Equilibrium : Notes and Study Materials -pdf
- Concepts of Equilibrium
- Equilibrium Master File
- Equilibrium Revision Notes
- Equilibrium MindMap
- NCERT Solution Equilibrium
- NCERT Exemplar Solution Equilibrium
- Equilibrium: Solved Example 1
- Equilibrium: Solved Example 2
- Equilibrium : Practice Paper 1
- Equilibrium : Practice Paper 2
- Equilibrium : Practice Paper 3
Subtopics included in NCERT Class 11 Chemistry Chapter 7 – “Equilibrium”
- Solid-liquid Equilibrium
- Liquid-vapour Equilibrium
- Solid – Vapour Equilibrium
- Equilibrium Involving Dissolution Of Solid Or Gases In Liquids
- General Characteristics Of Equilibria Involving Physical Processes
- Equilibrium In Chemical Processes – Dynamic Equilibrium
- Law Of Chemical Equilibrium And Equilibrium Constant
- Homogeneous Equilibria
- Equilibrium Constant In Gaseous Systems
- Heterogeneous Equilibria
- Applications Of Equilibrium Constants
- Predicting The Extent Of A Reaction
- Predicting The Direction Of The Reaction
- Calculating Equilibrium Concentrations
- Relationship Between Equilibrium Constant K, Reaction Quotient Q And Gibbs Energy G
- Factors Affecting Equilibria
- Effect Of Concentration Change
- Effect Of Pressure Change
- Effect Of Inert Gas Addition
- Effect Of Temperature Change
- Effect Of A Catalyst
- Ionic Equilibrium In Solution
- Acids, Bases And Salts
- Arrhenius Concept Of Acids And Bases
- The Bronsted-lowry Acids And Bases
- Lewis Acids And Bases
- Ionization Of Acids And Bases
- The Ionization Constant Of Water And It’s Ionic Product
- The Ph Scale
- Ionization Constants Of Weak Acids
- Ionization Of Weak Bases
- The Relation Between Ka And Kb
- Di- And Polybasic Acids And Di- And Polyacidic Bases
- Factors Affecting Acid Strength
- Common Ion Effect In The Ionization Of Acids And Bases
- Hydrolysis Of Salts And The Ph Of Their Solutions
- Buffer Solutions
- Solubility Equilibria Of Sparingly Soluble Salts
- Solubility Product Constant
- Common Ion Effect On Solubility Of Ionic Salts
Equilibrium Class 11 Notes Chemistry Chapter 7
• Chemical Equilibrium
In a chemical reaction chemical equilibrium is defined as the state at which there is no further change in concentration of reactants and products.
At equilibrium the rate of forward reaction is equal to the rate of backward reaction. Equilibrium mixture: The mixture of reactants and products in the equilibrium state is called an equilibrium mixtures.
Based on the extent to which the reactions proceed to reach the state of equilibrium, these may be classified in three groups:
(i) The reactions which proceed almost to completion and the concentrations of the reactants left are negligible.
(ii) The reactions in which most of the reactants remains unchanged, i.e. only small amounts of products are formed.
(iii) The reactions in which the concentrations of both the reactants and products are comparable when the system is in equilibrium.
• Equilibrium in Physical Processes
(i) Solid-Liquid Equilibrium: The equilibrium is represented as
Rate of melting of ice = Rate of freezing of water.
The system here is in dynamic equilibriums and following can be inferred.
(a) Both the opposing processes occur simultaneously
(b) Both the processes occur at the same rate so that the amount of ice and water – remains constant.
(ii) Liquid-Vapour Equilibrium
The equilibrium can be represented as
Rate of evaporation = Rate of condensation
When there is an equilibrium between liquid and vapours, it is called liquid-vapour equilibrium.
(iii) Solid-Vapour Equilibrium
This type of equilibrium is attained where solids sublime to vapour phase. For example, when solid iodine is placed in a closed vessel, violet vapours start appearing in the vessel whose intensity increases with time and ultimately, it becomes constant.
• Equilibrium involving Dissolution of Solid in Liquid
Solution: When a limited amount of salt or sugar or any solute dissolves in a given amount of water solution is formed.
