EQUILIBRIUM-

  • Equilibrium- It is a state of balance where everything is stable, nothing changes.
  • It is represent is double arrow(⇌) in chemical reaction. It showing the process is reversible and happening in both direction.
  • Equilibrium Mixture- Mixture of reactant and product in equilibrium state. That means ratio of reactant and product is constant.
  • Dynamic equilibrium- When 2 opposite processes happening at same time and have a same rate. (ex.- reactant become product and product reverse to become reactant, then overall amount are same)
  • Chemical equilibrium- When a chemical reaction reaches a point where the amount of reactant and product remain constant because the forward and reverse reactions happen at same rate.
  • Ionic equilibrium- When the rate of ionization equal to the rate of recombination in a solution, maintaining a stable concentration of ions.

EQUILIBRIUM IN PHYSICAL PROCESSES-

  1. Solid-Liquid Equilibrium-
  • It state that where ice and water coexist at 0°C, with the rate of melting equal to the rate of freezing.
  • For any pure substance at atm (atmospheric pressure), the temperature at which the solid and liquid phases are at equilibrium is called normal melting or freezing point of substance.
  • Both process occur at same rate, so the amount of ice and water remains constant.

2. Liquid-Vapour Equilibrium-

  • In this equilibrium state, the liquid and its vapor state present in closed system (or container) in equilibrium position.
  • It this condition, the rate of liquid to vapour and vapour to again liquid is same.
  • It is in cyclic form.
  • Factors affecting liquid-vapour equilibrium-
    • Temperature,
    • Nature of liquid and
    • Surface area.

3. Solid-Vapour Equilibrium-

  • In this equilibrium state, the solid and its vapour exist in same closed system (or container) and both have dynamic balance.
  • During this equilibrium, solid particle having vaporize and vapour particles again having solidify with same rate. Having same rate, the overall composition is same.
  • Condition required for this equilibrium-
    • Closed system,
    • Constant temperature and
    • Dynamic nature.

SOLID IN LIQUID-

  • In this state, the rate of solid dissolved in liquid is equal to the rate at which the dissolved ions or molecules re-precipitate back into the solid form.
  • This type of equilibrium is dynamic.
  • Ex- NaCl(s) ⇋ Na+(aq) + Cl–(aq)
  • Characteristic of this equilibrium-
    • Saturated solution- The liquid contains maximum amount of dissolved solid at a given temperature.
    • Dynamic nature- The processes of dissolution and precipitation are continuous having same rate.
    • Dependence on temperature- The solubility of most solids increases with temperature, then shifting the equilibrium.

GASES IN LIQUID-

  • Physical Equilibria:
  • 1. Opening Soda Bottles:
    • Carbon dioxide gas escapes due to a decrease in solubility at lower pressures.
    • Governed by Henry’s Law: The mass of gas dissolved in a liquid is proportional to the pressure of the gas above it.
    • Higher temperatures reduce gas solubility.
    • Soda turns “flat” as CO₂ escapes to reach equilibrium with atmospheric pressure.
  • 2. Equilibria Types and Observations:
    • Solid-Liquid Equilibrium:
      • Melting point occurs at constant temperature and pressure.
      • Mass of both phases remains constant in a closed system.
    • Liquid-Vapour Equilibrium:
      • Vapour pressure is constant at a given temperature.
    • Solid in Liquid (Solubility):
      • Solubility is temperature-dependent.
    • Gas in Liquid (Henry’s Law):
      • Gas concentration in a liquid is proportional to its partial pressure above the liquid.

Characteristics of Physical Equilibria-

  1. Conditions for Equilibrium:
    • Requires a closed system at a constant temperature.
    • Opposing processes occur at the same rate.
    • Dynamic but stable condition.
  2. Observables:
    • System’s measurable properties (e.g., pressure, temperature) remain constant at equilibrium.
    • Quantitative measures, such as vapour pressure or solubility, indicate the extent of equilibrium.

Chemical Equilibrium (Dynamic Nature):

  1. Definition:
    • Occurs when the forward and reverse reaction rates are equal.
    • Concentrations of reactants and products remain constant over time.
    • Dynamic because reactions continue at the molecular level.
  2. Examples:
    • General Reaction: A+B⇌C+D
    • Forward reaction slows as reactants decrease.
    • Reverse reaction speeds up as products increase.
    • Equilibrium can be achieved starting from either reactants or products.
  3. Haber Process (Synthesis of Ammonia):
    • Reaction: N2(g)+3H2(g)⇌2NH3(g)
    • Demonstrates equilibrium by showing a constant mixture composition.Use of isotopes (e.g., deuterium, D2​) confirms that forward and reverse reactions persist even at equilibrium.
  4. Isotope Evidence:
    • Mixing H2 and D2 , then results in all possible isotope combinations (e.g., NH3, NH2​D, NHD2, ND3).
    • Demonstrates ongoing reaction processes despite constant macroscopic properties.
  5. Key Characteristics of Chemical Equilibrium:
    • Can be reached from either side of a reaction.
    • No net change in the system’s composition at equilibrium.
    • Reactions remain dynamic, evidenced by molecular-level exchanges.

