Oxidation Number OR Oxidation State:

The donation of electrons is called the oxidation and the gain of electrons is called the reduction. Oxidation and reduction can further be explained by a knowledge of “Oxidation number”.

The oxidation state of an atom in its free or ground state is taken as zero. When the atom loses electrons its oxidation state increases and when the atom gains electrons its oxidation state decreases. The term oxidation-number represents the positive or negative character of the atom in a compound.

Oxidation number is defined as the charge an atom appears to have when electrons are assigned in accordance with the following arbitrary rules.

Electrons shared by two like atoms are divided equally between the two atoms. Electrons shared between two unlike atoms are assigned to the more electronegative atom of them.

Conventions Used in Assigning Oxidation Number or Oxidation State:

• The oxidation number of an element in a free atomic state (Na, H, Cl, O, P etc) or in its poly-atomic state (graphite, H₂, Cl₂, O₂ etc) is always zero.
• The oxidation number of hydrogen is always +1 in its compounds.  However, in metal hydrides like NaH, MgH₂  etc. the oxidation number of hydrogen is -1 because metals are more electropositive than hydrogen.
• O.N. of oxygen is always -2 in its compounds.  However, in peroxides like H₂O₂, Na₂O₂, BaO₂ etc. the oxidation number of oxygen is -1. In OF the oxidation number of oxygen is +2 because F is more electronegative than O.
• O.N. of group IA element i.e. Li, Na, K etc is always +1 in their compounds.
• O. N. of group IIA elements i.e. Be, Mg, Ca, Sr and Ba are always +2 in their compounds.
• O. N. of F is always -1 in its compounds because it is most highly electronegative.  Oxidation O. N. of other elements of group VIIA. (17) i.e. Cl, Br and I are also generally –1.
• In an ion, the sum of the oxidation numbers of different atoms is equal to charge over the ion.
• In a complex compound (involving co-ordination by ligands) it is more convenient to use oxidation number of group (ligand) as a whole instead of the oxidation number of individual atoms. For example, in HCN the oxidation number of CN- ion is –1. Here CN-  as a whole is considered and not of individual C or N.
• on the basis of the above standard oxidation numbers, which may be taken as rules, the oxidation, a number of a particular given atom in a compound can be determined.

Valency and Oxidation State:

Valency is a different term than oxidation number though sometimes the valency and the oxidation number of an element are same in a compound. Valency of an element is given by the number of electrons it actually loses or gains or shares during the formation of a compound, Whereas oxidation number is just the apparent charge (not necessarily actual) over the atom when the electrons are counted according to the arbitrary rules given earlier.

In most of the cases, the valency of an element is constant whereas the oxidation state of an element may vary in its different compounds. Valency and oxidation states of carbon in its different compounds give a good example of this. In CH₄, CH₃Cl, CH₂Cl₂, CHCl₃ and CCl₄ the valency of carbon is always four (due to sharing of four electrons) but its oxidation number is – 4, -2, 0, +2 and +4 respectively.

Oxidation-Reduction in Terms of Oxidation Number:

On the basis of oxidation number a reaction involving the increase in oxidation number is called as oxidation while a reaction involving the decrease in oxidation number is called as reduction (Remember increase in O.N. means increase in positive O.N. or decrease in negative O.N., while decrease in O.N. means decrease in positive O.N. or increase in negative O.N.).

For example, in the reaction, 2Mg + O₂ →  2MgO, The O.N. of Mg increase from 0 to +2.while the O.N. of O decreases from 0 to -2.  Thus, magnesium is oxidised while oxygen is reduced.

Science > Chemistry > Redox Reactions > Oxidation Number or Oxidation State

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In this article we shall study about redox reactions, in which both the oxidation and reduction reactions take place simultaneously.

Oxidation:

Old Concept:

• It is a process in which addition of oxygen takes place.

2Mg + O₂  →      2MgO

• It is a process in which addition of electronegative radical takes place.

2FeCl₂ +  Cl₂ →  2FeCl₃

• It is a process in which removal of hydrogen takes place.

H₂S + 2 [Cl] →  S + 2HCl

• It is a process in which removal of electropositive radical takes place.

