Topic 6 Equilibria

6 Equilibria

6.1 Chemical equilibria

Candidates should be able to:

(a) describe a reversible reaction and dynamic equilibrium in terms of forward and backward reactions;
(b) state mass action law from stoichiometric equation;
(c) deduce expressions for equilibrium constants in terms of concentrations, Kc, and partial pressures, Kp, for homogeneous and heterogeneous systems;
(d) calculate the values of the equilibrium constants in terms of concentrations or partial pressures from given data;
(e) calculate the quantities present at equilibrium from given data;
(f) apply the concept of dynamic chemical equilibrium to explain how the concentration of stratospheric ozone is affected by the photodissociation of NO2, O2 and O3 to form reactive oxygen radicals;
(g) state the Le Chatelier’s principle and use it to discuss the effect of catalysts, changes in concentration, pressure or temperature on a system at equilibrium in the following examples:
(i) the synthesis of hydrogen iodide,
(ii) the dissociation of dinitrogen tetroxide,
(iii) the hydrolysis of simple esters,
(iv) the Contact process,
(v) the Haber process,
(vi) the Ostwald process;

(h) explain the effect of temperature on equilibrium constant from the equation, ln K= -H/RT + C

6.2 Ionic equilibria

Candidates should be able to:

(a) use Arrhenius, BrØnsted-Lowry and Lewis theories to explain acids and bases;
(b) identify conjugate acids and bases;
(c) explain qualitatively the different properties of strong and weak electrolytes;
(d) explain and calculate the terms pH, pOH, Ka, pKa, Kb, pKb, Kw and pKw from given data;
(e) explain changes in pH during acid-base titrations;
(f) explain the choice of suitable indicators for acid-base titrations;
(g) define buffer solutions;
(h) calculate the pH of buffer solutions from given data;
(i) explain the use of buffer solutions and their importance in biological systems such as the role of H2CO3 / HCO3 in controlling pH in blood.

6.3 Solubility equilibria

Candidates should be able to:

(a) define solubility product, Ksp;
(b) calculate Ksp from given concentrations and vice versa;
(c) describe the common ion effect, including buffer solutions;
(d) predict the possibility of precipitation from solutions of known concentrations;
(e) apply the concept of solubility equilibria to describe industrial procedure for water softening.

6.4 Phase equilibria

Candidates should be able to:

(a) state and apply Raoult’s law for two miscible liquids;
(b) interpret the boiling point-composition curves for mixtures of two miscible liquids in terms of ideal  behaviour or positive or negative deviations from Raoult’s law;
(c) explain the principles involved in fractional distillation of ideal and non ideal liquid mixtures;
(d) explain the term azeotropic mixture;
(e) explain the limitations on the separation of two components forming an azeotropic mixture;
(f) explain qualitatively the advantages and disadvantages of fractional distillation under reduced pressure.

6 Equilibria

 

6.1 Chemical equilibria

Candidates should be able to:

(a) describe a reversible reaction and dynamic equilibrium in terms of forward and backward reactions;

 

(b) state mass action law from stoichiometric equation;

(c) deduce expressions for equilibrium constants in terms of concentrations, Kc, and partial pressures, Kp, for homogeneous and heterogeneous systems;

(d) calculate the values of the equilibrium constants in terms of concentrations or partial pressures from given data;

(e) calculate the quantities present at equilibrium from given data;

(f) apply the concept of dynamic chemical equilibrium to explain how the concentration of stratospheric ozone is affected by the photodissociation of NO2, O2 and O3 to form reactive oxygen radicals;

(g) state the Le Chatelier’s principle and use it to discuss the effect of catalysts, changes in concentration, pressure or temperature on a system at equilibrium in the following examples:

(i) the synthesis of hydrogen iodide,

(ii) the dissociation of dinitrogen tetroxide,

(iii) the hydrolysis of simple esters,

(iv) the Contact process,

(v) the Haber process,

(vi) the Ostwald process;

(h) explain the effect of temperature on equilibrium constant from the equation, ln K= -H/RT + C

6.2 Ionic equilibria

Candidates should be able to:

(a) use Arrhenius, BrØnsted-Lowry and Lewis theories to explain acids and bases;
(b) identify conjugate acids and bases;
(c) explain qualitatively the different properties of strong and weak electrolytes;
(d) explain and calculate the terms pH, pOH, Ka, pKa, Kb, pKb, Kw and pKw from given data;
(e) explain changes in pH during acid-base titrations;
(f) explain the choice of suitable indicators for acid-base titrations;
(g) define buffer solutions;
(h) calculate the pH of buffer solutions from given data;
(i) explain the use of buffer solutions and their importance in biological systems such as the role of H2CO3 / HCO3 in controlling pH in blood.

6.3 Solubility equilibria

Candidates should be able to:

(a) define solubility product, Ksp;
(b) calculate Ksp from given concentrations and vice versa;
(c) describe the common ion effect, including buffer solutions;
(d) predict the possibility of precipitation from solutions of known concentrations;
(e) apply the concept of solubility equilibria to describe industrial procedure for water softening.

6.4 Phase equilibria

Candidates should be able to:

(a) state and apply Raoult’s law for two miscible liquids;

(b) interpret the boiling point-composition curves for mixtures of two miscible liquids in terms of ideal  behaviour or positive or negative deviations from Raoult’s law;
(c) explain the principles involved in fractional distillation of ideal and non ideal liquid mixtures;

(d) explain the term azeotropic mixture;

(e) explain the limitations on the separation of two components forming an azeotropic mixture;

(f) explain qualitatively the advantages and disadvantages of fractional distillation under reduced pressure.

 

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