Key Concepts for Competitive Exams:

  1. Henry’s Law:
    • Solubility of a gas increases with pressure and decreases with temperature.
    • Formula- C=Kp​⋅ where C is concentration, P is pressure, and Kp​ is the proportionality constant.
  2. Equilibrium Constants:
    • Kc​: Ratio of concentrations of products to reactants at equilibrium.
    • Kp​: Pressure-based constant for gaseous reactions.
    • Equilibrium position depends on initial conditions but achieves the same constant under identical conditions.
  3. Le Chatelier’s Principle:
    • Predicts the shift in equilibrium due to changes in pressure, concentration, or temperature.
  4. Dynamic Nature of Equilibrium:
    • Reactants and products continuously interconvert.
    • Balance between forward and reverse reaction rates maintains constancy.

Practice Tips:

Understand Physical vs. Chemical Equilibria:

  • Compare how systems behave under varying conditions.
  • Focus on constants like vapour pressure, solubility, or equilibrium constants.

Memorize Key Laws:

  • Henry’s Law, Le Chatelier’s Principle, and laws governing phase transitions.

Apply to Real-World Examples:

  • Relate concepts to everyday phenomena, such as soda bottles or ammonia synthesis.

Law of Chemical Equilibrium and Equilibrium Constant:

Equilibrium Mixture

  • A mixture containing both reactants and products at equilibrium is termed an equilibrium mixture.
  • Key questions regarding equilibrium mixtures:
    1. Relationship between concentrations of reactants and products at equilibrium.
    2. Determination of equilibrium concentrations from initial values.
    3. Factors influencing the equilibrium composition (important for industrial processes like synthesis of H2, NH3CaO, etc.).

Law of Mass Action

  • Proposed by Guldberg and Waage in 1864, stating that at equilibrium, the ratio of the product of concentrations of products to the product of concentrations of reactants, each raised to their respective stoichiometric coefficients, is constant. This is the equilibrium constant (Kc_cc​).

Equilibrium Constant Expression-

  • For a general reaction aA + bB ⇌ cC + dD:
    • Kc = [C]c[D]d​ / [A]a[B]
  • Equilibrium concentrations are used in the expression, denoted in mol/L.
  • The equilibrium constant depends only on the reaction’s temperature.

Important Properties of Kc​:

  1. Reversing the Reaction:
    • Kc​ for the reverse reaction is the reciprocal of Kc for the forward reaction:
      • K’c​= 1 / Kc
  2. Changing Stoichiometric Coefficients:
    • Multiplying the equation by a factor nnn changes the equilibrium constant as follows:
      • K’c  ​=(Kc​)n
  3. Phases and States:
    • While writing Kc​, phases (solid, liquid, gas) are usually omitted, but only species in the gaseous or aqueous phase are included.

Heterogeneous Equilibrium-

  1. Definition:
    • Equilibrium in a system involving more than one phase is called heterogeneous equilibrium.
  2. Examples:
    • Water and vapor equilibrium:
      • H2​O(l) ⇌ H2​O(g)
      • Two phases: liquid and gas.
    • Solid-liquid solution equilibrium:
      • Ca(OH)2​(s) ⇌ Ca2+(aq) + 2OH−(aq)
  3. Key Characteristics:
    • Involves pure solids or liquids, whose concentrations are constant and omitted from the equilibrium constant expression.
    • Concentrations of gases [X(g)]and aqueous solutions [X(aq)] vary with the amount of substance.
  4. Equilibrium Expressions:
    • For thermal dissociation of calcium carbonate:
      • CaCO3​(s) ⇌CaO(s) + CO2​(g)
      • Equilibrium constant:
        • Kc​ =[CO2​] orKp​ = Pco2
        • At 1100 K: Kp=2 (Pa)
  5. General Rules:
    • The equilibrium constant (K) excludes concentrations of pure solids or liquids.
    • Units of Kc and Kp:
      • Depend on the reaction stoichiometry.
      • Are dimensionless when concentrations/pressures are expressed relative to their standard states.

