Physical processes: Equilibrium
* Equilibrium is a condition in which two opposing tendencies balance one another.
* The opposing tendencies of a reversible change continue to take place with equal rates simultaneously at the equilibrium stage also. Hence, the equilibrium is called dynamic equilibrium.
* Equilibrium can be established for physical as well as chemical processes.
* At the equilibrium state, if there is a change in the physical state, the equilibrium is called physical equilibrium.
* For any pure substance at atmospheric pressure, the temperature at which the solid and liquid are at equilibrium is called the normal melting point or normal freezing point of the substance.
* The pressure exerted by the vapour over a liquid when it is in equilibrium with liquid at a given temperature, is termed as vapour pressure of the liquid at that temperature.
* For any pure and liquid at one atmospheric pressure the temperature at which the liquid and vapours are at equilibrium is called the normal boiling point of the liquid.
* The process of direct phase change from solid to vapour state without the intermediate liquid state is called sublimation.
Solids or Gases in Liquids and General Characteristics of Physical Equilibria
* At a given temperature, a solution that remains in contact with excess solute is said to be a saturated solution.
* At a given temperature, solubility remains constant for the dissolution of solids in liquids.
* The phenomenon of fizz for carbonated drinks is due to the difference in the solubility of carbon dioxide at different pressures.
* Henry’s law states that the mass of a gas dissolved in a given amount of liquid at any temperature is directly proportional to the pressure of the gas present in equilibrium with the liquid.
* At a given temperature, equilibrium can be achieved only in a closed system.
* At equilibrium, the two opposing tendencies occur at the same rate, and hence, a dynamic yet stable status is established.
* At a given temperature, the attainment of equilibrium for a physical process is characterised by constant value of one of its parameters such as vapour pressure, melting point, solubility or
* concentration.
* At equilibrium, all the measurable properties of the system remain constant.
Equilibrium in Chemical Processes
* The reactions whose products cannot be converted back into the reactants under any conditions are known as irreversible reactions.
* The reactions whose products, under suitable conditions, can be converted back into the reactants are known as reversible reaction.
* Reversible reactions are represented by a pair of half headed arrows (⇌).
* In reversible reactions, the reaction in which the reactants react to give the products is called the forward reaction, and the reaction in which the products react back to give the reactants is called the backward reaction.
* When a reversible reaction is carried out in a closed vessel, a state is reached when the rate of the forward reaction becomes equal to the rate of backward reaction. This stage is known as chemical equilibrium.
* Equilibrium state can be approached from both the sides.
* The equilibrium is a dynamic equilibrium as both the forward and backward reactions continue to take place with the same speed.
* At equilibrium state, properties such as pressure, concentration, density or colour remain unchanged over time.
* Equilibrium is possible only in a closed system.
Law of Chemical Equilibrium
* The law of chemical equilibrium states that the product of concentrations of the products raised to their respective stoichiometric coefficients in the balanced chemical equation divided by the product of the concentrations of the reactants raised to their respective stoichiometric coefficients has a constant value at a given temperature.
* The equilibrium constant for the backward reaction is the inverse of the equilibrium constant for the forward reaction.
* The equilibrium in a system with all the reactants and products in the same phase is known as homogeneous equilibrium.
* The equilibrium in a system which has more than one phase is called heterogeneous equilibrium.
* For a reversible reaction involving gases, the equilibrium constant is conveniently expressed in terms of partial pressures of the reactants and products instead of their molar concentrations.
Kc and Kp are related to each other by the equation Kp = Kc (RT)∆n.
When ∆n = 0, Kp = Kc.
When ∆n is negative, Kp < Kc.
When ∆n is positive, Kp > Kc.
Equilibrium Constant: Characteristics
* The equilibrium in a system having two or more phases is known as heterogeneous equilibrium.
* The value of the equilibrium constant for a reaction is independent of the initial concentrations of the reactants and products.
* The equilibrium constant for a particular reaction at a particular temperature is always constant and changes only with change in the temperature.
* The equilibrium constant of a forward reaction and that of its backward reaction are reciprocals of each other.
* If a chemical equation is multiplied by a certain factor, then its equilibrium constant must be raised to a power equal to that factor in order to obtain the equilibrium constant for the new reaction.
Equilibrium Constant: Applications
* The knowledge of equilibrium constant helps us in:
• Predicting the extent of a reaction.