At a given temperature state is reached when no more solute can be dissolved then the solution is called saturated solution.
The equilibrium between a solid and its solution is indicated by the saturated solution and may be represented as
Here dissolution and precipitation takes place with the same speed.
On adding a small amount of radioactive sugar to the saturated solution it will be found that the sugar present in the solution as well as in the solid state is radioactive.
• Equilibrium between a Gas and its Solution in Liquid
This type of equilibrium can be seen by the following example:
Let us consider a sealed soda water bottle in which C02 gas is dissolved under high pressure. A state of equilibrium is attained between CO2 present in the solution and vapours of the gas.
Henry’s law: The solubility of a gas in a liquid at a certain temperature is governed by Henry’s law. It states that the mass of a gas that dissolves in a given mass of a solvent at any temperature is proportional to the pressure of the gas above the surface of the solvent.
• Characteristics of Equilibria Involving Physical Processes
(i) The equilibrium can be attained only in closed systems at a given temperature.
(ii) At the equilibrium the measurable properties of the system remain constant.
(iii) The equilibrium is dynamic since both the forward and backward processes occur at same rate.
(iv) At equilibrium, the concentrations of substances become constant at constant temperature.
(v) The value of equilibrium constant represents the extent to which the process proceeds before equilibrium is achieved.
• Equilibrium in Chemical Processes
Like equilibria in physical systems it can also be achieved in chemical process involving reversible chemical reactions carried in closed container.
The dynamic nature of chemical equilibrium can be demonstrated in the synthesis of ammonia by Haber’s process. Haber started his experiment with the known amounts of N2 and H2 at high temperature and pressure. At regular intervals of time he determined the amount of ammonia present. He also found out concentration of unreacted N2 and H2.
After a certain time he found that the composition of mixture remains the same even though some of the reactants are still present. This constancy indicates the attainment of equilibrium. In general, for a reversible reaction the chemical equilibria can be shown by
After a certain time the two reactions occur at the same rate and the system reaches a state of equilibrium. This can be shown by the given figure.
• Equilibrium in Homogeneous System
When in a system involving reversible reaction, reactants and products are in the same phase, then the system is called as homogeneous system.
After some time it can be observed that an equilibrium is formed. The equilibrium can be seen by constancy in the colour of the reaction mixture.
• Law of Chemical Equilibrium
At a constant temperature, the rate of a chemical reaction is directly proportional to the product of the molar concentrations of the reactants each raised to a power equal to the corresponding stoichiometric coefficients as represented by the balanced chemical equation. Let us consider the reaction,
• Relationship between Equilibrium constant K, reaction Quotient Q and Gibbs energy G.
A mathematical expression of thermodynamic view of equilibrium can be described by tine equation.
• Factors Affecting Equilibria
Le Chatelier’s principle: If a system under equilibrium is subjected to a change in temperature, pressure or concentration, then the equilibrium shifts in such a manner as to reduce or to counteract the effect of change.
Effect of Change of Concentration: When the concentration of any of the reactants or products in a reaction at equilibrium is changed, the composition of the equilibrium changes so as to minimise the effect.
Effect of Pressure Change
If the number of moles of gaseous reactants and products are equal, there is no effect of pressure.
When the total number of moles of gaseous reactants and total number of moles of gaseous products are different.
On increasing pressure, total number of moles per unit volume increases, thus the equilibrium will shift in direction in which number of moles per unit volume will be less.
If the total number of moles of products are more than the total number of moles of reactants, low pressure will favour forward reaction.
If total number of moles of reactants are more than total number of moles of products, high pressure is favourable to forward reaction.
Effect of Inert Gas Addition
If the volume is kept constant there is no effect on equilibrium after the addition of an inert gas.
Reason: This is because the addition of an inert gas at constant volume does not change the partial pressure or the molar concentration.
The reaction quotient changes only if the added gas is involved in the reaction.
Effect of Temperature Change
When the temperature of the system is changed (increased or decreased), the equilibrium shifts in opposite direction in order to neutralize the effect of change. In exothermic reaction low temperature favours forward reaction e.g.,
but practically very low temperature slows down the reaction and thus a catalyst is used. In case of endothermic reaction, the increase in temperature will shift the equilibrium in the direction of the endothermic reaction.