Applications of Equilibrium Constant

Equilibrium Constant (K):

  • The equilibrium constant (K) reflects the balance between reactants and products in a reaction at equilibrium.
  • Changing the reaction equation by multiplying or dividing by a number affects the equilibrium constant.

Applications of Equilibrium Constant:

  • Predicting the extent of a reaction: The value of K shows whether the products or reactants dominate at equilibrium. A large K indicates more products, while a small K suggests more reactants.
  • Predicting the direction of the reaction: The reaction will move forward or backward to restore equilibrium.
  • Calculating equilibrium concentrations: If we know the initial concentrations and K, we can calculate the concentrations of reactants and products at equilibrium.

Predicting the Extent of a Reaction:

  • If K > 10³, products are favored.
  • If K < 10⁻³, reactants are favored.
  • If 10⁻³ < K < 10³, both products and reactants are present in significant amounts.

Equilibrium Composition:

  • For the reaction to reach equilibrium, the concentrations of reactants and products must balance.
  • If Q (reaction quotient) ≠ K, the reaction will shift to restore equilibrium.

Predicting the Direction of the Reaction:

  • Q > K: The reaction will go in the reverse direction (towards reactants).
  • Q < K: The reaction will go in the forward direction (towards products).
  • Q = K: The system is already at equilibrium, and no change occurs.

Calculating Equilibrium Concentrations:

  • Steps to calculate equilibrium concentrations:
    1. Write the balanced equation.
    2. Create a table for initial concentrations, changes, and equilibrium concentrations.
    3. Solve for the change (x) and equilibrium concentrations.
    4. Check the results.

Thermodynamic Connection with K:

  • The value of K is related to Gibbs free energy (ΔG).
    • If ΔG < 0, the reaction is spontaneous and favors products.
    • If ΔG > 0, the reaction is non-spontaneous and favors reactants.
    • If ΔG = 0, the reaction is at equilibrium.

Le Chatelier’s Principle:

  • States that a system at equilibrium will adjust to counteract changes in concentration, pressure, or temperature.
  • Effect of concentration: Adding reactants/products shifts the equilibrium to consume the added substance and restore balance.
  • Effect of pressure: In reactions with different numbers of gaseous reactants and products, changing pressure can shift the equilibrium.

Conceptual Insights for Competitive Exams:

  1. Pure solids/liquids: Always excluded from K expressions, simplifying calculations.
  2. Reaction quotient Qc​: Key to predicting direction at non-equilibrium states.
  3. Impact of Temperature: Endothermic and exothermic shifts are commonly tested.
  4. High and Low K: Memorize reaction examples with extreme K values for inference questions.
  5. Unit Analysis: Check if K is dimensionless or has units based on stoichiometry.

Effect of Temperature on Equilibrium:

  1. Exothermic Reactions (NO2 to N2O4):
    • At low temperatures, the reaction shifts towards the formation of N2O4 (colorless), because this is an exothermic reaction.
    • At high temperatures, the reaction shifts back to NO2 (brown color), as higher temperatures favor the reverse reaction.
  2. Endothermic Reactions (Cobalt Complex):
    • In reactions like the one involving cobalt, the color of the solution changes with temperature. At higher temperatures, the solution becomes blue (due to the formation of [CoCl4]2-), while at lower temperatures, it turns pink (due to the formation of [Co(H2O)6]3+).

Effect of a Catalyst:

  1. Catalysts and Reaction Rates:
    • A catalyst speeds up both the forward and reverse reactions by providing an alternative pathway with lower energy, but it does not affect the equilibrium itself (it doesn’t change the position of equilibrium or the final concentration of reactants and products).
    • Example: The Haber process for ammonia synthesis uses iron as a catalyst to increase the reaction rate at temperatures that still allow reasonable ammonia yields.
  2. Effect of Temperature on Equilibrium:
    • In the ammonia synthesis process, low temperatures result in slow rates, but higher temperatures improve reaction rates, though they may reduce yield.
    • Optimum conditions for the Haber process are 500°C and 200 atm pressure.
  3. Catalyst in Sulfuric Acid Production:
    • In the contact process for making sulfuric acid, platinum or V2O5 (vanadium pentoxide) acts as a catalyst to speed up the conversion of SO2 to SO3.

Ionic Equilibrium:

  1. Electrolytes vs Non-Electrolytes:
    • Electrolytes are substances that dissociate into ions in water and can conduct electricity. They are either strong (fully dissociate) or weak (partially dissociate).
    • Non-electrolytes like sugar do not dissociate in water and do not conduct electricity.
  2. Example of Electrolyte (NaCl):
    • When sodium chloride (NaCl) dissolves in water, it separates into Na+ and Cl- ions, which are stabilized by water molecules.
  3. Ionic Equilibrium:
    • In weak electrolytes (e.g., acetic acid), equilibrium is established between the dissociated ions and the undissociated molecules.

