OXIDATION AND REDUCTION

OXIDATION AND REDUCTION IMPORTANT FACTS: IMPORTANT DEFINATIONS Many chemical reactions involve the addition of oxygen or hydrogen to the reactants. The reaction in which oxygen is added is called oxidation whereas the reaction which involves the addition hydrogen is called reduction. The two reactions always occur together, these reactions are called Redox reactions. 1.1 ELECTRONIC INTERPRETATION OF OXIDATION AND REDUCTION: Oxidation and reduction can be predicted in terms of the transference of electrons in the reactants. This is called as electronic concept of oxidation and reduction. According to this concept oxidation and Reduction are defined as follows. a) Oxidation is a process in which an atom or ion loses one or more electrons. It is also called de-electronation. Example: Na ------------ Na + + e- S 2- ------------ S + 2e- During oxidation there is an increase in positive valency or decrease in negative valency. The substances which can lose electrons during redox reactions are called reducing agents. b) Reduction is a process in which an atom or ion gains one or more electrons. It is also called electronation. Example: Cl + e- ---------------- Cl - Sn 4+ + 2e- ---------------- Sn 2+ During oxidation there is an increase in negative valency or decrease in positive valency. The substance which can gain the electrons during redox reactions are called oxidizing agents. Consider the following reaction occurring in aqueous medium. Zn (s) + 2HCl (aq) --------------- ZnCl 2(aq) + H 2(g)

In water medium it occurs as follows Zn (8) + 2H + (aq) + 2C1 - (aq) ------------- Zn 2+ (aq) + 2C1- (aq) + H 2(g) In this reaction Zn metal gets oxidized and hydrogen gets reduced. Lose Two electrons Zn(s) + 2H +( aq) Zn 2+ (aq) + H 2(g) Gain Two Electrons The two electrons from Zn(8) is transferred 2H(8) is transferred to 2H+ ions, hence it is a redox reaction. (Oxidation is also addition of electronegative element and reduction is addition of electropositive element. 1. The Oxidation State or oxidation number is the Charge an atom of an element would have if it existed as an ion in a compound. IN TERM OF OXIDATION REDUCTION ELECTRONS The Loss of electrons The Gain of electrons HYDROGEN The LOSS of hydrogen The GAIN of hydrogen OXYGEN The Gain of oxygen The Loss of oxygen OXIDATION STATE The INCREASE in the oxidation The DECREASE in the oxidation state of

state of the element the element 3. An Oxidizing Agent is a substance that cause oxidation on another substance. 4. A Reducing Agent is a substance that cause reduction on another substance. 5. A Redox Reaction is a chemical reaction in which oxidation and reduction both takes place. TEST YOURSELF: Define oxidation and reduction ( redox ) in terms of gain and loss of oxygen hydrogen, electron transfer and charges in oxidation state 1.1 OXIDATION AND REDUCTION RECTIONS: 1. A substance is Oxidised if it there is a Gain of Oxygens in a chemical reaction. EXAMPLE1: Oxidation ( Gain of Oxygen ) 4Na (s) + O 2 (g) 2 Na 2 O (s) Sodium Gains Oxygen and is Oxidised to sodium oxide. Sodium oxygen sodium oxide 2. A substance is Oxidised if it there is Loss of Hydrogens in a reaction.

EXAMPLE 2: Oxidation ( Loss of Hydrogen ) H 2S Looses Hydrogen and is Oxidised to sulfur H 2 S (g) + CI 2 (g) 2 HCI (g) + S(s) Hydrogen sulfide chlorine hydrogen chloride sulfur 3. A substance is Oxidised if there is loss of Electrons in a chemical reaction. EXAMPLE 3: Oxidation ( Loss of Electrons ) Mg (s) + CI 2 (g) MgCI 2 (s) Magnesium chlorine Magnesium chloride Electrons have been transferred during this reaction, shown by the following half equations. Mg (s) CI 2 (g) 4. A substance is Reduced if there is Loss of Oxygens in a chemical reaction.