• Predicting the direction of a reason.
• Calculating the equilibrium concentrations.
If Kp > 103, the reaction proceeds to completion.
If Kp < 10-3, the reaction proceeds rarely.
If Kp = 10-3 to 103, appreciable quantities of both reactants and products are present.
* At any point in a reversible reaction the ratio of concentration of products to the reactants is known as the reaction quotient Q (Qc
with molar concentrations and Qp with partial pressures).
* The reaction quotient is calculated using the concentrations or pressure at that given time of the reaction and not the equilibrium concentrations or pressures.
If Qc > Kc, backward reaction
If Qc < Kc, forward reaction
If Qc= Kc, at equilibrium
K, Q, G Relationship and Le Chatelier’s Principle
* The change in the Gibbs free energy of a chemical reaction can be defined as the difference in the Gibbs free energy of the products and the Gibbs free energy of the reactants.
∆G = 0 indicates the system at equilibrium.
∆G > 0 indicates a non-spontaneous process.
∆G < 0 indicates a spontaneous process.
* Reaction spontaneity can be interpreted in terms of value of ∆Go using the equation
* K_{c} = e \frac{ - \triangle G° }{RT}
* If ∆G°< 0, then \frac{ - \triangle G° }{RT}becomes positive and e \frac{ - \triangle G° }{RT} > 1.
* This makes Kc > 1, which indicates a spontaneous reaction.
If ∆G°> 0, then ∆ \frac{ - \triangle G° }{RT} becomes negative and e \frac{ - \triangle G° }{RT} < 1.
* This makes Kc < 1, which indicates a non-spontaneous reaction.
* Le Chatelier’s principle may be stated as “If a system at equilibrium is subjected to a change in the temperature, pressure or concentration of the reactants or the products that govern the equilibrium, then the equilibrium position shifts in the direction in which this change is reduced or nullified”.
Effect of Change in Concentration and Pressure on Equilibrium
* When a system at equilibrium is disturbed by the addition or removal of any reactant or product, Le Chatelier’s principle predicts that the position of equilibrium shifts in the direction in which this change is reduced or nullified.
* Pressure has no effect on reactions in which there is no change in the number of moles of the reactants and products.
* Removing or decreasing the concentration of any one of the reactants always favours the backward reaction.
* Adding a product to the equilibrium mixture results in the backward reaction.
* Removing a product from the equilibrium mixture results in the forward reaction.
* Pressure has no effect on reactions in which there is no change in the number of moles of the reactants and products.
Effect of Change in Temperature and Presence of Catalyst on Equilibrium
* The value of equilibrium constant depends on temperature.
* Equilibria in which the forward reaction is exothermic have equilibrium constants that decrease with an increase in temperature.
Equilibria in which the forward reaction is endothermic have equilibrium constants that increase with an increase in temperature.
* A catalyst has no effect either on the position of equilibrium or on the equilibrium composition of a reaction mixture.
* A catalyst increases the rate of reaction by providing a new low-energy pathway for the conversion of reactants into products.
* A catalyst lowers the activation energy for both forward and reverse reactions by exactly the same amount, and thus, makes it easy for the reaction to go faster in either direction.
Acids and Bases: Theories
* The equilibrium involving the ions in an aqueous solution of weak electrolytes is called ionic equilibrium.
* The equilibrium between a dissolved undissociated molecule and its ions is known as dissociation equilibrium.
* Properties of acids:
• They are sour in taste.
• They turn blue litmus red.
• They react with active metals like zinc and liberate hydrogen gas.
* Properties of bases:
• They are bitter in taste.
• They turn red litmus blue.
• They are soapy to touch.
* According to the Arrhenius theory:
• An acid is a substance that contains hydrogen and ionises in an aqueous solution to give hydrogen ions.
• A base is a substance that contain hydroxyl group and ionises in an aqueous solution to give hydroxide ions.
* According to the Bronsted-Lowry theory, proton donors are acids and proton acceptors are bases.
* A related pair of an acid and a base which differ by a single proton is called a conjugate acidbase pair.
* According to the Lewis theory, an acid is any molecule or ion that can accept an electron pair to form a coordinate covalent bond with the donor.
Acids and Bases: Ionisation and pH
* According to the Arrhenius theory, the strength of an acid and a base is related to the ease with which they function as a source of plus ions and minus ions in an aqueous medium.