Acids, Bases, and Salts:

  1. Arrhenius Concept:
    • Acids dissociate in water to produce H+ (or H3O+), and bases produce OH-.
    • In aqueous solutions, H+ is always bonded with water to form H3O+ (hydronium ion).
    • For example: HCl → H+ + Cl-, and NaOH → Na+ + OH-.
  2. Brønsted-Lowry Theory:
    • Acids donate protons (H+), and bases accept protons.
    • Example: Ammonia (NH3) acts as a base by accepting a proton, while HCl acts as an acid by donating a proton.
  3. Conjugate Acid-Base Pairs:
    • An acid and its conjugate base differ by one proton (H+).
    • Example: In the reaction between ammonia and water, ammonia (NH3) accepts a proton to become ammonium (NH4+), making NH3 and NH4+ a conjugate acid-base pair.
  4. Dual Role of Water:
    • Water can act as both an acid and a base, depending on the reaction. In HCl + H2O → H3O+ + Cl-, water is a base because it accepts a proton.

Key Knowledge for Competitive Exams:

  • Le Chatelier’s Principle: Temperature, pressure, and concentration changes shift the equilibrium in a direction that counteracts the change.
  • Effect of Temperature: For exothermic reactions, lowering temperature favors the forward reaction. For endothermic reactions, raising temperature favors the forward reaction.
  • Catalysts: They increase the rate of a reaction without affecting the equilibrium composition.
  • Conjugate Acid-Base Pairs: Important in understanding acid-base equilibria.

Lewis Acids and Bases

  1. Lewis Definition of Acids and Bases:
    • G.N. Lewis (1923) defined acids as species that accept an electron pair, and bases as species that donate an electron pair.
    • For bases, there is no major difference between Lewis and Brönsted-Lowry definitions. Both describe bases as donating lone pairs of electrons.
    • In the Lewis concept, acids do not need to contain protons. For example, BF3 does not have protons but acts as an acid by accepting a lone pair of electrons from NH3.
  2. Examples of Lewis Acids and Bases:
    • BF3 (a Lewis acid) reacts with NH3 (a Lewis base) as follows:
      • BF3 + :NH3 → BF3:NH3
    • AlCl3, Co3+, and Mg2+ are other examples of Lewis acids.
    • H2O, NH3, and OH– are examples of Lewis bases because they can donate lone pairs.
  3. Classifying Species as Lewis Acids or Bases:
    • HO–: Lewis base (donates an electron pair).
    • F–: Lewis base (donates an electron pair).
    • H+: Lewis acid (accepts an electron pair).
    • BCl3: Lewis acid (accepts an electron pair).

Understanding Acid-Base Strength

  1. Arrhenius Concept:
    • Useful in aqueous solutions where strong acids like HCl, HNO3, and H2SO4 fully dissociate into H+ ions.
    • Strong bases like NaOH and KOH dissociate completely into OH– ions.
  2. Brönsted-Lowry Concept:
    • Strong acids donate protons (H+) easily, and strong bases accept protons.
    • In the equilibrium of weak acids (e.g., HA), the equilibrium shifts towards the weaker acid.
  3. Conjugate Acids and Bases:
    • A strong acid has a very weak conjugate base.
    • A weak acid has a strong conjugate base.
    • Example: HCl (strong acid) → Cl– (weak conjugate base).

Ionization of Water and the Ionic Product (Kw)

  1. Water’s Unique Property:
    • Water can act as both an acid and a base:
      • H2O(l) + H2O(l) ⇌ H3O+(aq) + OH–(aq)
    • In pure water, [H3O+] = [OH–] = 1 × 10–7 M at 298 K.
    • The ionic product of water (Kw) is 1 × 10–14 at 298 K, which means the product of [H+] and [OH–] is always constant.
  2. Determining pH and pOH:
    • pH = –log[H+] and pOH = –log[OH–].
    • For pure water, pH = 7 (neutral), acidic solutions have pH < 7, and basic solutions have pH > 7.
  3. Relation between pH and pOH:
    • pKw = 14 at 298 K, where:
      • pH + pOH = 14

The pH Scale

  1. Definition of pH:
    • pH is a logarithmic scale used to measure the concentration of hydrogen ions [H+] in a solution.
    • pH = –log[H+].
    • For example:
      • HCl (10–2 M) has pH = 2.
      • NaOH (10–4 M) has pH = 10.
    • Pure water at 25°C has [H+] = 10–7 M, so pH = 7.
  2. Measuring pH:
    • pH paper and pH meters are used to measure the pH of a solution.
    • pH meters provide high precision, with an accuracy of up to 0.001.
  3. Common pH Values:
    • Strong acids like HCl have very low pH values (~0).
    • Acids like tomato juice have a pH around 4.2.
    • Bases like NaOH have high pH values (~13).

