Electrochemistry

Arrhenius Theory of Ionization 

(Theory of Electrolytic Dissociation)

Ionization

Ionization is the process of dissociation or splitting up of an electrolyte into its constituting positive and negative ions in aqueous solution or in fused state. 

 

Introduction                                                              A Swedish Chemist, Arrhenius in 1884 presented Theory of Electrolytic Dissociation to account for mechanism of conductivity of electrolytes, behaviour of electrolytic dissociation and phenomenon of electrolysis. [It gives not only a simple mechanism of conductivity of electrolytes but also a clear explanation of deviation of electrolytes observed in dilute solution].

 Main Postulates

The main points of this theory are summarized as follows:

1. When electrolytes (acids, bases and salts) are dissolved in water, they yield two types of charged particles called cations (positive particles) and anions (negative particles) which are free to move in solution. [The number of oppositely charged ions are equal so that the solution as a whole is neutral].e.g.

                                    AB_(aq)        ¾¾¾¾®   A⁺_(aq)     +    B^–_(aq)   

                                    HCl_(aq)      ¾¾¾¾®   H⁺_(aq)     +  Cl^–_(aq) 

                                    NaOH_(aq)   ¾¾¾¾®   Na⁺_(aq)   +  OH^–_(aq) 

                                    NaCl_(aq)     ¾¾¾¾®   Na⁺_(aq)   +  Cl^–_(aq) 

 

2. The oppositely charged ions present in solution constantly recombine on colliding to form unionized neutral molecules and thus there is a state of equilibrium between ionized and non-ionized molecules i.e.

                                    AB_(aq)   ⇌  A⁺_(aq)     +    B^–_(aq)   

      e.g.           

                                    NaCl_(aq) ⇌ Na⁺_(aq)   +   Cl^–_(aq) 

Thus ionization is a reversible process to which Law of Mass

 Action can be applied:

Where;

K    =    Ionization constant                   ,         

[  ]  =    Molar conc. of ions or unionized molecules.

3.  When an electric current is passed through an electrolytic solution, cations move towards the   cathode and anions move towards anode. This movement of ions is called Electrolysis. [Although this term is wrong in actual meanings because ionization has already been occurred and is not due to electric current. Thus electric current does not cause ionization. The current only directs or shifts the ions towards their respective electrodes]. This movement of ions is responsible for the conduction of electricity through electrolytic solutions. These ions lose their charges on electrodes and change into neutral atoms or molecules by the gain or loss of electrons.

4.  The properties of an electrolyte in solution are the properties of the ions produced in the solution. [In aqueous solution, acidic properties are due to H⁺ ions, basic properties are due to OH^– ions and the properties of salts are the properties of ions].

5. The electrical conductivity (i.e. ability to conduct electricity which is reciprocal to resistivity) of an  electrolytic solution depends upon the number of ions present in the solution which in turn is proportional to the degree of dissociation, at a particular concentration and temperature.

6.  The degree of dissociation is the extent to which the substance ionizes or the fraction of the total   number of molecules present in the ionic form.

It depends upon:

(i)         Nature of electrolyte.    

(ii)        Nature of solvent. 

(iii)       Degree of dilution    

(iv)       Temperature.

Strong electrolytes (e.g. HCl, NaCl etc) dissociates completely in water as compared to weak electrolytes (e.g. CH₃COOH, NH₄OH) which ionizes to lesser extent.

The behaviour of electrolyte differs in different solvent, thus HCl is dissociated in water due to high dielectric constant (80) of water.

The more dilute the solution is, the greater will be degree of dissociation of an electrolyte as more dilute solution produces greater number of ions.

    The greater the temperature, the more is the degree of dissociation.

                   Electrolysis                           

Definition of Electrolysis

Electrolysis is the process of chemical decomposition of an electrolyte in aqueous solution or molten (fused) state into its constituent ions and their deposition as neutral species at their respective electrodes by the passage of electric current with all the chemical process involved in it.

The process of the movement of cation and anion of an electrolyte in fused state or in its aqueous solution and their conversion into neutral species (atoms or molecules) at their respective electrodes under the influence of an applied electric field is called electrolysis.

OR

The process of decomposition of an electrolyte by the passage of electricity through its molten state or aqueous solution form in an electrolytic cell is called electrolysis.