EXAMPLE4 : Reduction ( Loss of Oxygen ) The copper (II) oxide loses oxygen and is reduced to copper metal CuO (s) + H 2 (g) Cu(s) + H 2 O(g) Copper (II) oxide hydrogen copper water 5. A substance is Reduced if it Gains Hydrogen in a reaction. EXAMPLE5 : Reduction ( Gain of Hydrogen ) Chlorine Gains Hydrogen and is reduced to hydrogen chloride H 2 (g) + CI 2 (g) 2HCI(g) Hydrogen chlorine hydrogen chloride 6. A substance is Reduced if there is Gain of Electrons in a chemical reaction.

EXAMPLE6 : > CI - ions spectator ions. Reduction ( Gain of Electrons ) > Iron (III) ion gains an electron and is reduced to iron (II) ion. 2FeCI 3 (aq) + H 2 S (g) 2FeCI 2 (aq) + 2HCI (aq) + S (s) Iron (III) hydrogen Iron (II) hydrogen sulfur Chloride sulfide chloride chloride The ionic equation for the reaction: 2Fe 3+ (aq) + H 2 S (g) 2Fe 2+ (aq) + 2H + (aq) + S (s) EXAMPLE7 : Which compound is (a) oxidized and (b) reduced in the following reaction? 2NH 3 (g) + 3CuO( s) N 2 (g) + 3CuO (s) + 3H 2 O(I) SOLUTION:

Oxidation ( Loss of Hydrogen ) 2NH 3 (g) + 3CuO( s) N 2 (g) + 3CuO (s) + 3H 2 O(I) Reduction ( Loss of Oxygen ) 7. OXIDATION STATE: The Oxidation State (oxidation number) is the charge an atom of an element would have if it existed as an ion in a compound. Oxidation state is also called Oxidation Number. The oxidation state can be a positive number (e.g.+3), a negative number (e.g.-3) or zero. The oxidation state of an element can be determined by using the following rules. RULE EXAMPLE OXIDATION The oxidation state of an element in the uncombined state (free state) is zero. Na C Mg O 2 0 0 0 0 The oxidant state of a simple ion is equal to the charge on the ion. Na + Fe 3+ O 2- P 3- +1 +3-2 -3

The oxidation states of Groups I and II elements in their compounds are fixed (refer to Periodic Table) Group I elements Group II elements +1 +2 The oxidation states of hydrogen and oxygen in their compounds are fixed H ( in H 2 O ) O ( in CuO ) +1-2 The oxidation states of the atoms present in the formula of the compound add up to zero. AI 2 O 3 2AI = 2 X (+3) = +6 3O = 3 X (2-) = -6 Total = (+6 ) + (-6) = 0 The oxidation states of all the atoms in a polyatomic ion is equal to the charge on the ion. SO 4 2- S = + 6 40 = 4 X (-2) = - 8 Total = (+6) + (-8) = -2 COMMON ERROR ACTUAL FACTS The reaction as shown below is a neutralization reaction : Acid + Base Salt and Water The reaction is not a neutralization reaction as chlorine is also produced other than salt and water. ( Comment : It is a redox reaction. MnO 2 oxidises 4HI(aq) + MnO 2 (s) MnCI2 (aq) + H2O (l) + Cl2 (g) HCI to CI 2 and itself reduced to MnCI 2. )