Ionization Constants and Weak Acids

  1. Ionization of Weak Acids:
    • A weak acid like HF does not fully dissociate. The dissociation is represented as:
      • HF ⇌ H+ + F–
    • The equilibrium constant for dissociation is called the acid dissociation constant (Ka).
  2. Calculating Degree of Dissociation (α):
    • The degree of dissociation (α) is the fraction of the acid that dissociates in solution.
    • For example, for HF in a 0.02 M solution with Ka = 3.2 × 10–4:
      • The quadratic equation is used to calculate α and determine the concentration of H+, F–, and HF.
  3. Example Calculation:
    • For HF (0.02 M) with Ka = 3.2 × 10–4, the degree of dissociation is α = 0.12, leading to:
      • [H3O+] = [F–] = 2.4 × 10–3 M and pH = 2.62.

Additional Concepts

  • pKa is the negative logarithm of Ka, and it is used to describe the strength of weak acids.
  • The pKa and Ka values help calculate the degree of ionization and the pH of weak acid solutions.

Di- and Polybasic Acids, Ionization, and Related Concepts

  1. Polybasic or Polyprotic Acids:
    • These acids have more than one ionizable proton per molecule (e.g., oxalic acid, sulphuric acid, phosphoric acid).
    • Example for a dibasic acid (H2X):
      • First ionization: H2X → H⁺ + HX⁻ (with equilibrium constant Ka₁).
      • Second ionization: HX⁻ → H⁺ + X²⁻ (with equilibrium constant Ka₂).
  2. Ionization Constants:
    • The first ionization constant (Ka₁) is for the first proton dissociation, and the second (Ka₂) is for the second proton dissociation.
    • The higher the order of ionization, the smaller the constant (Ka₂ is smaller than Ka₁).
  3. Factors Affecting Acid Strength:
    • Bond Strength: As the strength of the H-A bond decreases (energy to break the bond decreases), the acid becomes stronger.
    • Bond Polarity: If the H-A bond becomes more polar (increased electronegativity difference between H and A), acidity increases.
    • In the same Group: Larger atoms (e.g., HI > HCl > HF) have weaker H-A bonds, making them stronger acids.
    • In the same Period: As electronegativity increases (e.g., HF > H2O > NH3 > CH4), acidity increases.
  4. Common Ion Effect:
    • This refers to the shift in equilibrium when an ion already present in the system is added. For instance, adding acetate ions (Ac⁻) to acetic acid reduces the dissociation of the acid (Le Chatelier’s principle).
  5. Polyprotic Acid Solutions:
    • Polyprotic acids like H₂A (diprotic) dissociate in multiple steps (e.g., H₂A → HA⁻ + H⁺ → A²⁻ + H⁺).
    • The equilibrium concentrations depend on the ionization constants (Ka₁, Ka₂).
  6. Hydrolysis of Salts and pH:
    • When salts dissolve in water, their cations and anions can hydrolyze, affecting pH.
    • Salts of weak acids and strong bases (e.g., CH₃COONa) make the solution alkaline (pH > 7).
    • Salts of strong acids and weak bases (e.g., NH₄Cl) make the solution acidic (pH < 7).
    • Salts of weak acids and weak bases (e.g., CH₃COONH₄) can make the pH depend on the relative strengths of the acid and base involved.
  7. Buffer Solutions:
    • These resist changes in pH when small amounts of acid or base are added. Buffers are made from weak acids and their conjugate bases, or weak bases and their conjugate acids.
    • Henderson-Hasselbalch Equation:
      • For acidic buffers: pH = pKa + log([Salt]/[Acid]).
      • For basic buffers: pOH = pKb + log([Conjugate acid]/[Base]).
      • If the acid and its conjugate base (or base and its conjugate acid) are present in equal concentrations, pH (or pOH) = pKa (or pKb).
  8. Solubility Product (Ksp):
    • The solubility product constant (Ksp) is used to describe the equilibrium between a sparingly soluble ionic compound and its ions in solution.
    • Example: For barium sulfate (BaSO₄), the equilibrium is:
      • BaSO₄(s) ⇌ Ba²⁺(aq) + SO₄²⁻(aq)
      • Ksp = [Ba²⁺][SO₄²⁻] (since BaSO₄ is solid, its concentration is constant).
    • The solubility product indicates the maximum concentration of ions that can exist in equilibrium with the undissolved solid.

