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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The conductance of an ion depends on its size in an aqueous medium or in the solvent. Bigger is the ionic size lesser is its conductance

Example: The order of size of hydrated ionic radii of alkali metal cations is as Li⁺_(aq) < Na⁺_(aq) < K⁺_(aq)< Rb⁺_(aq)< Cs⁺_(aq). Hence the ease of ionic conductance is Li⁺_(aq) > Na⁺_(aq) > K⁺_(aq) > Rb⁺_(aq) > Cs⁺_(aq)

Concept of Molar Conductivity of an Electrolyte (Λ):

The different solutions may have different concentrations and hence contain a different number of ions. Hence electrolytic conductivity is not a suitable quantity to compare conductance of different solutions. In 1880 the German physicist George Kohlrausch introduced the concept of molar conductivity which is used to compare conductance of different solutions.

The molar conductivity of an electrolyte is defined as the electrolytic conductivity divided by the molar concentration C of the dissolved electrolyte.

Λ = κ / C    or   Λ  = κV

S.I. unit of electrolytic conductivity is siemens per metre (Sm^-1) or S cm^-1. S.I. unit of molar conductivity is siemens square metre per mole (S m² mol^-1). or S cm² mol^-1

If concentration C is measured in M i.e. mol L^-1 or mol dm^-3, then the relationship can be written as

If normality of solution is given then the conductivity is called equivalent conductivity and the relation can be written as

The relation between molar conductivity and equivalent conductivity is

Λ _M =   n Λ_E

Where n is total positive or negative valencies.

Variation of Electrolytic Conductivity with Concentration:

The electrolytic conductivity depends on the number of ions present in a unit volume of a solution. on dilution the degree of dissociation increases. Thus the number of current-carrying ions in the solution increases. But actually, the number of current-carrying ions per unit volume decreases. Hence the activity of the number of ions decreases and hence the electrolytic conductivity also decreases.

For the strong electrolyte, the electrolytic conductivity increases sharply with increasing concentration. For the weak electrolyte, the electrolytic conductivity is very low in dilute solutions and increases much more gradually with increase in the concentration. and this increase is due to an increase in active ions in the solution.

Variation of Molar Conductivity with Concentration:

The molar conductivity of both strong and weak electrolytes increases with dilution i.e. decrease in the concentration.

The molar conductivity is the conductance of all the ions produced by one mole of the electrolyte. Due to an increase in dilution degree of dissociation increases and which results in an increase in the molar conductivity.

For the strong electrolyte, the molar conductivity increases sharply with increasing concentration. Similarly weak electrolyte the molar conductivity increases gradually with an increase in the concentration.

Friedrich Kohlrausch Relation:

Friedrich Kohlrausch performed repeated experiments and plotted a graph of molar conductivity versus the square root of the concentration of a solution.

They showed that the molar conductivity of strong electrolytes varies linearly with the square root of concentration and established the following relation

Where Λ = Molar conductivity at given concentration
Λ_o = Molar conductivity at zero concentration or infinite dilution
C = Concentration of solution
α = constant.

The graph of molar conductivity versus the square root of the concentration of a solution is linear for a strong electrolyte. But such a graph for weak electrolytes is not a straight line.

Kohlrausch Law:

The law states that at infinite dilution, each ion migrates independently of its co-ion and makes its own contribution to the total molar-conductivity of an electrolyte. irrespective of the nature of the other ion with which it is associated.

Thus according to the law at infinite dilution, the total molar conductivity is the algebraic sum of molar conductivities of cation and anion.

Where, Λ = Molar conductivity of a solution
λ ₊^o = Molar conductivity of a cation
λ _–^o = Molar conductivity of an anion

For electrolyte A_mB_n, the molar conductivity at infinite dilution is

Illustration:

In both the cases the difference in of K and Na salt is the difference between Λ_o values of K and Na ions, and it is constant. This illustrates the law.

Applications of Kohlrausch Law:

• The law can be used to calculate the molar-conductivity of any electrolyte at zero concentration.
• The law is particularly useful in the calculation of Λ_o of weak electrolyte for which extrapolation method is not useful.
• Using the extrapolation method value of Λo for strong electrolytes is found and using that value of Λ_o weak electrolyte can be calculated.