Effect of a Catalyst
Catalyst has no effect on the equilibrium composition of a reaction mixture.
Reason: Since catalyst increases the speed of both the forward and backward reactions to the same extent in a reversible reaction.
• Ionic Equilibrium in Solution
Electrolytes: Substances which conduct electricity in their aqueous solution.
Strong Electrolytes: Those electrolytes which on dissolution in water are ionized almost completely are called strong electrolytes.
Weak electrolyte: Those electrolytes which on dissolution in water partially dissociated are called weak electrolyte.
Ionic Equilibrium: The equilibrium formed between ions and unionised substance is called ionic equilibrium, e.g.,
Acids: Acids are the substances which turn blue litmus paper to red and liberate dihydrogen on reacting with some metals.
Bases: Bases are the substances which turn red litmus paper blue. It is bitter in taste. Common Example: NaOH, Na2C03.
• Arrhenius Concept of Acids and Bases
Acids: According to Arrhenius theory, acids are substances that dissociates in water to give hydrogen ions H+(aq).
Bases: Bases are substances that produce OH–(aq) after dissociation in water.
• Limitations of the Arrhenius Concept
(i) According to the Arrhenius concept, an acid gives H+ ions in water but the H+ ions does not exist independently because of its very small size (~H-18 m radius) and intense electric field.
(ii) It does not account for the basicity of substances like, ammonia which does not possess a hydroxyl group.
• The Bronsted-Lowry Acids and Bases
According to Bronsted-Lowry, an acid is a substance which is capable of donating a hydrogen ion H+ and bases are substances capable of accepting a hydrogen ion H+.
In other words, acids are proton donors and bases are proton acceptors. This can be explained by the following example.
• Acid and Base as Conjugate Pairs
The acid-base pair that differs only by one proton is called a conjugate acid-base pair.
Let us consider the example of ionization of HCl in water.
Here water acts as a base because it accepts the proton.
CL is a conjugate base of HCl and HCl is the conjugate acid of base CL. Similarly, H20 is conjugate base of an acid H30+ and H30+ is a conjugate acid of base H2O.
• Lewis Acids and Bases
According to Lewis, acid is a substance which accepts electron pair and base is a substance with donates an electron pair.
Electron deficient species like AlCl3, BH3, H+ etc. can act as Lewis acids while species like H20, NH3 etc. can donate a pair of electrons, can act as Lewis bases.
• Ionization of Acids and Bases
Strength of acid or base is determined with the help of extent of ionization in aqueous solution.
pH Scale: Hydrogen-ion concentration are measured as the number of gram ions of hydrogen ions present per litre of solution. Since these concentrations are usually small, the concentration is generally expressed as the pH of the solution. pH being the logarithm of the reciprocal of the hydrogen ion concentration.
• Di and Polybasic Acids
Acids which contain more than one ionizable proton per molecule are called Dibasic acids or polybasic acids or polyprotic acids.
Common examples are oxalic acid, sulphuric acid, phosphoric acid etc.
Factors Affecting Acid Strength
When the strength of H-A bond decreases
The energy required to break the bond decreases, H-A becomes a stronger acid.
As the size of A increases down the group, H-A bond strength decreases and so the acid strength increases.
In a period, as the electronegativity of A increases, the strength of the acid increases.
• Common Ion Effect
If in a aqueous solution of a weak electrolyte, a strong electrolyte is added having an ion common with the weak electrolyte, then the dissociation of the weak electrolyte is decreased or suppressed. The effect by which the dissociation of weak electrolyte is suppresed is known as common ion effect.
• Hydrolysis of Salts and the pH of their Solutions
• Solubility Products
It is applicable to sparingly soluble salt. There is equilibrium between ions and unionised solid substance.
• Equilibrium: It can be established for both physical and chemical processes. At the state of equilibrium rate of forward and backward reactions are equal.
• Equilibrium constant: Kc is expressed as the concentration of products divided by reactants each term raised to the stoichiometric coefficients. For reactions,
• Le Chatelier’s principle: It states that the change in any factor such as temperature, pressure, concentration etc., will cause the equilibrium to shift in such a direction so as to reduce the effect of the change.