Key Points for Competitive Exam Preparation:

  • Focus on understanding how the acid dissociation constants (Ka) and the bond strength and polarity influence acidity.
  • Use the Henderson-Hasselbalch equation to calculate pH and pOH of buffer solutions.
  • Learn how hydrolysis of salts affects pH (acidity or alkalinity), and know how to identify the type of salt.
  • Remember the solubility product (Ksp) and its application to sparingly soluble salts.

Common Ion Effect on Solubility of Ionic Salts”

  1. Le Chatelier’s Principle:
    • If the concentration of one ion in a saturated solution increases, it will shift the equilibrium to reduce this increase by precipitating some salt, so that the product of ion concentrations (Ksp) equals the ionic product (Qsp).
    • If the concentration of an ion decreases, more salt will dissolve to increase the ion concentration until equilibrium is re-established.
  2. Effect of Adding Common Ion:
    • When a common ion is added (like chloride ion in a sodium chloride solution), the solubility of a salt decreases because the increased concentration of the common ion reduces the dissociation of the salt.
  3. Purification Using the Common Ion Effect:
    • The common ion effect can be used to purify salts like sodium chloride by precipitating other salts, like sodium or magnesium sulphates, which are less soluble under high chloride ion concentrations.
  4. Gravimetric Estimation:
    • The common ion effect is useful for precipitating ions in quantitative estimations. For example, silver ion can be precipitated as silver chloride, ferric ion as ferric hydroxide, and barium ion as barium sulphate using the common ion effect.
  5. Effect of pH on Solubility:
    • Solubility of salts like phosphates increases at lower pH because the concentration of anions (like phosphate ions) decreases due to protonation, thus shifting the equilibrium to dissolve more salt and restore the Ksp.
  6. Calculating Solubility:
    • The solubility of salts in a solution can be influenced by pH. For a salt like Ni(OH)2 in a solution containing NaOH, the solubility is affected by the existing concentration of OH– ions, and you can calculate solubility by solving equilibrium equations.
  7. Ionic Product and Solubility Product (Ksp):
    • Ksp is used to describe the equilibrium between the solid salt and its dissolved ions. If the ionic product (Qsp) exceeds Ksp, precipitation occurs. If Qsp is less than Ksp, more salt dissolves.
  8. Equilibrium Concepts:
    • Equilibrium is a dynamic state where the rate of forward and reverse reactions are equal. The equilibrium constant (Kc or Kp) expresses the ratio of product concentrations to reactant concentrations raised to their respective stoichiometric coefficients.
    • If temperature, pressure, or concentration changes, Le Chatelier’s principle can predict how the equilibrium will shift to counteract the changes.
    • A catalyst increases the rate of reaction but does not affect the equilibrium position or the equilibrium constant.
  9. Electrolytes and Ionization:
    • Electrolytes conduct electricity when dissolved in water due to the dissociation into ions. Strong electrolytes fully dissociate, while weak electrolytes establish an equilibrium between the ions and undissociated molecules.
    • Acid-Base Definitions: According to Arrhenius, acids release hydrogen ions, and bases release hydroxide ions. Brønsted-Lowry defined acids as proton donors and bases as proton acceptors, while Lewis defined acids as electron pair acceptors and bases as electron pair donors.
  10. Ionization Constants:
    • Weak acids and bases have ionization constants (Ka and Kb) that describe their degree of dissociation. These can be used to calculate pH, pKa, and pKb.
    • The ionization of water (Kw) is always 1 × 10^-14 at 25°C, and the relationship between pH, pOH, and Kw is given by: pH + pOH = 14.
  11. Hydrolysis of Salts:
    • Salts formed from strong acids and weak bases, or weak acids and strong bases, can undergo hydrolysis in water, affecting the pH of the solution.
  12. Buffer Solutions:
    • Buffers are solutions that resist changes in pH when small amounts of acid or base are added. They are crucial in maintaining stable pH in biological and chemical systems.
  13. Solubility Equilibrium:
    • The solubility product (Ksp) describes the equilibrium between a solid salt and its ions in solution. The solubility of sparingly soluble salts can be affected by the presence of common ions, which decrease solubility.

