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Chapter 7 : Equilibrium

7.1 Equilibrium in Physical Processes arrow_upward

  • The Most familiar examples are Physical processes are:
    • Solid-Liquid Equilibrium
    • Liquid-Gas Equilibrium
    • Solid-Gas Equilibrium

    7.1.1 Solid-Liquid Equilibrium

  • Ice and Water kept in an insulated thermos flask ate 273K and the atmospheric pressure are in equilibrium state we observe that.
  • Both the opposing processes occur simultaneously
  • Both the processes occur at the same rate so that the amount of ice and water remains constant.

  • 7.1.2 Liquid-Vapour Equilibrium

  • Drying agent like anhydrous calcium chloride (or phosphorus penta-oxide) is placed for a few hours in the box.
  • Watch glass (or petri dish) containing water is quickly placed inside the box.
  • The rate of evaporation is constant
  • The rate of increase in pressure decreases with time due to condensation of vapour into water.

  • 7.1.3 Solid-Vapour Equilibrium

  • If we place solid iodine in a closed vessel, after sometime the vessel gets filled up with violet vapour and the intensity of colour increases with time.
  • After certain time the intensity of colour becomes constant and at this stage equilibrium is attained.
  • Solid iodine sublimes to give iodine vapour and the iodine vapour condenses to give solid iodine.

  • 7.1.4 Equilibrium Involving Dissolution of Solid or Gases in Liquids

  • For dissolution of solids in liquids, the solubility is constant at a given temperature.
  • Dissolution of gases in liquid, the concentration of a gas in liquid is proportional to the pressure of the gas over the liquid.

  • 7.1.5 General Characteristics of Equilibria Involving Physical Processes

  • Equilibrium is possible only in a closed
  • system at a given temperature.
  • All measurable properties of the system
  • remain constant.

  • 7.2 Equilibrium in Chemical Processes-Dynamic Equilibrium arrow_upward




    H2 O (I)         H2 O(g)

    pH2 0 constant at given temperature


    H2 O (s)         H2 O(I)

    Melting point is fixed at constant pressure


    Sugar(s)-----Sugar (solution)

    Concentration of solute in solution is constant at given temperature


    CO2 ----------CO2 (aq)

    [gas(aq)]/[gas(g)] is constant at a given temperature

    [CO2 (aq)]/[CO2 (g)] is constant at a given temperature

    7.3 Law of Chemical Equilibrium and Equilibrium Constant arrow_upward

  • Let us consider a general reversible reaction:
  •                    A + B 􀀀 C + D

  • Where A and B are the reactants, C and D are the products in the balanced chemical equation.
  • The concentrations in an equilibrium mixture are related by the following equilibrium equation,
  • [C] [D]

                        Kc = -----------

    [A] [B]

  • The product of concentrations of the reaction products raised to the respective stoichiometric coefficient in the balanced chemical equation.
  • Divided by the product of concentrations of the reactants raised to their individual stoichiometric coefficients has a constant value. This is known as the Equilibrium Law or Law of Chemical Equilibrium.

  • 7.4 Homogeneous Equilibria arrow_upward

  • All the reactants and products are in the same phase.
  • CH3 COOC2 H5 (aq) + H2 OCH3 COOH(aq) + C2 H5 OH(aq)

  • For reactions involving gases, however, it is usually more convenient to express the equilibrium constant in terms of partial pressure.
  • The ideal gas equation is:

    7.5 Heterogeneous Equilibria arrow_upward

  • Equilibrium in a system having more than one phase is called heterogeneous equilibrium.
  • The equilibrium between water vapour and liquid water in a closed container.
  • H2 O(I)  H2 O(g)

    7.6 Applications of Equilibrium Constants arrow_upward

  • Let us consider applications of equilibrium constant to:
    • Predict the extent of a reaction on the basis of its magnitude,
    • Predict the direction of the reaction, and
    • Calculate equilibrium concentrations.