• Electrolytes: Substances that conduct electricity in aqueous solutions are called electrolytes.
• Arrhenius Concept: According to Arrhenius, acids give hydrogeneous while bases produce hydroxyl ions in their aqueous solution.
• Bronsted-Lowry concept: Bronsted-Lowry defined acid as proton donor and a base as a proton acceptor.
• Conjugate base and Conjugate acid: When a Bronsted-Lowry acid reacts with a base it produces its conjugate base and conjugate acid.
• Conjugate pair of acid and base: Conjugate pair of acid and base differs only by one proton.
• Lewis acids: Define acid as an electron pair acceptor and a base as an electron pair donor.
• pH Scale: Hydronium ion concentration in molarity is more conveniently expressed on a logarithmic scale known as the pH scale. The pH of pure water is 7.
• Buffer solution: It is the solution whose pH does not change by addition of small amount of strong acid or base.
For example: CH3COOH + CH3COONa.
• Solubility product (Ksp): For a sparingly soluble salt, it is defined as the product of molar concentration of the ions raised to the power equal to the number of times each ion occurs in the equation for solubilities.
CBSE Class 11 Chemistry Chapter-7 Important Questions
1 Marks Questions
1.Define dynamic equilibrium.
Ans. When the reactants in a closed vessel at a particular temperature react to give products, the concentrations of the reactants keep on decreasing, while those of products keep on increasing for sometime after which there is no change in the concentrations of either the reactants or products. This stage of the system is the dynamic equilibrium.
2.What is physical equilibrium? Give an example.
Ans .Physical equilibrium is an equilibrium between two different physical states of same substance e.g. H2 O(s) H2O(l)
3.What is meant by the statement ‘Equilibrium is dynamic in nature’?
Ans .At equilibrium, reaction does not stop rather it still continues, the equilibrium is dynamic in nature. It appears to stop because rate of forward reaction is equal to the rate of backward reaction.
4..How does dilution with water affect the pH of a buffer solution?
Ans .Dilution with water has no effect on the pH of any buffer. This is because pH of a buffer depends on the ratio of the salt, acid or salt base and dilution does not affect this ratio.
5.On what factor does the boiling point of the liquid depends?
Ans .Boiling point depends on the atmospheric pressure.
6..State Henry’s law.
Ans .The mass of a gas dissolved in a given mass of a solvent at any temperature is proportional to the gas above the solvent.
7.What happens to the boiling point of water at high altitude?
Ans .Boiling point of water depends on the altitude of the place. At high altitude atmosp here pressure thetore is less boiling point decreases.
8.On which factor does the concentration of solute in a saturated solution depends?
Ans . The concentration of solute in a saturated solution depends upon the temperature.
Sugar (soln.) sugar (solid).
9.What conclusion is drawn from the following –
H2O(s) H2O (l)
Ans . Melting point is fixed at constant pressure.
10.State the law of chemical equilibrium.
Ans. At a given temperature, 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.
11. Write the equilibrium constant for the following equation :
aA +bB cC + dD
Ans .The equilibrium constant for a general reaction
aA + bB cC +dD
is expressed as
Where [A], [B], [C] and [D] are the equilibrium concentrations of the reactants and products.
12.Write the chemical equation for the following chemical constant.
Ans .The chemical equation is given by
H2 (g) +I2(g) 2HI(g)
13.Write the expression for equilibrium constant Kp for the reaction
3Fe (s) + 4H2O (g)Fe3O4 (s) + 4H2(g)
14..The equilibrium constant for the reaction H2O + CO H2 + CO2
Is 0.44 at 1260k. What will be the value of the equilibrium constant for the reaction : 2H2 (g) + 2CO (g) 2CO(g) + 2H2O (g) at 1260 K
Ans .The reaction is reversed and also doubled,
15.Define reaction quotient.
Ans .The reaction quotient, Q is same as equilibrium constant Kc, except that the concentrations in
Qc are not necessarily equilibrium values.
16..If Qc > Kc, what would be the type of reaction?
