Redox Reactions

Redox Reactions

Oxidation Number
The net actual or imaginary charge on any atom in its any state is called its oxidation number.
 

Note:

  1. In covalent compounds, partial charges are taken as complete charge - 2



  2. If more than one atom of an elements are present in a molecule, their average oxidation number is reported, which may be fractional and called oxidation state.

    But the terms 'oxidation number' and oxidation-state are moreover used interchangeably.

    Example: In Fe3O4. Oxidation number of individual iron atoms are +2 and +3, since it is actually (FeO.Fe2O3), and average oxidation number of iron in Fe3O4 is + 8/3.

Valency-factor in Redox Reactions

  1. For all compounds involved in balanced redox reaction: Transfer of electrons for formula unit of compound in balanced equation. (Except intramolecular reaction)



  2. For oxidizing or reducing agents: Change in oxidation number for formula unit of oxidizing or reducing agent.
     

    Note: Only product is required, balanced equation is not required.

Chemical and Ionic Equilibrium
Rate of reaction (-dx/dt): Rate of change in concentration of any of the reactant or product per mole in a reaction.

Rate constant or specific rate of reaction (k): Rate of reaction at unit concentration of all the reactants is called rate constant or specific rate of reaction.

Equilibrium Constant: (K) Ratio of rate constant for forward and backward reactions in a reversible reaction is called equilibrium constant.

Equation aA  +  bBcC  +  dD

Here - kf and kb are rate constant fro forward and backward reactions.

Equilibrium Constant

Equilibrium State: Condition in a reversible reaction at which all the macroscopic properties of reaction like concentration, density, color etc. becomes constant but both forward and backward reactions takes place with equal rate is called equilibrium state.

Note that, it is condition but not a numerical value.

Factors affecting above properties
 

  Rate of reaction Rate Constant Equilibrium Constant Equilibrium Sate
1. Temperature
On increasing the Temperature

Always increases
Temperature quotient

Always increases
Arrhenius equation
K = Ae-Ea/RT

Increases if Hg = +ve
- Van't Hoff equation
K = qe-H/RT

Endothermic shift takes place.
- Le Chateler principle
2. Change in concentration
On increasing concentration of reactant

Changes as per order of reaction
x x
Forward shift takes place
3. Addition of catalyzed
Arrhenius equation
x x
4. Introducing noble gas

At constant Pressure

At constant volume

Decreases

x

x



x

x



x

If Hg > 0 Then forward shift

x

Ionic Equilibrium

  1. Ionic product of water  Kw  =  [H+] [OH-]

  2. Dissociation of constant of water

  3. Degree of dissociation of water

  4. pKw  =  pH  +  pOH at 25C
     
    (i) For neutral solutions -   [H+] = [OH-] = 10-7
    Kw = 10-14
    ph = pOH = 7
    pKw = 14
    (ii) For acidic solution -
    and
    [H+] > [OH-] but Kw = 10-14
    pH < pOH but pKw = 14
    (iii) For basic solution -
    and
    [H+] < [OH-] but Kw = 10-14
    pH > pOH but pKw = 14
  5. Calculation of pH and pOH
    pH = - log [H]  and  pOH = - log [OH-]
     
    1. Strong acids - (H+) = Molarity x V.F. = Normality

    2. Strong bases - (OH-) = Molarity x V.F. = Normality

    3. Weak acids - (H+) = C =  

    4. Weak bases - (OH-) = C =   

       = Degree of dissociation
      C = Molarity
      Ka and Kb = Dissociation constant of weak acid and weak base

    5. Acidic buffer = pH = pKa + log 

    6. Basic buffer = pOH = pKb + log 

    7. Salt hydrolysis
       
      1. Strong acid-weak base Ex. NH4Cl
        pH = 7 - pKb - log C

      2. Weak acid-strong base Ex. NaCN
        pH = 7 + pKa + log C

      3. Weak acid-weak base
        pH = 7 + pKa - log Kb

      4. Strong acid-strong base-No-hydrolysis
        ph = 7

Thermodynamics
G = G + 2.303 RT log Q

Q = Reaction Quotient

At equilibrium G = 0 and Q = Kc

G = - nFEcell

and G = -nFEcell

At equilibrium G and Ecell 0

For concentration cell

Ecell = 0 & Ecell 0

For spontaneous process

Stotal > 0, G < 0 and Ecell > 0

For non-spontaneous processes

Stotal < 0, G > 0 & Ecell < 0

For reversible reactions at equilibrium

Stotal = G = Ecell = 0

Sates of Matter

van der Waal's equation

For ideal gases PV = nRT

Critical constants

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