Summary of Key Concepts for Competitive Exams:

  • Le Chatelier’s Principle governs the shifts in equilibrium with changes in ion concentration.
  • The common ion effect reduces the solubility of salts by increasing the concentration of one of its ions.
  • Ksp helps predict whether a salt will precipitate or dissolve.
  • The solubility of salts can be affected by pH, and this is key for many applications like purifications and gravimetric estimations.
  • Equilibrium constants (Kc, Kp) and reaction quotient (Q) help determine the direction of chemical reactions.
  • Electrolytes, acids, bases, and buffers are important concepts for understanding how substances interact in solution, affecting pH and solubility.

These all are the notes of chapter 6 in chemistry. And after some time you get important questions HERE. *#THANKS FOR VISITING, VISIT AGAIN#* 😊

Important Questions And Answers Which Strong Your Concept Are-

  1. What is equilibrium?
    • Answer- Equilibrium is a state of balance where nothing changes, and everything is stable.
  2. How is equilibrium shown in chemical reactions?
    • Answer- It is shown with a double arrow (⇌), meaning the process is reversible and happens in both directions.
  3. What are the types of equilibrium?
    • Answer- 3
  4. What is dynamic equilibrium?
    • Dynamic equilibrium is when two opposite processes happen at the same time and at the same rate, so the overall amounts stay constant.
  5. What is chemical equilibrium?
    • Answer- Chemical equilibrium happens when the forward and reverse reactions in a chemical reaction occur at the same rate, so the amounts of reactants and products stay constant.
  6. What is ionic equilibrium?
    • Answer- Ionic equilibrium is when the rate of ionization equals the rate of recombination, keeping the ion concentration stable.
  7. What happens in solid-liquid equilibrium?
    • Answer- Ice and water coexist at 0°C because the rate of melting equals the rate of freezing.
  8. What happens in liquid-vapour equilibrium?
    • Answer- In a closed system, liquid turns to vapour, and vapour turns back to liquid at the same rate.
  9. What happens in solid-vapour equilibrium?
    • Answer- A solid changes to vapour, and vapour changes back to solid at the same rate in a closed system.
  10. What happens when a solid dissolves in a liquid?
    • Answer- The solid dissolves and re-precipitates at the same rate, creating a dynamic balance.
  11. What is a saturated solution?
    • Answer- A solution that contains the maximum amount of dissolved solid at a specific temperature.
  12. What happens when a gas dissolves in a liquid?
    • Answer- Gas dissolves in proportion to its pressure above the liquid (Henry’s Law). Increasing pressure increases gas solubility, but increasing temperature decreases solubility.
  13. What factors affect equilibrium in physical processes?
    • Answer-
      • Temperature: Changes shift equilibrium (e.g., solubility or melting/freezing).
      • Pressure: Affects gas equilibria.
      • Surface Area: Affects liquid-vapour equilibrium.
  14. What is the equilibrium constant (K)?
    • Answer- It is a ratio of the concentrations of products to reactants at equilibrium.
  15. How does changing the reaction affect K?
    • Answer-
      • Reversing the reaction: Knew= 1/K.
      • Multiplying coefficients: Knew=K n (where nnn is the multiplier).
  16. What happens if Q>K?
    • Answer- The reaction goes forward toward the products.
  17. What happens if Q=K?
    • Answer- The system is at equilibrium.
  18. What does Le Chatelier’s Principle say?
    • Answer- If a system at equilibrium is disturbed, it will shift to counteract the change.
  19. What does Henry’s Law state?
    • Answer- The amount of gas dissolved in a liquid is proportional to its pressure above the liquid.
  20. What happens at the molecular level during equilibrium?
    • Answer- The forward and reverse reactions continue, but their rates are equal, so concentrations of reactants and products stay constant.
  21. What happens to equilibrium in an exothermic reaction when the temperature is lowered?
    • Answer- The reaction shifts to produce more products (forward direction) because it releases heat.
  22. What happens to equilibrium in an endothermic reaction when the temperature is increased?
    • Answer- The reaction shifts to produce more products because it absorbs heat.
  23. Does a catalyst change the equilibrium position of a reaction?
    • Answer- No, a catalyst only speeds up the reaction rate but does not change the equilibrium composition.
  24. In the reaction of NO₂ ⇌ N₂O₄, what happens at low temperatures?
    • Answer- The equilibrium shifts to form more N₂O₄ (colorless).
  25. What happens to the color of a cobalt solution at high temperatures?