* In terms of the Bronsted-Lowry theory, a strong acid readily donates the proton and a strong base readily accepts a proton.
* Strong acids have weak conjugate bases and weak acids have strong conjugate bases.
* The ionic product of water, Kw, at given temperature, is defined as the product of the concentrations of H plus and OH minus ions in water or in aqueous solutions.
* Aqueous solutions can be categorised as neutral solutions, acidic solutions and basic solutions, depending upon the relative number of H plus and OH minus ions.
* The pH of a solution is defined as negative logarithm to the base 10… of the molar hydrogen concentration in it.
* Based on the value of the ionic product of water, the pH scale has values from 1 to 14. A solution with pH value seven is neutral, less than seven is acidic, and more than seven is basic.
Acids and Bases: Dissociation Constants
* The dissociation constant is the equilibrium constant that is affected only by a change in temperature and not by a change in concentration.
* The stronger the acid, the greater is the extent of dissociation, … the greater are the concentration of H3o plus and A minus ions, … and the larger is the value of ka.
* The stronger the weak acid, the higher is the value of ka and the lower is the value of pKa.
The stronger the weak base, the higher is the value of kb and the lower is the value of pKb.
ka= \frac{ ca² }{( 1 - a)}
pKa= -log Ka
kb = \frac{ ca² }{( 1 - a)}
pKb = – log Kb
Relation between Ka and Kb, Common Ion Effect
* The equilibrium constant for a net reaction obtained after adding two or more reactions equals the product of the equilibrium constants for individual reactions.
pKa+ pKb = pKw = 14
* Acids that can donate more than one proton to a solution are known as polyprotic acids or polybasic acids.
* Sulphuric acid and oxalic acid, which have two ionisable protons per molecule that can be donated, are called dibasic acids.
* Phosphoric acid is called a triprotic acid as it can donate three protons.
* The extent of dissociation of an acid depends on the strength and polarity of the H-A bond.
* In general, the weaker the H-A bond, the less is the energy required to break the bond and the stronger is the acid.
* In a group, acidic strength increases with an increase in the size of the atom directly attached to the hydrogen atom.
* The common ion effect may be defined as the shift in the position of an ionic equilibrium caused by the addition of a solute that provides an ion that is a part of the equilibrium.
Hydrolysis of Salts and Buffer Solutions
* Salts are formed by the neutralisation of an acid with a base, or vice versa.
* The process in which the cation or the anion of a salt reacts with water to produce an acidic or alkaline solution is called salt hydrolysis.
* Hydrolysis is the reverse process of neutralisation.
* The aqueous solution of a salt of a weak acid and a strong base is alkaline due to the hydrolysis of the anion.
* The aqueous solution of a salt of a weak base and a strong acid is acidic due to the hydrolysis of its cation.
* The aqueous solution of a salt of a weak acid and a weak base is almost neutral due to the hydrolysis of the anion and the cation.
* The aqueous solution of a salt of a strong acid and a strong base is neutral, as neither the anion nor the cation undergoes hydrolysis.
* A solution that resists any change in its pH on dilution or on the addition of a small amount of a strong acid or a strong alkali is called a buffer solution.
Solubility Equilibria of Sparingly Soluble Salts
* The energy required to overcome the attractive forces between the ions in a solution is called the lattice enthalpy.
* The energy released during the process of solvation is known as solvation enthalpy.
* A salt is soluble in a solvent only when its solvation enthalpy is higher than its lattice enthalpy.
* Depending on their solubility, salts are classified as soluble, slightly soluble and sparingly soluble.
* The solubility product of a sparingly soluble salt at a given temperature may be defined as the product of the concentrations of its ions in a saturated solution, with the concentration terms of each ion raised to the power equal to the number of times that the ion occurs in the equation representing the dissociation of the salt.
Application of Common Ion Effect
* The solubility of an electrolyte in water decreases on addition of an electrolyte that has one ion in common with the electrolyte. This is known as the common ion effect.
* The common ion effect is also used for the complete precipitation of one of the ions present in a solution as its sparingly soluble salt.
* The solubility of salts of weak acids increases with an increase in the concentration of H plus ions or a decrease in the PH.
* The common ion effect is one of the applications of the Le Chatelier’s principle.