Electrolysis is the process in which electrical energy is used to bring about a non-spontaneous chemical change.

Explanation

1.  Electrolysis is an endothermic chemical change (decomposition) caused by the passage of an electric current through fused ionic compounds or their aqueous solutions. [Note that in the electrolytic conduction, the electrolyte is decomposed].

2.  During electrolysis, metals are deposited on a cathode or hydrogen gas is released while non-metals (gases) are liberated at anode or metal making anode dissolves.

  

Conditions for Electrolysis

1)     Electrolyte should be in molten state or in solution form.

2)     Solution should be diluted, so degree of dissociation is greater.

3)    The high temperature favours the electrolysis.

 

Factors affecting products of electrolysis

The products of electrolysis mainly depends upon

(i)   nature of electrolyte   (strong or weak)

(ii)Nature of electrode (Inert or non-attacked or reactive electrodes)

(iii)  Concentration of ions

(iv)  Over potential of ions

Inert or non-attacked Electrodes

These electrodes do not react with solution of electrolyte or with the

 product of electrolysis. e.g. C (graphite), Pt

 

reactive electrodes

They are not used in electrolysis.

 

Preferential Discharge of Ions

When there are more than one cation or anion the process of

 discharge becomes competitive in nature. Discharge of any ion

 requires energy and in case of several ions being present the

 discharge of that ion will take place first which requires the energy.

 

Uses and Applications of Electrolysis

The process of electrolysis is used for following purposes:

1.   For electroplating of metals

2.   For extraction of metals from their ores

3.   For refining of impure metals

 

Uses and Applications of Electrolysis

 

1)   For electroplating of metals

The process of electrolysis is used for electroplating baser metals

 with superior metals to make them more attractive and resistant for

 corrosion. Nickel plating, chromium plating, tin plating, gold plating

 etc. are various types of electroplating.

 

2)    For extraction of metals from their ores

Electrolysis is used for the extraction of certain metals from their

 ores. Mostly representative metals of groups A like Na, K, Ca, Mg,

 Al, etc. are extracted from their ores by electrolysis

 (electrometallurgy). Many normal metals are extracted by the

 electrolysis of their molten compounds usually their chlorides as

 their chlorides have low melting points. e.g. sodium metal is

 extracted by the electrolysis of molten sodium chloride in Down’

 Cell by Down’s process.

 

[Previously (before 1886) aluminium metal is extracted from its salt

 AlCl₃ by heating it with sodium metal. Sodium itself was expensive

 making aluminium even more expensive. Now-a-days Al is

 extracted by the electrolysis of purified alumina (Al₂O₃) which is

 obtained from its chief ore bauxite (Al₂O₃.nH₂O) by purification.

 Due to the process of electrolytic extraction of aluminium, its price

 quickly falls].

3)    For refining of impure metals

Many metals are purified into pure metals by process of electrolysis.

 For example, impure copper (blister copper) is purified by the

 process of electrolysis.

Construction and working of an electrolytic cell

An electrolytic cell consists of electrolyte in a vessel, electrodes and a battery. 

The figure shows that electrons from battery enter through cathode at which positive ions are reduced by accepting electrons. At anode negative ions loses electrons and undergoes oxidation. It means at cathode reduction occurs and oxidation takes place at anode.

       Changes accompanied by Conduction of

Electricity through Electrolytic Solution

Electrolytes & its composition

The compounds which ionize into positive and negative ions in

 aqueous solution conduct electricity are called electrolytes, which

 always contain two types of ions namely Positively charged ion

 called cation, (Na⁺, K⁺, Mg²⁺ etc. ) and Negatively charged ion

 called anion (Cl^–, SO₄^2– etc.). These ions move randomly in their solution.

Construction of Electrolytic Cell

Electrolysis is carried out in a specialized apparatus called

 Electrolytic Cell. Electrolysis is the process in which electrical

 energy is used to bring about a non-spontaneous chemical change.

 An electrolytic cell is composed of two metallic plates called

 Electrodes which are immersed in electrolytic solution and are

 connected to +ive and –ive terminals of a battery. The electrode

 connected with +ive terminal of battery is called as Anode where

 process of oxidation occurs. The electrode connected with –ive

 terminal of battery is called as Cathode where process of reduction

 occurs. [In an electrochemical cell anode is –ive whereas cathode is

 +ive].