2KI + H₂O₂ → I₂ + 2KOH

Moden Concept:

According to the electronic concept, a reaction in which loss of electrons from an atom or an ion takes place is called oxidation. Consider reaction, 

2Mg + O₂ →  2MgO

In this reaction, the valency of magnesium is increased’ from zero (in the atomic state) to + 2 (in MgO).

i.e.  Mg⁰ →  Mg²⁺ + 2 e^–

In this reaction, magnesium is losing electrons. And hence oxidation of magnesium takes place.

Reduction:

Old Concept:

• It is a process in which addition of hydrogen takes place.

Cl₂ + H₂  →  2HCl

• It is a process in which addition of electropositive radical takes place.

2HgCl₂ +  SnCl₂ → Hg₂Cl₂ + SnCl₄ .

• It is a process in which removal of oxygen takes place.

CuO + 2 [H] → Cu + 2H₂O

• It is a process in which removal of electronegative radical takes place.

FeCl₃ + H → FeCl₂ + HCl

Modern Concept:

According to the electronic concept, a reaction in which the gain of electrons by an atom or an ion takes place is called reduction. Consider reaction, 

2HgCl₂ +  SnCl₂ → Hg₂Cl₂ + SnCl₄

In this reaction, the valency of mercury is decreased’ from +2  (in HgCl₂) to +1 (in Hg₂Cl₂).

i.e.   Hg²⁺ +   e^–  →   Hg ⁺

In this reaction mercury is gaining electron. And hence reduction of mercury takes place.

Redox Reaction:

In any of a chemical reaction if one of the reactants is oxidized, other is surely reduced, i.e. oxidation and reduction always take place simultaneously.

Example – 1: 

2Mg + O₂ →  2MgO  

Mg is oxidized to MgO (addition of oxygen, i.e. increase in positive valency of Mg i.e. loss of electrons by Mg), whereas oxygen is reduced to MgO (addition of positive radical, i.e. increase in negative valency of oxygen, i.e. gain of electrons by oxygen)

Example – 2: 

2HgCl₂ +  SnCl₂ → Hg₂Cl₂ + SnCl₄

HgCl₂ is reduced to Hg₂Cl₂ whereas SnCl₂ is oxidised to SnCl₄.

Thus oxidation and reduction take place simultaneously.  Therefore, all such reactions are called as reduction-oxidation reactions or redox reactions.  In all such reactions, one of the reactants loses the electrons (oxidized) while other gains those electrons (reduced)

However, it should be remembered that all the chemical reactions are not redox reactions.  There are several other types of reactions also.

NaCl + AgNO₃ → AgCl + NaNO₃

In such reactions none of’ the reactants is oxidized or reduced; simply the exchange of cation or anion takes place.

Oxidizing Agent (Oxidant):

The substance which excepts electrons and makes the other substance to lose electrons is called oxidizing agent or oxidant.

Example – 1:

Consider reaction, 

2Mg + O₂ →  2MgO

In this reaction, oxygen is making magnesium to lose electrons and hence in this reaction oxygen is the oxidizing agent.

Example – 2:

Consider reaction,

2K + Cl₂  → 2KCl

In this reaction, chlorine is making potassium to lose an electron and hence in this reaction chlorine is the oxidizing agent.

Characteristics of Oxidizing Agent:

• The substance which excepts electrons and makes the other substance to lose electrons is called oxidizing agent or oxidant.
• In a reaction, the oxidizing agent oxidizes the other substance but is itself reduced.
• Oxygen, or a substance capable of giving oxygen, is always a good oxidizing agent.
• According to electron concept, an oxidizing agent is that which is capable of de-electronating the other substance.
• An oxidizing agent is an electron acceptor and during the redox reaction, it is electronated.
• Fluorine (F) has a maximum tendency to accept electrons hence it is the strongest oxidizing agent.

Examples of Common oxidizing Agents:

Oxygen (O or O₂), Ozone (O₃), Hydrogen peroxide (H₂O₂), Sulphuric acid (H₂SO₄), Nitric acid (HNO₃), Perchloric acid (HClO₄), Potassium chlorate (KClO₃), Acidified potassium dichromate (K₂ Cr₂ O₇ + H₂SO₄), Acidified potassium permanganate (KMnO₄ + H₂SO₄), Alkaline potassium permanganate (KMnO₄ + KOH)

Reducing Agent (Reductant):

The substance which loses electrons and makes the other substance to gain electrons is called reducing agent or reductant.