    7.7 Relationship Between Equilibrium Constant K, Reaction quotient Q and Gibbs energy G arrow_upward

  • ΔG is negative, then the reaction is spontaneous and proceeds in the forward direction
  • ΔG is positive, and then reaction is considered non-spontaneous. Instead, as reverse reaction would have a negative ΔG, the product of the forward reaction shall be converted to the reactants.
  • ΔG is 0. Reaction has achieved equilibrium; at this point, there is no longer any free energy left to drive the reaction.
  • ΔG =ΔGΘ + RT InQ,

        where, GΘ is standard  Gibbs energy.

  • At equilibrium, when ΔG=0 and Q=Kc , the equation becomes.
  •       ΔG= ΔGΘ +RT In K=0

          ΔGΘ =-RTInK

          Ink=-ΔGΘ /RT


    7.8 Factors Affecting Equilibrium arrow_upward

  • Some common factors affecting to equilibrium are:
    • Effect of Concentration Change
    • Effect of Concentration Change
    • Effect of Inert Gas Addition
    • Effect of a Catalyst

    7.8.1 Effect of Concentration Change

  • The concentration stress of an added reactant/product is relieved by net reaction in the direction that consumes the added substance.

  • 7.8.2 Effect of Concentration Change

  • A pressure change obtained by changing the volume can affect the yield of products in case of a gaseous reaction where the total number of moles of gaseous reactants and total number of moles of gaseous products are different.

  • 7.8.3 Effect of Inert Gas Addition

  • If the volume is kept constant and an inert gas Such as argon is added which does not take part in the reaction, the equilibrium remains undisturbed.

  • 7.8.4 Effect of a Catalyst

  • A catalyst increases the rate of the chemical reaction by making available a new low energy pathway for the conversion of reactants to products. It increases the rate of forward and reverse reactions that pass through the same transition state and does not affect equilibrium.

  • 7.9 Ionic Equilibrium in Solution arrow_upward

  • It is well known that the aqueous solution of sugar does not conduct electricity.
  • Common salt is added to water it conducts electricity. Also, the conductance of electricity increases with an increase in concentration of common salt.

  • 7.10 Acids, Bases and Salts arrow_upward

  • Acids, bases and salts find widespread occurrence in nature.
  • Hydrochloric acid present in the gastric juice is secreted by the lining of our stomach in a significant amount of 1.2-1.5 L/day and is essential for digestive Processes.
  • Lemon and orange juices contain citric and ascorbic acids and tartaric acid is found in tamarind paste.
  • Dissolution of sodium chloride in water. Na+ and Cl ions are stabilized by their hydration with polar water molecules.

  • 7.10.1 Arrhenius Concept of Acids and Bases

  • According to Arrhenius theory acids are substances that dissociates in water to give hydrogen ions H+ (aq) and bases are substances that produce hydroxyl ions OH(aq).

  • 7.10.2 The Bronsted-Lowry Acids and Bases

  • According to theory “acid is a substance that is capable of donating a hydrogen ion H+ and bases are substances capable of accepting a hydrogen ion, H+ ”.
  • The example of dissolution of NH3 in H2 O represented by the following equation:

  • 7.10.3 Lewis Acids and Bases

  • G.N. Lewis in 1923 defined an acid as a species which accepts electron pair and base which donates an electron pair.
  • Typical example is reaction of electron deficient species BF3 with NH3 . BF3 does not have a proton but still acts as an acid and reacts with NH3 by accepting its lone pair of electrons.

  • 7.11 Ionization of Acids and Bases arrow_upward

  • Strong acids like perchloric acid (HClO4 ), hydrochloric acid (HCl), hydrobromic acid (HBr), hydroiodic acid (HI), nitric acid (HNO3 ) and sulphuric acid (H2 SO4 ) are termed strong.
  • Because they are almost completely dissociated into their constituent ions in an aqueous medium, thereby acting as proton (H+ ) donors.

