². The approach demonstrates that once the equilibrium concentrations are known, solving for solubility is straightforward by taking the appropriate root of the Ksp expression.

Balancing Complex Salts and Practical Derivations For more complex salts such as CaF₂, the dissolution must account for multiple ions, leading to unequal stoichiometric contributions. Here, CaF₂ dissociates into one Ca²⁺ and two F⁻ ions, so the solubility product is expressed as 4s³, with the concentration of F⁻ adjusted to 2s. This derivation underscores the importance of balancing the ionic equation and offers clear formulas for calculating solubility, emphasizing repeated practice to master numerical challenges.