Ans . If Qc > Kc, the reaction will proceed in the direction of the reactants (reverse reactions)
17.What inference you get when Qc = Kc?
Ans . If Qc = Kc, the reaction mixture is already at equilibrium.
18.State Le chatelier’s principle.
Ans . It states that a change in any of the factors that determine the equilibrium conditions of a system will cause the system to change in such a manner so as to reduce or to counteract the effect of the change.
19.Can a catalyst change the position of equilibrium in a reaction?
Ans .No, a catalyst cannot change the position of equilibrium in a chemical reaction. A catalyst, however, affects the rate of reaction.
20.What is the effect of reducing the volume on the system described below?
2C(s) + O2(g) 2CO(g)
Ans. The forward reaction is accompanied by increase in volume. Hence according to Chatelier’s principle, reducing the volume will shift the equilibrium in the forward direction.
21..What happens when temperature increases for a reaction?
Ans .The equilibrium constant for an exothermic reaction decreases as the temperature increases.
22.Can a catalyst change the position of equilibrium in a reaction?
Ans .No, a catalyst cannot change the position of equilibrium in a chemical reaction. A catalyst affects the rate of reaction.
23.If Qc < Kc, when we continuously remove the product, what would be the direction of the reaction?
Ans.Continuous removal of a product maintains Qc at a value less than Kc and reaction continues to move in the forward direction.
24.Define strong and weak electrolyte.
Ans . Those electrolytes which dissociate almost completely into ions in aqueous solutions are
Known as strong electrolytes while those which show poor dissociation into ions in aqueous
solutions are called weak electrolytes.
25.Write the conjugate acids for the following Bronsted bases : NH2, NH3 and HCOO–.
26.Which conjugate base is stronger CN– or F– ?
Ans .F – < CN – basic character.
27.What is the difference between a conjugate acid and a conjugate base?
Ans.A conjugate acid and base differ by a proton.
28.Select Lewis acid and Lewis base from the following :
Cu2+ , H2O, BF3 OH–
Ans .Lewis acids : Cu2+, BF3
Lewis bases : H2O, OH–
29.The dimethyl ammonium ion, (CH3)2 NH2+, is a weak acid and ionizes to a slight degree in water what is its conjugate base?
Ans . (CH3)2 NH
30.If pH of a solution is 7, calculate its pOH value.
Ans .pH +pOH = 14
pOH = 14 – pH
= 14 – 7
31.What happens to the pH if a few drops of acid are added to CH3COONH4 solution?
Ans . pH will almost remain constant.
32.What is the concentration of H3O+ and OH– ions in water at 298K?
Ans .[H3O+] = [OH–] = 1 x 10-7 mol-1
33.The pka of acetic acid and pkb of ammonium hydroxide are 4.76 and 4.75 respectively. Calculate the pH of ammonium acetate solution.
Ans .PH = 7+ [Pka – pkb]
= 7 +
= 7 + 0.005
34.Calculate the pH of the solution
Ans .pH value of 0.002M HBr.
HBr + H2O
pH = – (log H3O+) = – log (2×10-3)
= (3-log2) = 3-0.3010
35.Define Buffer solution.
Ans . The solutions which resist change in pH on dilution or with the addition of small amounts of acid or alkali are called Buffer solutions.
36.When is a solution called unsaturated?
Ans .When the ionic product is less than the solubility product the solution is unsaturated.
37.Give an example of acidic buffer?
Ans.CH3 COOh + CH3 COONa.
38.Calculate the solubility of Ag Cl (s) in pure water.
Ans . Let the solubility of Ag Cl in water be S mol L-1
Ag Cl (S) Ag + + Cl–
[Ag+] = S; [Cl–] =S
ksP = [Ag+] [Cl–]
2.8 x 10-10 = s x s
Or S =
39.Name a basic buffer having pH around 10.
Ans .Basic buffer
Na2B4O7 + Na OH
Borax sodium hydroxide.
2 Marks Questions
1.Mention the general characteristics of equilibria involving physical processes.
Ans .(a) For solid liquid equilibrium, there is only one temperature at 1 atm at which two phases can co-exist. If there is no exchange of heat with the surroundings, the mass of the two phases remain constant.