    • Answer- It turns blue due to the formation of [CoCl₄]²⁻.
  26. What is an electrolyte?
    • Answer- A substance that dissociates into ions in water and conducts electricity, e.g., NaCl.
  27. What is the difference between strong and weak electrolytes?
    • Answer- Strong electrolytes dissociate completely (e.g., NaCl), while weak electrolytes dissociate partially (e.g., acetic acid).
  28. What happens in the ionization of water?
    • Answer- Water molecules dissociate into H₃O⁺ and OH⁻ ions: H₂O + H₂O ⇌ H₃O⁺ + OH⁻.
  29. According to Arrhenius, what are acids and bases?
    • Answer- Acids produce H⁺ in water, and bases produce OH⁻.
  30. What does the Brønsted-Lowry theory say about acids and bases?
    • Answer- Acids donate protons (H⁺), and bases accept protons.
  31. What is a conjugate acid-base pair?
    • Answer: An acid and its conjugate base differ by one proton, e.g., NH₃ (base) and NH₄⁺ (acid).
  32. What is the Lewis definition of acids and bases?
    • Answer-
      • Lewis acid: Accepts an electron pair (e.g., BF₃).
      • Lewis base: Donates an electron pair (e.g., NH₃).
  33. What is pH?
    • Answer- A scale that measures the concentration of H⁺ ions, defined as: pH=−log[H+]
  34. What is the pH of pure water at 25°C?
    • Answer- 7 (neutral).
  35. What is the relationship between pH and pOH?
    • Answer- pH+pOH=14
  36. If a solution has [H⁺] = 10−4 M, what is its pH?
    • Answer- pH = 4 (acidic).
  37. What happens if you increase the concentration of reactants in a system at equilibrium?
    • Answer- The equilibrium shifts to produce more products.
  38. How does pressure affect gaseous equilibria?
    • Answer- Increasing pressure favors the side with fewer gas molecules.
  39. What happens when you increase temperature in an exothermic reaction?
    • Answer- The equilibrium shifts toward the reactants (reverse direction).
  40. Why are catalysts used in industrial reactions like the Haber process?
    • Answer- Catalysts speed up the reaction to reach equilibrium faster without changing the equilibrium composition.
  41. Which catalyst is used in sulfuric acid production?
    • Answer- Vanadium pentoxide (V₂O₅) in the contact process.
  42. What is a weak acid?
    • Answer- A weak acid only partially dissociates in water. For example, HF dissociates as: HF ⇌ H⁺ + F⁻.
  43. What is Ka ?
    • Answer- Ka​ (acid dissociation constant) measures the strength of a weak acid. A larger Ka means a stronger acid.
  44. How do you calculate the degree of dissociation (α)?
    • Answer- α is the fraction of the acid that dissociates. It is calculated using Ka​ and the initial concentration of the acid.
  45. If HF has Ka= 3.2×10−4 and the initial concentration is 0.02 M, what is the pH?
    • α=0.12 (calculated from a quadratic equation).
    • [H+]= 0.02×α= 2.4×10−3M.
    • pH=−log[H+]=2.62.
  46. What are polyprotic acids?
    • Answer- Acids that release more than one proton (H⁺) during dissociation. For example, H₂SO₄ (sulfuric acid).
  47. What are K1​ and K2​?
    • Answer-
      • K1​: The constant for the first dissociation step (stronger).
      • K2​: The constant for the second dissociation step (weaker).
  48. How does bond strength affect acidity?
    • Answer- Weaker bonds between H and the other atom make the acid stronger. For example, HI > HCl > HF.
  49. How does electronegativity affect acidity?
    • Answer- More electronegative atoms make the bond more polar, increasing acidity. For example, HF > H₂O > NH₃ > CH₄.
  50. What is the common ion effect?
    • Answer- It occurs when adding an ion that is already part of the equilibrium. This reduces the dissociation of the weak acid or base.
  51. What is a buffer?
    • Answer- A solution that resists pH changes when small amounts of acid or base are added. It is made from a weak acid and its conjugate base or a weak base and its conjugate acid.
  52. What is the Henderson-Hasselbalch equation?
    • Answer-
      • For acidic buffers:
      • pH = pKa​ + log (Salt/Acid).
  53. What is Ksp.
    • Answer- It is the equilibrium constant for the dissociation of a sparingly soluble salt into its ions.
  54. What happens when the ionic product (Qsp​) exceeds Ksp​?
    • Answer- Precipitation occurs because the solution contains more ions than it can hold in equilibrium.
  55. How does pH affect solubility?
    • Answer-
      • Lower pH (acidic) increases solubility for salts of weak acids, like phosphates, by reducing the concentration of anions.
      • Higher pH decreases solubility.
  56. How is the common ion effect used in salt purification?
    • Answer- Adding a common ion like Cl⁻ can decrease the solubility of impurities, helping to purify salts like sodium chloride.