1.3 OXIDATION NUMBER: Oxidation number is defined as the charge present on the atom of an element which is present in the combined state. It may have a positive or negative value. The oxidation number may be real or apparent charge. Difference between valency and oxidation number: Valency Oxidation number 1. It is the combining capacity of an element. 1. It is a charge assigned to an atom or ion ion in a molecule by using arbitrary rules. 2. It has no negative or positive sign 2. It carries a sign ve or + ve. 3. Valency is a fixed value 3. It is not fixed for an element. It Depends on the compound. 4. It is a whole number but never zero 4. Is is a whole number, fraction or Even zero. Table 4.1 : Differences between Valency and Oxidation Number 1.4 RULES FOR COMPUTING OXIDATION NUMBER: The following arbitrary rules have been adopted, to calculate oxidation number or Element on the basis of periodic properties of elements. 1. In an uncombined state or free state the ON in zero. 2. In the combined state of elements the ON s are a) F = -1 b) O = -2 (peroxides O O - =1) In F2O it is +2 c) H = +1 d) Metals always + Ve e) Alkali and Alkaline earth metals = +1 and + 2. f) Halogen = -1 g) Sulphur = - 2

3. The algebraic sum of all oxidation numbers of elements in a compound is zero Example in K 2 MnO 4 4. The algebraic sum of all oxidation numbers of elements in a radical is equal to net charge on that radical Example C2O4 2-5. Maximum oxidation number (except O and F) = Group number. Minimum oxidation number (except metals) = group number 8. (Note group number is an Mendeleev s modern periodic table). 1.4 CALCULATION OF OXIDATION NUMBERS: Examples: 1. KMnO 4 Let on O.N. of Mn be x 1 + x + 4 (-2) = 0 1 + x 8 = 0 1 + x =8 x = 8-1 x = +7 The ON of Mn = + 7 2. H 2 SO 4 Let the O.N. of s be x 2 + x + 4 (-2) = 0 2 + x -8 = 0 x = +6 3 K 2 Cr 2 O 7 Let the O.N. of Cr be x 2+2x+7(-2)=0 2 + 2x 14 = 0 x = +6 3. H 3 POa Let the O.N. of P be x 3 + x + 4 (-2) = 0 3 + x + (-8) = 0 x = -5

4. H 3 PO 4 Let the O.N. of P be x 3 + x + 4(-2) =0 3 +x+(-8) =0 x= -5 5. HNO 3 Let the O.N. of N be x 1 + x-6=0 X=6-1 X = +5 6. (Fe(CN) 6 )3- Let the O.N. of Fc be x X + 6(-1) = -3 x-6 = -3 x = -3+6 x = +3 7. MnO 4 - Let the O.N. of Mn be x X + 4 (-2) = 0 X +(-8) = 0 X 8 = 0 x = +8 8. Cr(H 2 O) 3+ Let the O.N. of Mn be x The ON of H 2 O is zero as it is neutral molecule X + 6 (0) = +3 x = +3 9. Na 2 S 4 O 6 Let the O.N. of S be x 2 + 4x +6 (-2) = 0 2+4x-12=0 2+4x=12 4x=12-2

4x=10 10 5 X = --- = -- x = + 5 ---- 2 4 2 10. FeSO 4 (NH 4 ) SO 4 6H 2 O Let O.N. of Fe be x X +(-2) +0+0=0 x = +2 1.5 BALANCING EQUATIONS: The chemical equations involving redox reactions can be balanced by following two Methods. a) Oxidation number method b) Ion electron Method (Half reaction method) We shall learn the balancing of equations involving redox reactions by oxidation number method. 1.6a) Oxidation number method: The following steps must be applied in balancing a redox equation by oxidation number method. 1. Write the skeleton equation. 2. Write the oxidation numbers of all elements on their symbols.

