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Introduction


A Redox reaction consists of two processes
  1. Oxidation - A process in which there is an increase in the oxidation number of the element. Occurs by removal of electrons
  2. Reduction - A process in which there is a decrease in the oxidation number of the element. Occurs by addition of electrons
    Oxidation and reduction processes are complementary in nature and do not occur singly.
Oxidizing agent: A Substance that undergoes reduction is called an oxidizing agent.
Reducing agent: A Substance that undergoes oxidation is called a reducing agent.
Oxidation number: Charge on an atom in a compound, when all the other elements are removed from it.

Rules to calculate oxidation number

  1. Oxidation number in elemental state is zero
  2. For alkali metals, the oxidation number is always taken as +1
  3. For alkaline earth metals, the oxidation number is always taken as 2, except in peroxides where the oxidation n umber is taken as +1
  4. Oxidation number of halogens is taken as -1. However, when Cl, Br or I form complexes with O, halogens take a positive oxidation number
  5. Oxidation number of O is taken as -2 except in peroxides, where oxidation number of O is taken as -1
  6. When O combines with F, the oxidation number of O becomes positive
  7. H gets a positive oxidation number in most of its compounds except in metal hydrides where its oxidation number is calculated as -1

Types of redox reactions

  1. Combination reactions: Two or more molecules combine to form a larger molecule
    A + B  C

    For example: C + O2  CO2
    On the reactant side: The oxidation number of C and O is zero as both are in the elemental state
    On the product side, the oxidation number of O is -2
    Since, there are two O atoms in CO2 the oxidation number of C in CO2 is +4
    Thus, C undergoes oxidation while O undergoes reduction in the combustion of Carbon
  2. Decomposition reactions: A larger molecule decomposes to two or more smaller molecules.
    A  B +C

    For example: 2 KClO3  2 KCl + 3 O2
    On the reactant side: the oxidation state of O is -2
    The oxidation state of K is +1
    The oxidation state of Cl is +5
    On the product side, the oxidation state of O is 0
    While the oxidation state of Cl is -1
    The oxidation state of K is +1
    Therefore, Cl has undergone reduction and O has undergone oxidation in the thermal decomposition of KClO3
    Note: At least one of the products needs to be in elementary state.
  3. Displacement reactions: X + YZ  ZX + Y
    For example: CuSO4 + Zn  ZnSO4 + Cu
    In this reaction, oxidation state of Cu changes from +2 to 0 (reduction)
    And that of Zn changes from 0 to +2(oxidation)

Balancing of redox reactions


There are two methods
  1. Oxidation number method
    1. Calculate the oxidation number of each element on either side of the equation
    2. Calculate the increase or decrease in the oxidation number per atom
    3. Multiply by suitable integer to equalize the increase or decrease in oxidation number
    4. Balance all the atoms other than O and H
    5. Finally, add O and H by adding water molecules to the side deficient in these atoms
    6. In case of Ionic reactions
      1. For acidic medium, first balance O atoms by adding water molecules to the side deficient in O atoms. Add H+ ions to the balance H atoms to the side deficient in H.
      2. For alkaline medium, first balance O atoms by adding water molecules to the side deficient in O atoms. H atoms are also balanced by adding H2O molecules to the side deficient in H and an equal number of OH- ions are added on the opposite side.
  2. Ion electron method - For any redox reaction, the number of electrons lost in oxidation and the number of electrons added in reduction must be same. This method makes use of this principle
    1. Write the oxidation number of each of the element on either side of the equation
    2. Identify the reduction half and oxidation half reactions
    3. Balance the two half reactions separately by applying the following rules
      1. Balance the atoms that have undergone a change in oxidation number
      2. Add electrons to the side to balance the difference in oxidation number
      3. Balance the charge by adding H+ for reactions occurring in acidic medium and OH- for reactions taking place in alkaline medium
      4. Balance O atoms by adding required number of H2O molecules to the side deficient in O atoms
      5. In acidic medium, H atoms are balanced by adding H+ ions to the side deficient in H. While in basic medium, H atoms are balanced by adding H2O molecules to the side deficient in H and an equal number of OH- ions are added on the opposite side
    4. The two half reactions are then added and multiplied by an integer to cancel out the electrons.

Redox reactions and Electrode Processes

An electrochemical cell has two electrodes

An electrode at which oxidation occurs - Anode

An electrode at which reduction occurs - Cathode

An electrochemical cell consists of two electrodes places in respective electrolytes connected through a salt bridge

At each electrode, a redox half reaction occurs

Electrons flow in the direction of anode to cathode

The flow of electricity is therefore in the opposite direction from cathode to anode

A redox couple is usually denoted as Mn+/M (Oxidized state/Reduced state)

For a Galvanic cell


At anode: Zn(s) -→ Zn2+(aq) + 2 e- (Oxidation)

At cathode: Cu2+(aq) + 2 e-  Cu(s) (Reduction)

Overall cell reaction: Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)

An electro chemical cell is therefore represented as

Zn2+(aq)/Zn(s)//Cu2+(aq)/Cu(s)

in general. Oxidation Half Cell (Oxidized form/Reduced form) salt bridge Reduction half cell (Oxidized form/Reduced form)

The sign // denotes salt bridge

Functions of Salt Bridge
  1. Allows movement of ions
  2. Maintains electrical neutrality


Electrode Potential: Tendency of an electrode to lose or gain electrons. Electrodes with higher reduction potential undergo reduction. Electrode potential is given a positive sign if oxidation occurs at the electrode. The reduction potentials are calculated against standard Hydrogen electrode, for which the electrode potential is taken as Zero at STP.