Calculation of the Molar Conductivity of any Electrolyte at Zero Concentration:

Let us calculate Λ_o for weak electrolyte acetic acid (CH₃COOH|) using Λ_o values of strong electrolytes sodium acetate (CH₃COONa|) and sodium chloride (NaCl).

The values of  Λ_o for strong electrolytes can be found by extrapolation method and using them for weak electrolyte Λ_o  can be calculated.

Relation Between Molar Conductivity and Dissociation Constant (Theory of Weak Electrolyte) :

Where α = degree of dissociation
Λ = Molar conductivity at concentration C

Λ_o  = Molar conductivity at zero concentration

Now, the dissociation constant k for weak electrolyte is given by

This is the relation between dissociation constant and molar conductivity of the weak electrolyte. This relation is called Ostwald’s equation.

Measurement of Conductivity:

The determination of conductivity and molar conductivity of a solution consists of a measurement of the resistance of the solution using Wheatstone’s metre bridge.

The cell used for measurement consists of a glass tube with two platinum plates coated with a thin layer of finely divided platinum called platinum black. The cell is to be dipped in a solution whose resistance is to be measured as shown in fig.

Now conductivity of a cell is given by

The quantity l/a  is constant and called cell constant and is defined as the ratio of the distance between the electrodes and the area of cross-section of the electrode. It is denoted by ‘b’

The resistance of the solution is found using Wheatstone’s metre bridge. Using the above relation the conductivity of the solution is calculated. The molar conductivity is obtained by using the formula and value of cell constant b can be obtained using the formula b = kR

The circuit arrangement is as shown below.

Types of Conduction:

Metallic Conduction:

The charge transfer through electronic conductors is called metallic conduction

Characteristics of metallic conduction:

• In this conduction, charge transfer occurs through metal.
• It involves the flow of electrons.
• There is no movement of metal atoms.
• There is no chemical change of metal.

Ionic or Electrolytic Conduction:

The charge transfer through electrolytic conductors is called electrolytic conduction

Characteristics of metallic conduction:

• In this conduction, charge transfer occurs through molten electrolyte or its aqueous solution
• It involves the motion of ions in the solution.
• There is a movement of ions.
• There is a chemical change in an electrolyte.

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Science > Chemistry > Electrochemistry > Ionic Conduction

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In this article, we shall study the concept of electrochemistry, its cause, and its terminology.

Electrochemistry is a branch of chemistry which deals with the interrelationship between chemical energy and electrical energy. The study of electrochemistry is broadly divided into two branches. a) Conversion of chemical energy into electrical energy and b) Conversion of electrical energy into chemical energy. Electrochemistry has wide applications in engineering and science. Michael Faraday is called the father of electrochemistry.

Basic Terminology of Electrochemistry:

Oxidation Reaction:

The loss of an electron or electrons by a species is called oxidation.  Example

Na    →    Na⁺ +    e^–

In oxidation, the oxidation number of elements increases as a result of the loss of electrons. In the above example, the oxidation number of sodium increases from 0 to +1. Thus the oxidation can also be defined as the process in which the oxidation number of an element increase.

Reduction Reaction:

The gain of an electron by a species is called reduction.  Example.

Cl    +     e^–   →     Cl^–

In reduction, the oxidation number of an element decreases as a result of the gain of electrons. In the above example the oxidation number of chlorine decreases from 0 to -1. Thus the reduction can also be defined as the process in which the oxidation number of an element decreases.

Oxidizing Agent (Oxidant):

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

Reducing Agent (Reductant):

The substance which loses electrons and makes the other substance to accept electrons is called a reducing agent or reductant. Consider reaction

2Mg     +   O₂    →   2MgO

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

Redox Reactions:

In any of a chemical reaction, if one of the reactants is oxidized, the other is surely reduced. Consider reaction

2Mg     +   O₂    →   2MgO

In this reaction, Mg is oxidized to MgO (loss of electrons by Mg), whereas oxygen is reduced to MgO (gain of electrons by oxygen). Hence 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). Such a reaction may be expressed as the sum of two half-reactions. One reaction involving loss of electrons by a species and another involving gain of electrons by a species. This is the basis of all electrochemical processes.

