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Oxidation Number and Oxidation State - (Concept)

Oxidation
It is a process that involves the loss of electrons by the atoms or ions.
Reduction
It is a process that involves the gain of electrons by the atoms or ions.

Any reaction, in which the electrons are exchanged between atoms or ions, represents a simultaneous process of oxidation and reduction and is called as a Redox Reaction.

In a Redox Reaction, the species that loses electron (i.e., gets oxidised) is known as reducing agent or reductant, (since it causes reduction of other species), the species which accepts electrons from reductant (i.e., gets reduced) is known as oxidising agent or oxidant (as it causes oxidation of other species).

Oxidation State (O.S.): It refers to the hypothetical charge on atoms in a compound if all the bonds were assumed to be 100% ionic.

Oxidation state, many times, is also referred to as Oxidation Number.
This means oxidation number of an element in a compound is equal to the oxidation state of that element multiplied by total atoms of that element in a particular compound.
(i) In ionic compounds, it is simply the charge on corresponding cation and anion which is expressed as the oxidation state of that particular element. For example, the oxidation state of potassium and chlorine in potassium chloride (KCl) is simply +1 and –1 respectively as KCl is treated as K⁺Cl^–.

Refer to the following examples where oxidation states are written above the atoms:

+2-1	+2-2	+3-1	+1+6-2
MgCl₂	CaS	AlCl₃	K₂SO₄

NOTE: (a) In MgCl₂ and AlCl₃, -1 is the oxidation state of Cl.
(b) In each of the cases, the sum of oxidation number of all atoms of all kinds is equal to zero since the compound is neutral.

(ii) In Covalent Compounds, it is not so easy to assign oxidation state of an atom. In order to simplify the concept, we are going to define a set of rules which would enable us to assign oxidation state to every element in any compound.

Rules for Assigning Oxidation State (O.S.) and Oxidation Number (O.N.):

Any element in free state is assigned an oxidation state of zero. For example: O.S. of H, P, S, O in H₂, P₄, S₈, O₂ respectively is zero.
The oxidation state of any cation or anion (of form A⁺ or B^-) is equal to the magnitude of its charge. For example: O.S of Ca in Ca²⁺= +2 and O.S of Al in Al³⁺= +3.
The algebraic sum of the oxidation number of all atoms in a neutral compound is equal to 0. The algebraic sum of the oxidation numbers of all atoms in an ion (like PO₄^3-) is equal to charge on the ion.
The oxidation states of Alkali Metals (Group IA) is +1 in all of their compounds and that of Alkaline Earth elements (Group IIA) is +2 in all of their compounds.
Hydrogen in almost all of its compounds is assigned an oxidation state of +1. The exception occurs when hydrogen forms compounds with strong metals like KH, NaH, MgH₂, CaH₂, etc. In all of these, the oxidation state of hydrogen is -1.
Oxygen in almost all of its compounds is assigned an oxidation state of -2. But in certain compounds like Peroxides(H₂O₂), the oxidation state of oxygen is -1. Another exception is OF₂, where O.S. is +2. O₂F₂, where O.S. is +1 and KO₂ in which O.S. is -1/2.
Fluorine is most electronegative element and is assigned an O.S. of -1, in all its compounds. For other halogens, O.S. is generally -1 except when they are bonded to a more electronegative halogen or oxygen. O.S. of iodine in IF₇ is +7, O.S. of chlorine in KClO₃ is +5.
Generally, an element with greater electronegativity is assigned -1 by hypothetical breaking of one covalent bond.

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Exam	Chapter
JEE MAIN	Redox Reaction and Electrochemistry

A.	$\mathrm{Cl_2}$
B.	$\mathrm{Br^-}$
C.	$\mathrm{Cl^-}$
D.	$\mathrm{Br_2}$

A.	Oxidation
B.	reduction
C.	both
D.	neither

A.	are reduced themselves
B.	show a decrease in oxidation number
C.	show electromotion
D.	All of these

A.
B.	$\mathrm{K _{4}\left[ V ( CN )_{6}\right]}$
C.	$\mathrm{VSO _{4}}$
D.	$\mathrm{VOSO _{4}}$

A.	are oxidised themselves
B.	show de-electronation
C.	show an increase in oxidisation number
D.	All of these

A.	$\mathrm{NaCl + KNO _{3} \rightarrow NaNO _{3}+ KCl}$
B.	$\mathrm{CaC _{2} O _{4}+2 HCl \rightarrow CaCl _{2}+ H _{2} C _{2} O _{4}}$
C.	$\mathrm{Mg ( OH )_{2}+2 NH _{4} Cl \rightarrow MgCl _{2}+2 NH _{4} OH}$
D.	$\mathrm{Z n+2 A g C N \rightarrow 2 A g+Z n(C N)_{2}}$

