Redox

• Oxidation : It is a process in which electron are lost by an atom, ion or molecule.
• Reduction : Reduction is a process in which electrons are gained by an atom, ion or molecule.
• Redox Reactions : Those reactions in which electrons are transferred from one substance to another are called Redox Reactions .
  

Oxidising Agents (OXIDANTS) and Reducing Agents (REDUCTANTS)

• Those species (atoms, molecules and ions) which have tendency to accept the electrons are known as oxidizing agents or oxidants, whereas those species (atoms, molecules or ions) which releases the electrons are called reducing agents or reductants. In other words, oxidizing agents are reduced and reducing agents are oxidized.

Oxidation Number or Oxidation State

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Oxidation state of an element in a particular species (atoms, molecules or ions) is the number of electrons gained or lost by that element during its change from free state into that species.


Rules for Calculation of Oxidation Number

• Following rules have been arbitrarily adopted to decide oxidation no. of elements on the basis of their periodic properties.

1. In uncombined state or free state, oxidation number of an element is zero.
2. In combined state oxidation no. of
a) F is always -1.
b) O is -2. In peroxides it is -1 in superoxides it is -1/2. However in it is +2.
c) H is +1. In ionic hydrides it is -1. (i.e., IA, IIA and IIIA metals)
d) Halogens as halide is always -1.
e) Sulphur as sulphide is always -2.
f) Metals is always +ve.
g) Alkali metals (i.e., I A group - Li, Na, K, Rb, Cs, Fr) is always +1.
h) Alkaline earth metals (i.e., II A group Be, Mg, Ca, Sr, Ba, Ra) is always +2.
3. The algebraic sum of the oxidation no. of all the atoms in a compound is equal to zero,
4. The algebraic sum of all the oxidation no. of elements in a radical is equal to the net charge on the radial.
5. Oxidation number can be zero, +ve, -ve (integer or fraction)
6. Maximum oxidation no. of an element is = Group no (Except O and F)
Minimum oxidation no. of an element is = Group no. -8 (Except metals)

• Redox reactions involve oxidation and reduction both. Oxidation means loss of electrons and reduction means gain of electrons. Thus redox reactions involve electron transfer and the number of electrons lost are same as the number of electrons gained during the reaction. This aspect of redox reaction can serve as the basis of a pattern for balancing redox reactions. There are two common and useful methods to balance redox reactions. These are
  a) Oxidation number method and
  b) Ion-electron method.
  
Balancing Redox Reaction by Oxidation Number method
• For balancing a redox reaction by oxidation number method, follow the order of steps as listed below (of course, all steps may not be required for balancing some reactions).
i) Find the oxidation numbers of the elements whose oxidation state is being changed. ii) Then, balance the number of atoms in both side of the element whose oxidation number is being charged. iii) Now, find the increase and decrease in ox. no. iv) To equalise the change in oxidation states multiply the species whose oxidation state is being changed, by a suitable integers. v) If the coefficient developed are not correct, then change them by inspection (such coefficient changes is required when an element from a compound goes in 2 different compounds, one with the same oxidation state & the other with changed oxidation state). vi) Count the charges on both sides of the equation and balance the charges in the equation by adding requisite H⁺ in acidic medium or OH- in basic medium to the required side. vii) Balance the hydrogens and oxygens by adding the appropriate number of H₂O molecules on the required side. Balancing Redox Reaction by ion-Electron Method
• This method of balancing redox reaction involves following steps.
1. Separate the reactants and products into half-reactions involving the elements that changes its oxidation number. Write the skeleton equations for each half-reaction. 2. Balance each half-reaction separately involving given steps. i) First balance the atoms of the element undergoing oxidation or reduction. ii) Then balance atoms of the elements other than hydrogen and oxygen. iii) For balancing oxygen atoms in acidic or neutral medium, add suitable number of H₂O molecules to the side deficient in O while in alkaline medium, add equal number of H₂O molecules as the excess of O on the side having excess of O atoms and add double the number of OH^- ions on the opposite side of the equation. iv) In order to balance the hydrogen atom in acidic or neutral medium, add required number of H+ to the side deficient in H while in alkaline medium, add equal number of OH^- ions as the excess number of atom on the side having excess H and add equal number of H₂O molecule on the opposite side of the equation. 3. Multiply each half-reaction by suitable integer to make the number of electrons lost and gained same and add both the half-equations to get a completely balanced reaction.
• Mostly, the medium in which a redox reaction is to be balanced is given in the problem but if the problem does not state the medium explicitly, then the medium is decided by looking at the reactants or products. If an acid or base is one of the reactants or products, then the medium is the same. If ammonia is present, the solution would be basic, for example ammonium ion is present, it would be acidic. If metals which forms insoluble hydroxides are shown in their ionic form, the solution is acidic.
Guidelines for the identification of Oxidizing and Reducing Agent 1. If an element is in its highest possible oxidation state in a compound, it can function as an oxidizing agent, e.g. KMnO₄, K₂Cr₂O₇, HNO₃, H₂SO₄, HClO₄ etc. 2. If an element is in its lowest possible oxidation state in a compound, it can function as a reducing agent, e.g. H₂S, FeSO₄, Na₂S₂O₃, SnCl₂ etc. 3. If an element is in its intermediate oxidation state in a compound, it can function both as an oxidizing agent as well as reducing agent, e.g. H₂O₂, H₂SO₃, HNO₂, SO₂ etc. 4. If highly electronegative element is in its higher oxidation state in a compound, that compound can function as a powerful oxidizing agent, e.g. KClO₄, KClO₃, KIO₃ etc. 5. If an electronegative element is in its lowest possible oxidation state in a compound or in free state, it can function as a powerful reducing agent, e.g. I^-, Br^-, N^3- etc. Common Oxidising and Reducing Agents

Oxidising agent	Effective Change	Decrease in Oxidation Number
KMn4 in acid solution	MnMn2-	5
KMnO4 in neutral solution	MnMnO2	3
K2Cr2O7 in acid solution	Cr2Cr3+	3
dilute HNO3	NO	3
concentrated HNO3	NO2	1
concentrated H2SO4	SO2	2
manganese (IV) oxide	MnO2 Mn2+	2
Chlorine	Cl2 Cl-	1
Chloric (I) acid	ClO-Cl-	2
KIO3 in dilute acid	I2	5
KIO3 in concentrated acid	I+	4

Reducing agent	Effective change	Increase in oxidation number
iron (II) salts (acid)	Fe2+ Fe3+	1
tin (II) salts (acid)	Sn2+ Sn4+	2
ethanedioates (acid)	C2CO2	1
sulphites (acid)		2
hydrogen sulphide	S2-S	2
odides (dilute acid)	I-I2	1
odides (concentrated acid)	I-I+	2
netals, e.g. Zn	Zn Zn2+	2
Hydrogen	H2 H+	1

NOTE: The complete theory about cells and electrochemistry is given in the chapter of Electro chemistry class [12^th] DESCRIBED ON THIS BLOG.