Example – 1:

Consider reaction, 

2Mg + O₂ →  2MgO

In this reaction, magnesium is making oxygen to gain electrons and hence in this reaction magnesium is reducing agent.

Example – 02:

Consider reaction, 2K + Cl₂  → 2KCl, In this reaction potassium, is making chlorine to gain an electron and hence in this reaction potassium is reducing agent.

Characteristics of Reducing Agent:

• The substance which loses electrons and makes the other substance to gain electrons is called reducing agent or reductant.
• In a reaction, the reducing agent reduces the other substance but is itself oxidised.
• Hydrogen, or a substance capable of giving hydrogen, is always a good reducing agent.
• According to electron concept, a reducing agent is that which is capable of electronating the other substance.
• A reducing agent is an electron donor and during the redox reaction, it is de-electronated.
• Sodium (Na) has a maximum tendency to donate electron hence it is the strongest reducing agent.

Examples of Common Reducing Agents:

Hydrogen (H or H₂), Hydrogen iodide (HI), Hydrogen sulphide (H₂S), Lithium aluminium hydride (LiAI H₄), Sodium borohydride (NaB H₄), Sulphur dioxide (SO₂), Carbon (C), Ozone (O₃), Hydrogen peroxide (H₂O₂), Tin & hydrochloric acid (Sn  + HCl), Sodium & alcohol (Na + C₂ H₅OH), Metallic salts (ous) like SnCl₂ , FeSO₄ etc.

Science > Chemistry > Redox Reactions > Introduction to Redox Reactions

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A series of electrodes or half cells arranged in order of their increasing standard oxidation potentials or in the decreasing order of their standard reduction potentials is called an electromotive force series or electrochemical series. Electrochemical series is also known as e.m.f. series

Characteristics Electrochemical Series:

• In this series, all reduction potentials are given on hydrogen scale whose, Eo is taken as zero.
• The standard reduction potential of an element is a measure of the tendency of that element to get reduced.
• The element which has greater reduction potential gets reduced easily.  While the elements with low reduction potential will get easily oxidized
• Elements that lose electrons more easily have lower (negative) reduction potential and those which lose electrons with greater difficulty or instead of losing they accept electrons more easily have a higher (positive) reduction potential.
• In EMF series elements having higher (+ ve), the reduction potential is placed at the top.  While those having lower (-ve) reduction potential are placed at the bottom.  SHE has the middle position in the electrochemical series.
• The substances which are stronger reducing agents than hydrogen are placed below the hydrogen in the series and have negative standard reduction potential. The substances which are weaker reducing agents than hydrogen are placed above the hydrogen in the series and have positive standard reduction potential. Thus as we move down the group strength of reducing agent increases while the strength of the oxidizing agent decreases.
• Metal at the bottom is the most active metal. As we move down in the series activity and electropositivity of metals increase. Nonmetal at the Top is the most active nonmetal. As we move down in the series activity and electronegativity of nonmetal decreases.

Applications of Electrochemical Series:

For Choosing Elements as Oxidising Agents:

The elements which have more electron-accepting tendency are oxidizing agents. Elements at the top of the electrochemical series have higher (+ ve) reduction potential.  Hence they gain an electron from other elements and oxidize them.  So they are good oxidizing agents.

Element (F₂) at the topmost position of electrochemical series which has the highest reduction potential is the strongest oxidizing agent. Oxidizing power decreases from top to bottom in the series.

e.g. The elements like Cu, Ag, Hg, Br₂, Cl₂, etc. are good oxidizing agents. F₂ is the strongest oxidizing agent.

For Choosing Elements as Reducing Agents:

The elements which have more electron losing tendency are reducing agents. The elements at the bottom in the electrochemical series have lower (- ve) reduction potential. Hence they lose electrons readily and supply to other elements and reduce them. So bottom elements in electrochemical series are reducing agents.  Reducing strength goes on increasing from top to bottom in the series.