A.	$\mathrm{2 HI + H _{2} SO _{4} \rightarrow I _{2}+ SO _{2}+2 H _{2} O}$
B.	$\mathrm{Ca ( OH )_{2}+ H _{2} SO _{4} \rightarrow CaSO _{4}+2 H _{2} O}$
C.	$\mathrm{NaCl + H _{2} SO _{4} \rightarrow NaHSO +4 HCl}$
D.	$\mathrm{2 PCl _{5}+ H _{2} SO _{4} \rightarrow 2 POCl _{3}+2 HCl + SO _{2} Cl _{2}}$

A.	$\mathrm{NH _{3}, N _{3} H, NH _{2} OH , N _{2} O}$
B.	$\mathrm{N _{2} O , N _{3} H , NH _{2} OH , NH _{3}}$
C.	$\mathrm{N _{2} O , NH _{2} OH , N _{3} H , NH _{3}}$
D.	$\mathrm{NH _{3}, NH _{2} OH , N _{3} H , N _{2} O}$

A.	$\mathrm{F e, Fe _{3} O _{4}, Fe _{2} O _{3}}$
B.	$\mathrm{Fe _{3} O _{4}, F e, Fe _{2} O _{3}}$
C.	$\mathrm{Fe _{2} O _{3}, Fe _{3} O _{4}, Fe}$
D.	$\mathrm{Fe , Fe _{2} O _{3}, Fe _{3} O _{4}}$

A.	$\mathrm{NO _{2}< NO < N _{2} O _{3}< N _{2} O}$
B.	$\mathrm{NO _{2}< N _{2} O _{3}< NO < N _{2} O}$
C.	$\mathrm{N _{2} O < N _{2} O _{3}< NO < NO _{2}}$
D.	$\mathrm{N _{2} O < NO < N _{2} O _{3}< NO _{2}}$

A.	O.N.(Ca)>O.N.(Li)
B.	O.N.(Li)=O.N.(Ca)
C.	O.N.(Li)>O.N.(Ca)
D.	None

A.	$+1,\: +1\: and\: +1$
B.	$+2,\: +1\: and\: +\frac{1}{2}$
C.	$+1,\: +2\: and\: +4$
D.	$+1,\: +4\: and\: +2$

A.	$\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-} \rightarrow 2 \mathrm{Cr}^{3+}$
B.	$\mathrm{MnO}_{4}^{-} \rightarrow \mathrm{Mn}^{2+}$
C.	$\mathrm{CrO}_{4}^{2-} \rightarrow \mathrm{Cr}^{3+}$
D.	$\mathrm{C}_{2} \mathrm{O}_{4}^{2-} \rightarrow 2 \mathrm{CO}_{2}$

A.	$\mathrm{+ 6 \: only}$
B.	$\mathrm{+5 \: only}$
C.	$\mathrm{0 \: and +5\: only}$
D.	$\mathrm{+3 \: and\: +5\: only}$

A.	+4,+5
B.	+5,+4
C.	+3,+4
D.	+3,+3

A.	-3,+2,+10
B.	$-2,\frac{+200}{94},+6$
C.	-2,+2,+6
D.	-2,+2,+10

A.	2
B.	1
C.	0.5
D.	3

A.	$\mathrm{Ca < Zn < Mg < Ni}$
B.	$\mathrm{Ni < Zn < Mg < Ca}$
C.	$\mathrm{Zn < Mg < Ni < Ca}$
D.	$\mathrm{Ca < Mg < Zn < Ni}$

A.	$\mathrm{Cl ^{-}, S ^{2-}, K ^{+}, B a^{2+}, C r^{6+}, M n^{7+}}$
B.	$\mathrm{S^{2-}, Cl ^{-}, K ^{+}, B a^{2+}, C r^{6+}, M n^{7+}}$
C.	$\mathrm{M n^{7+}, C r^{6+}, B a^{2+}, K^{+}, C l^{-}, S^{2-}}$
D.	$\mathrm{S ^{2-}, Cl ^{-}, B a^{2+}, C r^{6+}, M n^{7+}, K ^{+}}$

A.	$-2,-1,\frac{-1}{2}$
B.	$-2,-1,+2$
C.	$-2,-1,\frac{-1}{2},+2$
D.	$-2,\frac{-1}{2},+2$

Compound(atom)	Oxidation State of Atom
(I) $\mathrm{NO}\mathrm{(N)}$	(A) +1
(II) $\mathrm{H}_{2}\mathrm{O}_{2}\mathrm{(O)}$	(B) -1
(III) $\mathrm{MnO}_{4}^{-}\mathrm{(Mn)}$	(C) +2
(IV) $\mathrm{N}_{2}\mathrm{H_{2}}\mathrm{(N)}$	(D) +7