(b) For liquid vapors equilibrium, the vapors pressure is constant at a given temperature.
(c)For dissolution of solids in liquids, the solubility is constant at a given temperature.
(d)For dissolution of gases in liquids, the concentration of a gas in liquid is proportional to pressure of the gas over the liquid.
2.Write the expression for the equilibrium constant for the reaction :
4NH3 (g) + 502 (g) 4NO (g) + 6H2 O (g)
Ans .The equilibrium constant is given by
3.When the total number of moles of product and reactants are equal, K has no unit. Give reason.
Ans .When the total number of moles of products is equal to the total number of moles of reactants the equilibrium constant k has no unit for eg.
H2(g) + I2(g) 2HI(g)
Units of .
4.What is the unit of equilibrium for the reaction N2(g) + 3H2(g) 2NH3 (g).
5.Give the relation Kp = Kc .
Ans .Let us consider a reaction
aA +bB cC + dD
Assuming the gaseous components to behave ideally,
Pi Vi = ni RT …
Where [i] is the molar concentration of the species i
6.The value of Kc for the reaction
2A B+C is 2×10-3. At a given time, the composition of the reaction mixture is [A] = [B] =[C] = 3×10-4 M. In which direction the reaction will proceed?
Ans . For the reaction the reaction Qc is given by
As [A] = [B] = [C] = 3×10-4M
As Qc > Kc so the reaction will proceed in the reverse direction.
7.Write the equilibrium constant expression for each of the following reactions. In each case, indicate which of the reaction is homogeneous or heterogeneous.
(a)2CO(g) +O2(g) 2CO2(g)
(b)N2O5(g) NO2 (g) + NO3(g)
(c)Zn(s) + 2HCl (g) ZnCl2(s) +H2(g)
2H2O(l) 2H2O (l) + O2(g)
Ans. .(a) Kc
Homogeneous : a, b
Heterogeneous : c, d
8.The dissociation of HI is independent of pressure, while dissociation of PCl5 depends upon the pressure applied. Why?
Ans.For 2HI H2 + I2
Where x is degree of dissociation
For PCl5 PCl3 + Cl2
Where x is degree of dissociation
Since Kc for HI does not have volume terms and thus dissociation of HI is independent of pressure. On the other hard Kc for PCl5 has volume in denominator and thus an increase in pressure reduces volume. And to have kc constant, x decrease.
9.On what factors does the value of the equilibrium constant of a reaction depend?
Ans. The equilibrium constant of a reaction depends upon
(ii) Pressure, &
(iii)Stoichiometry of the reaction
10.Why the addition of inert gas does does not change the equilibrium?
Ans .It is because the addition of an inert gas at constant volume does not change the partial pressures or the molar concentrations of the substance involved in the reaction.
11.The equilibrium constant of a reaction increases with rise in temperature. Is the reaction exo – or endothermic?
Ans. The equilibrium constant increases with a rise in temperature. Therefore, the reaction is endothermic.
12.Using Le – chatelier principle, predict the effect of
(a)decreasing the temperature
(b)increasing the temperature
in each of the following equilibrium systems:
(i) N2 (g) + 3H2(g) 2NH3(g) +
(ii) N2(g) + O2(g) + 2NO(g)
Ans .(i) For an exothermic reaction increase in temperature shifts the equilibrium to the left and decrease in temperature shifts it to the left.
(ii) For an endothermic reaction increase in temperature shifts the equilibrium to the right and decrease in temperature shifts it to the right.
13.(i) In the reaction equilibrium
A + B C +D,
What will happen to the concentrations of A, B and D if concentration of C is increased.
(ii) what will happen if concentration of A is increased?
Ans .(i) For an equilibrium reaction
A + B C + D
If the concentration of a product is increased, the concentration of other components changes in such a way that the conc of C decreases and vice – versa.
If the conc of C is increased the conc of D will decrease and those of A and B will increase simultaneously so that the numerical value of Kc is the same and vice – versa. The equilibrium shifts to the left.
(ii) If the conc of A is increase, conc of B will decrease and those of C and D will increase simultaneously so that the numerical value of Kc is the same and vice – versa. The equilibrium shifts to the right
14.Give two examples of actions which can act as Lewis acids.