Element (Li) having the bottom-most position has the lowest reduction potential hence it is the strongest reducing agent.

e.g. The element like Zn, Cd, Ni, K, etc. are good reducing agents.

For Studying displacement reaction:

One metal can be displaced from a salt solution by another metal is known as a redox reaction. Elements having higher reduction potential will gain electrons and that having lower reduction potential will lose electrons. Hence element lower in electrochemical series can displace an element placed higher in electrochemical series from its salt solution.

Example-1:

Zn displaces Cu from CuSO₄, because, zinc is placed lower in electrochemical series and has lower reduction potential while Cu is placed higher in electrochemical series and has higher reduction potential.  Hence zinc can easily displace copper from CuSO₄

Zn + CuSO₄ → ZnSO₄ + Cu    i.e.

Zn + Cu⁺⁺_(aq) → Zn ⁺⁺_(aq) +  Cu

Example-2:

Fe displaces Cu from CuSO₄ because Fe is placed lower in electrochemical series and has lower reduction potential while Cu is placed higher in electrochemical series and has higher reduction potential.  Hence Fe can easily displace copper from CuSO₄.

Fe + CuSO₄ → FeSO₄(aq)+ Cu  i.e.

Fe + Cu⁺⁺_(aq) → Fe⁺⁺_(aq) +  Cu

To predict whether a given metal will displace another, from its salt solution:

A metal lower in the series will displace the metal from its solution which is higher in the series, i.e., the metal having low standard reduction potential will displace the metal from its salt’s solution which has a higher value of standard reduction potential. A metal lower in the series has a greater tendency to provide electrons to the cations of the metal to be precipitated.

Displacement of one nonmetal from its salt solution by another nonmetal:

A nonmetal higher in the series having the high value of standard reduction potential will displace another nonmetal with lower reduction potential i.e., occupying the position below in the series. The nonmetal’s which possess high positive reduction potentials have the tendency to accept electrons readily. These electrons are provided by the ions of the nonmetal having the low value of reduction potential. Thus, Cl2 can displace bromine and iodine from bromides and iodides.

Displacement of hydrogen from dilute acids by metals:

The metal which can provide electrons to H⁺ ions present in dilute acids for reduction evolve hydrogen from dilute acids. The metal having negative values of reduction potential possess the property of losing electron or electrons. Thus, the metals occupying lower positions in the electrochemical series readily liberate hydrogen from dilute acids and on ascending in the series tendency to liberate hydrogen gas from dilute acids decreases.

The metals which are above hydrogen in electrochemical series like Cu, Hg, Au, Pt, etc., do not evolve hydrogen from dilute acids.

Displacement of hydrogen from water:

Iron and the metals below iron are capable of liberating hydrogen from water. The tendency increases from top to bottom in electrochemical series. Alkali and alkaline earth metals liberate hydrogen from cold water but Mg, Zn and Fe liberate hydrogen from hot water or steam.

For Calculation of standard EMF of cell ( E^o_cell):

From the standard electrode potential values, it is easy to calculate EMF of cell. Standard oxidation potential values are given in EMF series. Eo  cell is calculated using formula:

E^o_cell =    E^o_{red (cathode)}    –    E^o_{red (anode)}

e.g. in Daniell cell,

Now, From the series, E^o_Zn  = – 0.763 V ,  E^o_Cu =  + 0.337 V

E^o_cell =    E^o_{red (cathode)}    –    E^o_{red (anode)}

∴  E^o_cell =    E^o_{red (Cu)}    –    E^o_{red (Zn)}

∴  E^o_cell =     0.337  –  ( -0.763)

∴  E^o_cell =     0.337  + 0.763

∴  E^o_cell =     1.10 V

For Checking Spontaneity of Redox Reactions:

If cell is assembled such that one electrode has higher positive oxidation potential and other has lower negative oxidation potential then redox cell reaction will be spontaneous and cell will have positive EMF.  On the contrary if EMF of cell is negative then redox cell reaction will be non spontaneous.

e.g. in Daniell cell,

Now, From the series, E^o_Zn  = – 0.763 V ,  E^o_Cu =  + 0.337 V

E^o_cell =    E^o_{red (cathode)}    –    E^o_{red (anode)}

∴  E^o_cell =    E^o_{red (Cu)}    –    E^o_{red (Zn)}

∴  E^o_cell =     0.337  –  ( -0.763)

∴  E^o_cell =     0.337  + 0.763

∴  E^o_cell =     1.10 V

Since cell has positive EMF, following redox cell reaction is spontaneous.