A.	I-A, II-B, III-C, IV-D
B.	I-C, II- B, III- D, IV- A
C.	I-D, II-B, III-A, IV-C
D.	I-C, II-A, III-D, IV-B

A.	$\mathrm{HNO}_3, \mathrm{NO}, \mathrm{N}_2, \mathrm{NH}_3$
B.	$\mathrm{N}_2, \mathrm{NO}, \mathrm{NH}_3, \mathrm{HNO}_3$
C.	$\mathrm{NO}, \mathrm{N}_2, \mathrm{NH}_3, \mathrm{HNO}_3$
D.	$\mathrm{NH}_3, \mathrm{~N}_2, \mathrm{NO}, \mathrm{HNO}_3$

A.	$\mathrm{Reduced: \mathrm{H}_2 \mathrm{~S} \rightarrow \mathrm{S}, Oxidised: \mathrm{KMnO}_4 \rightarrow \mathrm{MnSO}_4}$
B.	$Reduced: \mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{MnSO}_4, Oxidised: \mathrm{KMnO}_4 \rightarrow \mathrm{MnSO}_4$
C.	$Reduced: \mathrm{KMnO}_4 \rightarrow \mathrm{MnSO}_4, Oxidised: \mathrm{H}_2 \mathrm{~S} \rightarrow \mathrm{S}$
D.	$Reduced: \mathrm{KMnO}_4 \rightarrow \mathrm{MnSO}_4, Oxidised: \mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{MnSO}_4$

compound	oxidation state
(I) $\mathrm{H}_2 \mathrm{~S}$	A) +6
(II) $\mathrm{H}_2 \mathrm{~S}_{2}\mathrm{~O}_{3}$	B) +1
(III) $\mathrm{S}_2 \mathrm{Cl}_{2}$	C) -2
(IV) $\mathrm{Na}_2 \mathrm{SO}_{3}$	D) +4

A.	I-A, II-B, III-C, IV-D
B.	I-B, II-A, III-D, IV-C
C.	I-D, II-C, III-B, IV-A
D.	I-C, II-A, III-B, IV-D

compound	oxidation state
(I) $\mathrm{MnCl}_2$	A) +2
(II) $\mathrm{Mn}_2\mathrm{O}_{7}$	B) +7
(III) $\mathrm{Mn}_2\mathrm{O}_{2}$	C) +6
(IV) $\mathrm {K}_{2}\mathrm{MnO}_{4}$	D) +4

A.	I-A, II- B, III- D, IV-C
B.	I-D, II- C, III-B, IV-A
C.	I-B, II-D, III-A, IV-C
D.	I-C, II-D, III-B, IV-A

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Oxidation Number and Oxidation State - (Concept)

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A.	0 → +2, 0 → +3
B.	0 → -2, 0 → +2
C.	0 → +2, 0 → -3
D.	0 →+1, 0 → -1

A.	Methylene blue
B.	Cerric ions.
C.	Potassium dichromate
D.	Xylenol orange

A.	$\mathrm{BrF}$
B.	$\mathrm{HCO}_3^-$
C.	$\mathrm{H}_3 \mathrm{PO}_3$
D.	$\mathrm{PCl}_3$

A.	$\mathrm{As}_2 \mathrm{~S}_3 \rightarrow \mathrm{H}_3 \mathrm{AsO}_4$
B.	$\mathrm{As}_2 \mathrm{~S}_3+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{H}_2 \mathrm{SO}_4$
C.	$\mathrm{As}_2 \mathrm{~S}_3+\mathrm{HNO}_3 \rightarrow \mathrm{NO}$
D.	$\mathrm{HNO}_3 \rightarrow \mathrm{NO}$

A.	$\mathrm{OF_2}$
B.	$\mathrm{O_2F_2}$
C.	$\mathrm{OCl_2}$
D.	$\mathrm{O_2Cl_2}$

Compound	Oxidation state
I) $\mathrm{HClO_4}$	A) $+1$
II) $\mathrm{HClO}$	B) $+5$
III) $\mathrm{HClO_3}$	C) $+3$
IV) $\mathrm{HClO_2}$	D) $+7$

A.	I-D, II-A, III-B, IV-C
B.	I-A, II-B, III-C, IV-D
C.	I-C, II-A, III-D, IV-B
D.	I-B, II-C, III-D, IV-A

A.	$\mathrm{HSOH}$
B.	$\mathrm{H_2S_2O_3}$
C.	$\mathrm{H_2SO_5}$
D.	$\mathrm{H_2S_2O_2}$

A.	$\mathrm{HIO_3}$
B.	$\mathrm{HIO}$
C.	$\mathrm{H_5IO_6}$
D.	$\mathrm{H_4IO_4}$

Compound	Oxidation State
I) $\mathrm{HRuO_4}$	A) $+3$
II) $\mathrm{HNO_2}$	B) $+7$
III) $\mathrm{HPO_3}$	C) $+6$
IV) $\mathrm{H_6O_sO_6}$	D) $+5$