Electrochemical series: A series which consists of arrangement of oxidising agents in their decreasing order of strength. It is also termed as electromotive series.

Applications of Electrochemical series

  1. To compare relative strengths of oxidising and reducing agents - elements with higher reduction potential will undergo reduction 
  2. To predict whether a metal will react with H from acids - metals that are placed below H in electrochemical series will liberate H from acids 
  3. Emf of a cell: the standard emf of a cell is calculated as -
    E0Cell = E0Cathode - E0Anode

Equilibrium in Physical and Chemical Process


Equilibrium represents a state of a process in which the properties such as temperature, pressure, and concentration do not show any noticeable change over time.

Equilibrium can be a physical or a chemical process

Examples of Physical equilibrium
(to be given clearly)

Characteristics of equilibrium

  1. The state of equilibrium is a dynamic state - the reaction does not stop at the state of equilibrium. But, only that the rate of conversion in either direction remains the same and hence, there is no visible change observed, because the rate of forward reaction equals the rate of backward reaction
  2. Equilibrium can be established only if none of the products formed escape.
  3. The equilibrium can be attained from either direction
  4. A catalyst does not alter the state of equilibrium
  5. Law of chemical equilibrium - definition?
 

Law of Mass Action


The rate of a chemical reaction is proportional to the product of the active masses of the reactants with each reactant term raised to the power of its stoichiometry in the reaction.

For a chemical reaction -


Equilibrium Constant = rate of forward reaction/rate of backward reaction
From the law of mass action,
rate of forward reaction/rate of backward reaction = ([C]c[D]d)/([A]a[B]b) = Kc
for gaseous reactions, Pressures are employed in place of Concentrations
therefore, Kp = (PCc×PDd)/(PAa×PBb)
Relationship between Kp and Kc
Characteristics of Equilibrium Constant:
From gas equation, we know that
PV = nRT
Since, n/V = C, concentration;
Kp = [(CCRT)c (CDRT)d]/[(CART)a(CBRT)b]
    = [(CCcCDd)/CAaCBb)](RT)[(c+d)-(a+b)]
    = KC(RT)Δn

Characteristics of Equilibrium Constant

  1. Equilibrium constant of a reaction is independent of the concentrations of the reactants
  2. Equilibrium constant is independent of the direction in which the state of equilibrium is approached
  3. Equilibrium constant for a given reaction is dependent on temperature
  4. If the reaction is reversed, the value of equilibrium constant is inversed
  5. When a reaction is multiplied or divided by a number, n; the new equilibrium constant is taken as the Power of that number, n.
  6. When the equilibrium equation is divided by a number n, the new equilibrium constant is taken as the nth root of the original value.
  7. For equations in which equilibrium process is written in more than one step, then, the final equilibrium constant is the product of the individual equilibrium constants.

Effect of Temperature on Equilibrium constant


The effect of temperature on K is given by The Vo't Hoff equation: (d ln Kp/ d T) = ΔH0/RT2

Where, ΔH0 = Standard enthalpy change
In the integrated form: log (Kp,2/Kp,1) = (ΔH0/2.303 R)[(T2-T1)/T1T2]
Where, KP,1 and KP,2 are equilibrium constants at Temperatures T1 and T2 respectively.
Applications of Equilibrium constant:
Predicting the extent of reaction:
Larger the value of equilibrium, greater the extent of reaction in the given direction
In addition, equilibrium concentrations can also be calculated from equilibrium constant values.

Factors affecting equilibrium constant

  1. Change of concentration of any product or reactant
  2. Change of temperature
  3. Change of pressure
  4. Effect of catalyst (does nor affect the value of equilibrium constant, however; helps in attaining of equilibrium rapidly)
  5. Addition of inert gas at constant volume does not affect the reaction. However, if the reaction takes place at constant pressure, shifts equilibrium towards larger number of moles.
Le Chattlier's principle
When a system at equilibrium is disturbed, the equilibrium shifts in a direction to nullify the change.

Effect of Temperature
In an equilibrium, if the forward reaction is an exothermic reaction, then, the backward reaction would be endothermic. From our knowledge of thermodynamics, we know that exothermic reactions are favoured at lower temperatures while endothermic processes are favoured at higher temperatures. Therefore, on increasing the temperature, equilibrium shifts in a direction that favours endothermic process.

Effect of Pressure
Increase in Pressure will lead to shift in the direction with lesser number of gaseous moles.

Effect of Volume
If the volume of the gaseous reaction is decreased, then, the pressure exerted by the molecules will increase and hence, the reaction will shift in a direction with lesser number of gaseous moles.

Effect of Concentration
Increase in the concentration of reactants would lead to formation of products and hence the forward reaction is favoured, whereas a decrease in concentration of reactants will force the equilibrium to shift to the backward direction.

Equilibrium Constant and Reaction coefficient
Equilibrium constant can be considered as a special case of reaction coefficient, when the reaction has attained the state of equilibrium.

A Reaction Quotient is also given by the law of mass action. However, the concentrations of the reactants and those of products have either not yet reached the equilibrium state or have crossed over the equilibrium state.

From Thermodynamics, ΔG = ΔG0 + RT ln Q, where Q is the reactant Quotient.

At equilibrium, ΔG = 0, Q = K

Therefore, ΔG0 = -RT ln Keq




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