Zn + Cu⁺⁺_(aq) → Zn ⁺⁺_(aq) +  Cu

Thus higher the positive EMF of the cell, the more is the spontaneity of the redox cell reaction.

For Construction of a Cell:

Various cells can be constructed by combining standard electrodes given in EMF series as per the requirement of e.m.f.

e.g If a cell of e.m.f. 1.1 V is required, then from e.m.f. series we can locate zinc and copper electrode whose combination gives required e.m.f.

For Corrosion Treatment:

When two metals which are in contact with each other are exposed to the atmosphere, the element lower in series will be oxidized. i.e. it is rusted and destroyed.

If there is a scratch on the galvanized sheet of iron, and iron is exposed then zinc is rusted and iron is protected. This is because in e.m.f. series zinc is below the iron. But if there is a scratch on the tin-plated iron, iron gets rusted because in e.m.f. series iron is below tin.

To Find Reactivity of Metals:

As we move down in the electrochemical series reactivity of metal increases. Alkali metals and alkaline metals at the bottom are highly reactive. They can react with cold water and evolve hydrogen. They can dissolve in acid-forming salt.

Metals like Fe, Pb, Sn, Ni, Co which are in little higher in the series do not react with cold water but react with steam and evolve hydrogen. Metals like Cu, Ag, and Au which lie above the hydrogen are less reactive and do not react with water in any form to evolve hydrogen.

To Ascertain Electropositivity of Metals:

Strongly electropositive metals:

Metals having standard reduction potential near about -2.0 volt or more negative like alkali metals, alkaline earth metals are strongly electropositive in nature.

Moderately electropositive metals:

Metals having values of standard reduction potentials between 0.0 and about -2.0 volt are moderately electropositive. Al, Zn, Fe, Ni, Co, etc., belong to this group.

Weakly electropositive metals:

The metals which are above hydrogen and possess positive values of standard reduction potentials are weakly electropositive metals. Cu, Hg, Ag, etc., belong to this group.

To Find Thermal Stability of Metallic Oxides:

The thermal stability of the metal oxide depends on its electropositive nature. As the electropositivity increases from top to bottom, the thermal stability of the oxide also increases from top to bottom.

The oxides of metals having high positive reduction potentials are not stable towards heat. The metals which are above copper form unstable oxides, i.e., these are decomposed on heating.

To Determine the Products of Electrolysis:

In case two or more types of positive and negative ions are present in solution, during electrolysis certain ions are discharged or liberated at the electrodes in preference to others.

In general, in such competition, the ion which is the stronger oxidizing agent (higher value of standard reduction potential) is discharged first at the cathode. e.g. In a mixture of copper and silver ions, silver will be deposited first because the reduction potential of silver is higher than copper.

• The increasing order of deposition of few cations is: K⁺, Ca⁺⁺, Na⁺, Mg⁺⁺+, Al⁺⁺⁺, Zn⁺⁺, Fe⁺⁺, H⁺, Cu⁺⁺, Ag⁺, Au⁺⁺⁺.
• The anion which is a stronger reducing agent (low value of standard reduction potential) is liberated first at the anode.
• The increasing order of discharge of few anions is SO₄^– –, NO₃^–, OH^–, Cl^–, Br^–, I^–
• When an aqueous solution of NaCl containing Na⁺, Cl^–, H^+, and OH- ions is electrolyzed, H+ ions are discharged at cathode and Cl- ions at the anode, i.e., H2 is liberated at cathode and Cl2 at the anode.
• When an aqueous solution of CuS04 containing Cu++, H+ and OH- ions is electrolyzed, Cu⁺⁺ ions are dis­charged at the cathode and OH^– ions at the anode.

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Science > Chemistry > Electrochemistry > Electrochemical Series

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