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Homogeneous and Heterogeneous Equilibria

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Introduction

A system in which all reactants are in same phase is known as homogeneous system.

For example in the reaction N2(g) + 3H2(g) → 2NH3(g), all the reactant and products are in gas phase so this is a homogeneous system.
 

homo_heterogenous_equilibria_02

Homogeneous Equilibria

homo_heterogenous_equilibria_03
When in a equilibrium reaction all the reactants and products are in same phase, it is known as homogeneous equilibrium.

Examples of equilibrium in gas phase are given below:

(I) The reactions in which number of moles of products are equal to number of moles of reactants

H2 + I→ 2HI

N2 + O2 → 2NO


(II) The reactions in which number of moles of products are not equal to number of moles of reactants


N2 + 3H
→ 2NH


2 SO2 + O2 → 2SO3

PCl
→ PCl3+ Cl2
 

Examples of equilibrium in liquid phase are:

CH3COOH + C2H5OH 
→ CH3COOC2H5+ H2O


Expression for equilibrium constant:


For the reaction of hydrogen and iodine to form hydrogen iodide


H2(g) + I2(g) 
→ 2HI (g)

homo_heterogenous_equilibria_04

 

For the reaction of formation of ammonia from hydrogen and oxygen

N2 + 3H
 → 2NH3

homo_heterogenous_equilibria_05

Heterogeneous Equilibria

homo_heterogenous_equilibria_06The equilibrium in which the reactants and products of a reaction are present in two or more than two phases, is called a heterogeneous equilibrium.

Some examples of heterogeneous equilibrium are:

CaCO3(s) 
 → CaO (s) + CO2 (g)


3Fe (s) + 4H2O (g)
  →  Fe3O4 (s) + 4H2 (g) 


H2O (l)
  → H2O(g)


Ag2O(s) + 2HNO3(aq)
  → 2AgNO3(aq) +H2O(l)


The position of the heterogeneous equilibrium is independent of the amount of pure solid or pure liquid present in the reaction mixture.

As the concentration of solids and liquids remains almost constant during the reaction they do not appear in the equilibrium expression.
 

Expression for equilibrium constant:

For the decomposition of calcium carbonate to calcium oxide and carbon dioxide

CaCO3(s) 
  →  CaO (s) + CO2 (g) 

homo_heterogenous_equilibria_07
But by convention [CaCO3 (s)] = 1, [CaO (s)] = 1

Hence, K = [CO2 (g)]


It is better to express the concentration of a gas in terms of partial pressure, the equilibrium constant of this reaction can be is expressed as

homo_heterogenous_equilibria_08


Above equation explains why concentration of CO2 becomes constant after the equilibrium is attained in the decomposition of calcium carbonate in a closed vessel.

(ii) For the equilibrium


H2O (l
  →  H2O(g)

homo_heterogenous_equilibria_09


But by convention [H2O(l)] = 1

Hence, Kc = [H2O(g)]


Or, in terms of pressure, homo_heterogenous_equilibria_010

This explains why vapour pressure of water is constant at constant temperature.
 

(iii) In the reaction of silver oxide with nitric acid

Ag2O(s) + 2HNO3(aq) 
  →  2AgNO3(aq) +H2O(l)


homo_heterogenous_equilibria_011as [Ag2O(s)] =1, [H2O (l)]=1


Why vapour pressure of water is is constant at constant temperature?

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Reference Links:

  1. http://www.chemguide.co.uk/physical/equilibria/kc.html
  2. http://www.brightstorm.com/science/chemistry/chemical-equilibrium/heterogeneous-equilibrium-homgeneous-equilibrium
  3. http://www.attanolearn.com/excel/5560_homogeneous-heterogeneous-equilibria.jsf
  4. http://www.freefictionbooks.org/books/p/23058-the-phase-rule-and-its-applications-by-findlay?start=8
  5. http://www.nyu.edu/classes/tuckerman/honors.chem/lectures/lecture_21/node6.html
    

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