3. Identify the elements that undergo change in oxidation number. 4. Calculate the increase or decrease in ON per gram atom, with respect to the reactants. If more than one atom is present, then mulply the number of atoms undergoing change to calculate the total change in oxidation number. 5. Equate the increase and decrease in ON on the reactant side by multiplying the oxidizing and reducing agents suitably. 6. Balance the equations except hydrogen and oxygen. Later balance hydrogen and oxygen also. 7. In reactions occurring in acid medium, balance Oxygen by adding H 2 O where Oxygen is less. Then balance Hydrogen atoms by adding H + where Hydrogen is less. 8. In Basic medium, balance negative charges by adding OH where negative charge is less. Then and H 2 O molecules to other side and balance the equation. EXAMPLE8: Calculate the oxidation states of nitrogen in the following compounds: a) NH 3 b) N 2 O 4 c) NO 3 - SOULATION: a) Let the oxidation state of N = x. x 3 X (+ 1) x + 3 X (+1) = 0 (Rule 5 ) x = - 3 NH 3 The oxidation state of nitrogen in NH 3 is 3. b) Let the oxidation state of N = y. 2y 4 X (-2) 2y + 4 X (-2) = 0 (Rule 5 ) Y = + 4 N 2 O 4

The oxidation state of nitrogen in N 2 O 4 is +4. c) Let oxidation state of N = z. z 3 X (-2) z + 3 X (-2) = -1 ( Rule 6 ) z = + 5 - NO 3 The oxidation state of nitrogen in NO - 3 is + 5, TIPS FOR STUDENT Oxidation state is not written in the same way as the charge on an ion. For example, in PbCl 2, oxidation state of Pb is +2, but the charge on Pb 2+ is 2+. Similarly, the oxidation state of phosphorus in Na 3 P is -3, but the charge on P 3- is 3-. 8. To check whether oxidation has taken place in any reaction, follow the three steps : a) Write the balanced equation for the reaction. b) Write the oxidation states of all the substances in the reaction. c) Compare the oxidation states to check which reactant has been oxidised. 9. Oxidation occurs when the oxidation state of an element increases. Reduction occurs when the oxidation state of an element decreases. EXAMPLE9 :

Oxidation ( Increase in Oxidation State ) +2 +3 2FeCl 2 (s) + Cl(g) 2FeCl 3 (s) 0-1 Reduction ( Decrease in Oxidation State ) In this reaction, the oxidation number of Iron increases from +2 in FeCl 2 to +3 in FeCl 3. This is an oxidation process. Chlorine decreases from O in Cl 2 to -1 in Cl -. This is a reduction process. 10. Oxidation and reduction always occur together. If one reactant is oxidised, the other reactant must be reduced. We call the combined process the redox reaction. TIP FOR STUDENT

COMMON ERROR ACTUAL FACTS The oxidation number of hydrogen is 1 and the oxidation number of Oxygen is 2. The + and signs must be shown. Thus, the oxidation number of hydrogen +1 the oxidation number of oxygen is -2. 10.2 OXIDISING AND REDUCING AGENTS: 1. A substance that oxidizes other substances is called an Oxidizing Agent. An oxidising agent is reduced when it oxidises another substance. 2. A substance that reduces other substances is called an reducing Agent. An reducing agent is oxidised when it reduces another substance. EXAMPLE10 :

In the extraction of iron from iron ( III ) oxide, the following reaction occurs between iron ( III ) oxide and carbon monoxide. Reduced ( Loss of Oxygen ) +3 0 Fe 2 O 2 (s) + 3CO(g) 2Fe( l ) + 3CO 2 (g) oxidising agent +2 +4 reducing agent Oxidised ( Gain Of Oxygen ) a) Iron ( III ) oxide is an oxidizing agent. It oxidises carbon monoxide to carbon dioxide and is itself reduced to iron. b) Carbon monoxide is a reducing agent. It reduces iron (III) oxide to iron and is itself oxidised to carbon dioxide. 3. In terms of electron transfer, a) An oxidising agent is an acceptor of electrons, b) A reducing agent is a donor of electrons.