Conductance:

The substances that allow the flow of electricity through them are called conductors.  The flow of electricity through a conductor involves the transfer of electrons from one point to the other. Depending on the mechanism of the transfer of electrons, the conductors are classified into two types.  a)  Electronic conductors b) Electrolytic conductors

Electronic Conductors:

The conductors through which the conduction of electricity occurs by direct flow of electrons under the influence of applied potential are known as electronic conductors. e.g. copper, aluminium, silver, mercury etc.

Characteristics of Electronic Conductors:

• In electronic conductors, the flow of electricity occurs by the migration of electrons through the conductor.
• In electronic conductors, the conduction does not involve the transfer of matter.
• In electronic conductors, the conduction process does not involve chemical change.
• The resistance of electronic conductors increases and their conductivity decreases with the increase in temperature.
• Ohm’s law is followed but Faraday’s laws are not followed.

Electrolytic Conductors:

The conductors through which the conduction of electricity occurs by the migration of positive and negative ions under the influence of applied potential are known as electrolytic conductors. e.g. electrolysis of fused NaCl.

Characteristics of Electrolytic Conductors:

• In electronic conductors, the flow of electricity occurs by the migration of positive and negative ions through the conductor.
• In electrolytic conductors, the conduction involves the transfer of matter.
• In electrolytic conductors, the conduction process always involves chemical change.
• The resistance of electronic conductors decreases and their conductivity increases with the increase in temperature.
• Both Ohm’s law and Faraday’s laws are followed.

Resistance and Conductance:

By ohm’s law, V = IR

Where R = Resistance of a conductor V = Potential difference across the conductor I = current through the conductor. S. I. unit of resistance is ohm (Ω), that of potential difference is volt (V) and that of current is ampere (A)

Resistance of an Electronic Conducting Wire:

Experimentally it is found that the value of resistance (R) depends on the length (L) of a conductor, the area of cross-section (A) of conductor and nature of a conductor as follows:-

The resistance is directly proportional to the length of a conductor.

R α  L   ………………. (1)

The resistance is inversely proportional to the area of a cross-section.

R α  1/A   ………………. (2)

The resistance depends on the nature of the conductor.

From equation (1) & (2)

R = ρl / A

This is an expression for the specific resistance or the resistivity of a material of a conductor.

Resistivity or Specific Resistance:

We have, ρ = RA / l

Let A = 1 unit  and L = 1 unit, then ρ = R

Thus specific resistance or resistivity of a material of a conductor is defined as that resistance of a conductor whose area of cross-section and its length is unity.

Unit of Resistivity or Specific Resistance:

We have, ρ = RA / l

Hence unit of ρ  = Unit of R x Unit of Area / Unit of Length = ohm x metre² / metre   = ohm metre

Therefore, S.I. unit of resistivity or specific resistance is ohm metre (Ωm)

Conductance:

Reciprocal of resistance is called conductance (K). Its S.I. unit is mho or siemens.

Conductivity:

Reciprocal of resistivity is called as conductivity (κ)

Conductance of an Electronic Conducting Wire:

Experimentally it is found that 1. The conductance (G) is directly proportional to the area of cross-section (A) of a conductor

G α  A   ………………. (1)

The conductance (G) is inversely proportional to the length (L) of the conductor  

G α  1/L     ………..  (2)

It also depends on the material of the conductor.

From equation (1) & (2)

G α  A / L

G =  κA / L

k is constant called specific conductance or conductivity.

This is an expression for the resistance of a conducting wire.

Now,         Let A = 1 unit  and L = 1 unit , then G = κ

Thus specific conductance or conductivity of a material of a conductor is defined as that conductance of a conductor whose area of cross-section and its length is unity. S.I. unit of conductivity is siemens per metre (S m-1). Other units used are S cm-1, W-1m-1, and W^-1cm^-1.

Next Topic: Ionic Conduction

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