Ans .Ag+, H+.
15.Justify the statement that water behaves like an acid and also like a base on the basis of protonic concept
Ans .Water ionizes as H2O + H2O H3O+ +OH–
With strong acid water behaves as a base and accept the proton given by the acid e.g. HCl + H2O Cl– + H3O+
While with strong base, water behaves as an acid by liberating a proton e.g. :
H2O + NH3 NH4+ + OH–.
16.The degree of dissociation of N2O4,
N2O4(g) 2NO2(g), at temperature T and total pressure is α. Find the expression for the equilibrium constant of this reaction at this temperature and pressure?
At eq 1-α2α
If p is the total pressure then
Then KP =
A solution give the following colors with different indicators. Methyl orange – yellow, methyl red – yellow, and bromothymol blue Orange . what is the pH of the solution?
17.Show that, in aqueous solutions
pH + pOH = pkw
What is the value of pH + pOH at 250c?
Ans .(i) The colors in methyl orange indicates that pH > 4.5
(ii) Colors in methyl red indicates that pH > 6.0 and
(iii) colors in bromothymol blue indicates that pH < 6.3.
Therefore, the pH of the solution is between 6.0 to 6.3.
18.Calculate the pH of the solution
Ans . pH value of 0.002M HBr.
HBr + H2O
pH = – (log H3O+) = – log (2×10-3)
= (3-log2) = 3-0.3010
19.The concentration of H+ in a soft drink is 3.8 x 10-3 M. what is its pH?
Ans .[H+] = 3.8 x10-3 M
pH = – log [H+] = log (3.8 x 10-3)
= – log 3.8 – log 10-3
= – 0.5798 + 3 .
20..Define solubility product.
Ans .The solubility product of a salt at a given temperature is equal to the product of the concentration of its ions in the saturated solution, with each concentration term raised to the power equal to the number of ions produced on dissociation of one mole of the substance.
21.Ksp for Hg SO4 is 6.4 x 10-5. What is the solubility of the salt?
Ans .S = (ksp)1/2
= (6.4 x 10-5)1/2
= (64 x 10-6)1/2
22.Calculate the pH of a buffer solution containing 0.1 mole of acetic acid and 0.15 mole of sodium acetate. Ionisation constant for acetic acid is 1.75 × 10-5.
Ans .pH = pka + log
pH = – log 1.75 x 10-5 + log
or, pH = – log 1.75 x 10-5 + log 1.5
3 Marks Questions
1.Name the three group into which chemical equilibrium can be classified.
Ans .Chemical equilibrium can be classified into three groups
(i)The reaction that proceeds nearly to completion and only negligible concentrations of the reactants are left.
(ii)The reactions in which only small amounts of products are formed and most of the reactants remain unchanged at equilibrium stage.
The reactions in which the concentrations of the reactants and products are comparable, when the system is in equilibrium.
2.Give the generalizations concerning the composition of equilibrium mixtures.
Ans.(i) If Kc > 103, products predominates over reactants i.e; if Kc is very large, the reaction proceeds nearly to completion.
(ii) If Kc < 10-3, reactants predominates over products i.e; if Kc is very small, the reaction proceeds rarely.
(iii) If Kc is in the range of 10-3 to 103, appreciable concentration of both reactants and products are present.
1.Predict if the solutions of the following salts are neutral, acidic or basic:
NaCl, KBr, NaCN, NaOH H2SO4, NaNO2, NH4 NO3, KF
NaCl – Neutral
KBr – Neutral
NaCN – Basic
NaOH – Basic
H2SO4 – Acidic
NaNO2 – Basic
NH4NO3 – Acidic
KF – Basic
5 Marks Questions
1.Find the oxidation state of sulphur in the following compounds :
H2S, H2SO4, S2O42-, S2O82- and HSO3–.
2 + x = o
X = -2
+ 1 + x – 6 = -1
or x – 5 = -1
or x = +4
+2 + x – 8 = 0
Or x = + 6
There is peroxide linkage, thus
oxidation state of S is
2x – 8 = -2
2x = 6
X = +3