EXAMPLE 11: Reaction between chlorine gas and potassium iodide solution. Cl2 (g) + 2KI (aq) 2KCl (aq) + I 2 (aq) Chlorine Potassium iodide Potassium chloride iodine Chlorine gas is the acceptor of electron, and is thus an oxidising agent. Cl 2 (g) + 2e - 2CI - (aq) reducation An oxidising agent undergoes reduction. The iodide ion is the donor of electrons, and is thus a reducing agent. 2I - (aq) I 2 (aq) + 2e - oxidation An reducing agent undergoes oxidation EXAMPLE 12: Identify the oxidising and reducing agents in the following reaction. a) PbO(s) + H 2 (g) Pb(s) + H 2 O (l) b) Zn(s) + CuSO 4 ( aq) ZnSO 4 (aq) + Cu (s) SOLUTION:

a) Reduction ( loss of oxygen ) PbO(s) + H 2 (g) Pb (s) + H 2 O(l) Oxidation ( gain of oxygen) Lead (ll) oxide, PbO, is the oxidising agent. It oxidizes hydrogen to water. Hydrogen is the reducing agent. It reduces leas (ll) oxide to lead. b) Oxidation ( donor of electrons ) Zn(s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu(s) Reduction ( acceptor of electrons ) In this reaction, Zn(s) Zn 2+ (aq) + 2e - ( electron acceptor ) Cu 2+ (aq) + 2e - Cu(s)

( electron acceptor ) Zinc metal is the reducing agent. It reduces Cu 2+ to Cu by donating electrons to Cu 2+. Cu 2+ ion is the oxidising agent. It oxidizes Zn to Zn 2+ by accepting electrons from Zn. 4. Some substances can act as both an oxidising agent and a reducing agent. EXAMPLE: Sulfur dioxide, SO 2 acts as an oxidising agent in reaction (1) but as a reducing agent in reaction (2). -2 0 SO 2 (g) + 2H 2 S (g) 3S (s) + 2H 2 O (I) ( 1 ) oxidising agent 0-2 2SO 2 (g) + O 2 (g) 2SO 3 (g) ( 2 ) Reducing agent 5. The table below shows some common oxidising and reducing agents. OXIDISING AGENT REDUCING AGENT Chlorine, CI 2 Potassium iodide KI

Bromine, Br 2 Carbon C Nitric acid HNO 3 Carbon monoxide, CO Hydrogen Peroxide, H 2 O 2 Ammonia, NH 3 Potassium manganate (VII), KMnO 4 Sulfur dioxide, SO 2 Potassium dichromate(vi), K 2 Cr 2 O 7 Hydrogen, H 2 Oxygen, O 2 Hydrogen sulfide, H 2 S Concentrated sulfuric acid, H 2 SO 4 Metals 6. TEST DIAGRAM OBSERVATION AND INFERENCE EXPLANATION Add potassium iodide solution (colourless). Potassium Iodide If the solution turns brown, The substance is an oxidising agent. An oxidising agent oxidises iodide ion to iodine. 2I - (aq) I 2 (aq) +2e - (Colourless) (Brown) Unknown Brown

Substance Solution Test with starch-iodide paper. Moist starch-iodide paper An oxidising agent changes the colour of moist starch-iodide paper from white to blue, The iodine reacts with the starch to give a blue colour. 7. TEST DIAGRAM OBSERVATION AND INFERENCE EXPLANATION Add acidified potassium dichromate (VI) solution. Acidified potassium dichromate (VI)solution If the orange solution turns green, The substance is a reducing agent. The reducing agent reduces the dichromate (VI) ion, Cr 2 O 7 2-, to chromium(iii) ion, Cr 3+. Dilute sulfuric acid is always used to acidify Potassium dichromate (VI) solution. Orange solution Green solution 2- Cr 2 O 7 (aq) + 14H + (aq) + 6e - (orange) 2Cr 3+ (aq) + H 2 O(I) (green)

Add acidified potassium managnate (VII) solution. Acidified potassium Managnate (VI)solution If the purple solution turns colourless (decolourisation), the substance is a reducing agent. The reducing agent reduces the managnate (VII) ion to managanese (II) ion. MnO 4 - (aq) + 8H + (aq) + 5e - (Purple) Mn 2+ (aq) + H 2 O(I) (colourless) Purple solution Colourless solution COMMON ERROR ACTUAL FACTS A substance is always an oxidising agent or a reducing agent in all reactions. Some substance, such as sulfur dioxide, SO 2, hydrogen peroxide, H 2 O 2, and sodium nitrite, NaNO 2, can act as both oxidising and reducting agents. SUMMARY AND KEY POINTS 1.)The reaction in which oxygen is added is called oxidation whereas the reaction which involves the addition hydrogen is called reduction. The two reactions always occur together, these reactions are called Redox reactions. 2.) Oxidation is a process in which an atom or ion loses one or more electrons. It is also called de-electronation.