  • 7.11.1 The Ionization Constant of Water and its Ionic Product

  • In presence of an acid, HA it accepts a proton and acts as the base while in the presence of a base. B it acts as an acid by donating a proton.
  • In pure water, one H2 O molecule donates proton and acts as an acid and another water molecules accepts a proton and acts as a base at the same time.

  • 7.11.2 The pH Scale

  • Hydronium ion concentration in molarity is more conveniently expressed on a logarithmic scale known as the pH scale.
  • The pH of a solution is defined as the negative logarithm to base 10 of the activity (H+) of hydrogen ion.
  • Acidic solutions possess a concentration of hydrogen ions, [H+ ] > 10–7 M, while basic solutions possess a concentration of hydrogen ions, [H+ ] < 10–7 M. thus, we can summaries that:
    • Acidic solution has pH<7
    • Basic solution has pH>7
    • Neutral solution has pH=7

    7.11.3 Ionization Constants of Weak Acids

  • Ka is the ionization constant of acid.
  • Some selected weak acids are:
    • Hydrofluoric acid (HF), Nitrous acid (HNO2 ) and Phenol (C6 H5 OH).

    7.11.4 Ionization of Weak Bases

  • Kb is the ionization constant of base
  • Some selected weak bases are:
    • Dimethylamine (CH3 )2 NH, Ammonia NH3 , Pyridine C6 H5 N.

    7.11.5 Relation between Ka and Kb

  • The equilibrium constant for a net reaction obtained after adding two (or more) reactions equals the product of the equilibrium constants for individual reactions:
  • KNET = K1 K2

  • pK values of the conjugated acid and base are related to each other by the equation:
  • pKa + pKb = pKw = 14(at 298K)

    7.11.6 Di- and Polybasic Acids and Di- and Polyacidic Bases

  • Some of the acids like oxalic acid, sulphuric
  • acid and phosphoric acids have more than one ionizable proton per molecule of the acid. Such acids are known as polybasic or
  •     polyprotic acids.

  •  7.11.7 Factors Affecting Acid Strength

  • An acid depends on the strength and polarity of the H-A bond.
  • Strength of H-A bond decreases, that is, the energy required to break the bond decreases, HA becomes a stronger acid.

  • 7.11.8 Common Ion Effect in the Ionization of Acids and Bases

  • An acetic acid solution results in decreasing the concentration of hydrogen ions, [H+ ].

  • 7.11.9 Hydrolysis of Salts and the pH of their Solutions

  • The cations/anions formed on ionization of salts either exist as hydrated ions in aqueous solutions or interact with water to reform corresponding acids/bases depending upon the nature of salts.

  • 7.12 Buffer Solutions arrow_upward

  • The solutions which resist change in pH on dilution or with the addition of small amounts of acid or alkali are called Buffer Solutions.
  • A mixture of acetic acid and sodium acetate acts as buffer solution around pH 4.75 and a mixture of ammonium chloride and ammonium hydroxide acts as a buffer around pH 9.25.

  • 7.13 Solubility Equilibrium of Sparingly Soluble Salts arrow_upward

  • The solubility depends on a number of factors important amongst which are the lattice enthalpy of the salt and the solvation enthalpy of the ions in a solution.

  • 7.13.1 Solubility Product Constant

  • Solid like barium sulphate in contact with its saturated aqueous solution.
  • The equilibrium between the undisolved solid and the ions in a saturated solution can be represented by the equation.
  • The equilibrium constant is given by the equation:
  • A pure solid substance the concentration remains constant and we can write:

  • 7.13.2 Common Ion Effect on Solubility of Ionic Salts

  • Addition of a soluble salt that contain one of the ions of the insoluble salt, decreases the solubility of the insoluble salt.
  • The common ion effect is also used for almost complete precipitation of a particular ion as its sparingly soluble salt.                               

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