Example: Na ------------ Na + + e- S 2- ------------ S + 2e- During oxidation there is an increase in positive valency or decrease in negative valency. The substances which can lose electrons during redox reactions are called reducing agents. 3.) Reduction is a process in which an atom or ion gains one or more electrons. It is also called electronation. Example: Cl + e- ---------------- Cl - Sn 4+ + 2e- ---------------- Sn 2+ During oxidation there is an increase in negative valency or decrease in positive valency. The substance which can gain the electrons during redox reactions are called oxidizing agents. 4.) The Oxidation State or oxidation number is the Charge an atom of an element would have if it existed as an ion in a compound. IN TERM OF OXIDATION REDUCTION ELECTRONS The Loss of electrons The Gain of electrons HYDROGEN The LOSS of hydrogen The GAIN of hydrogen OXYGEN The Gain of oxygen The Loss of oxygen OXIDATION STATE The INCREASE in the oxidation state of the element The DECREASE in the oxidation state of the element

5.) An Oxidizing Agent is a substance that cause oxidation on another substance. 6.) A Reducing Agent is a substance that cause reduction on another substance. 7.) A Redox Reaction is a chemical reaction in which oxidation and reduction both takes place. 8.) A substance is Oxidised if it there is a Gain of Oxygens in a chemical reaction. 9.) A substance is Oxidised if it there is Loss of Hydrogens in a reaction. 10.) A substance is Oxidised if there is loss of Electrons in a chemical reaction. 11.) A substance is Reduced if there is Loss of Oxygens in a chemical reaction. 12.) A substance is Reduced if it Gains Hydrogen in a reaction. 13.) A substance is Reduced if there is Gain of Electrons in a chemical reaction. 14.) OXIDATION STATE: The Oxidation State (oxidation number) is the charge an atom of an element would have if it existed as an ion in a compound. Oxidation state is also called Oxidation Number. The oxidation state can be a positive number (e.g.+3), a negative number (e.g.-3) or zero. The oxidation state of an element can be determined by using the following rules.

15.) Oxidation number is defined as the charge present on the atom of an element which is present in the combined state. It may have a positive or negative value. The oxidation number may be real or apparent charge. 16.) The algebraic sum of all oxidation numbers of elements in a compound is zero. Example in K 2 MnO 4 17.) The algebraic sum of all oxidation numbers of elements in a radical is equal to net charge on that radical. 18.) Maximum oxidation number (except O and F) = Group number. Minimum oxidation number (except metals) = group number 8. (Note group number is an Mendeleev s modern periodic table). KEY POINTS: i)oxidation state is not written in the same way as the charge on an ion. ii)for example, in PbCl 2, oxidation state of Pb is +2, but the charge on Pb 2+ is 2+. Similarly, the oxidation state of phosphorus in Na 3 P is -3, but the charge on P 3- is 3-. 19.) Oxidation occurs when the oxidation state of an element increases. Reduction occurs when the oxidation state of an element decreases.

20.) Oxidation and reduction always occur together. If one reactant is oxidised, the other reactant must be reduced. We call the combined process the redox reaction. 21.) A substance that oxidizes other substances is called an Oxidizing Agent. An oxidising agent is reduced when it oxidises another substance. 22.) A substance that reduces other substances is called an reducing Agent. An reducing agent is oxidised when it reduces another substance.