1. The electrolytic cell consists of a container having electrodes that

 are connected to the terminals of a battery.

 

2. Electrodes are the conductors (metallic or non-metallic solid rods)

 through which an electric current enters or leaves the solution of

 electrolytes. They are of two types i.e. Anode and Cathode.

3. Anode is the positive electrode by the help of which current enters

the electrolytic solution. It is the  electrode towards which anions are

 attracted and on which process of oxidation takes place.

4.  Cathode is the negative electrode by the help of which current

leaves the electrolytic solution. It is the electrode towards which

cations are attracted and on which process of reduction occur.

Working of Electrolytic Cell or 

Mechanism of Electrolytic Conduction

1. When an electrolyte is fused or dissolved in water, it is ionized

 and its constituents ions are free to move through the molten

 electrolyte or its solution.

 

2.  When two electrodes, connected to the two poles of a battery are

 immersed in the electrolyte, the battery supplies electrons to one

 electrode (cathode).

 

3. The cations of the electrolyte move to cathode (negative electrode)

 where they receive electronsthereby changing into neutral atoms.

 

4.  At the other electrode (anode), the battery withdraws electrons.

 The anions of the electrolyte rush to anode where they transfer their

 electrons thereby changing into neutral atoms.

 

5. It is this movement of ions (in opposite directions), which

 constitutes the electrolytic current.

 

6. the battery provides a source of electrical energy, pushing

 electrons onto the cathode and making it negatively charged

. Electrons are also drawn out of the anode, making it positively

 charged. 

 

Migration of Ions towards respective Electrodes

The ions of an electrolyte move toward their respective electrodes by

 the passage of electric current through its aqueous solution. The

 cations are attracted towards cathode while anions are attracted

 towards anode.

 

Discharge and Deposition of ions into Neutral Species 

at respective Electrodes

On reaching the electrodes, these migrating ions change into neutral

 species by losing or gaining electrons.

 

Redox Reaction

The electrolytic reaction occurring at electrodes involves reduction at

 cathode and oxidation at anode is called Redox Reaction or

 Oxidation-Reduction Reaction (ORR).

 

Reduction

The process that involves the gain of electrons is called Reduction. 

It always takes place at cathode.

 

Oxidation

The process that involves the loss of electrons is called Oxidation. 

It always takes place at anode.

 

Overall Reaction

The overall reaction is obtained by adding two half reactions i.e.

 oxidation and reduction.

Predictions of Products of Electrolysis Using Inert Electrodes

 (Platinum or Graphite)

Predictions of Products of Electrolysis 

Using Active Electrodes

Predictions of Products of Electrolysis 

Electrolytic Reactions or Cell Reactions 

Electrolysis of Aqueous CuCl₂

1.      When cupric chloride; CuCl₂ (an ionic compound) is dissolved in water, it dissociates into its ions i.e. Cu²⁺ and Cl ^–. Before application of current, these ions move randomly in water. But when the current is passed, then these ions start moving towards their respective electrodes.

2.       On passing electricity through the aqueous copper chloride, electrolysis starts with the movement of cations (Cu²⁺) towards cathode and anions (Cl^–) towards anode where they are discharged into copper metal and chlorine gas respectively.

3.       When ions are changed to neutral particles, the current can no longer flow.

 

Cell or Electrolytic Reaction

Electrolysis of Molten Sodium Chloride

1.        Solid sodium chloride does not conduct electricity (because its ions are held together tightly in regular lattice arrangement and are not free to move in the crystal) but when sodium chloride is fused at 800°C, their ions Na⁺ and Cl^– are freed from their lattice moving freely to conduct electricity.

2.      On passing electricity through the molten sodium chloride, electrolysis starts with the movement of cations (Na⁺) towards cathode and anions (Cl^–) towards anode where they are discharged into sodium metal and chlorine gas respectively.

3.       The potential required to oxidize Cl^- ions to Cl₂ is -1.36 volts and the potential needed to reduce Na⁺ ions to sodium metal is -2.71 volts. The battery used to drive this reaction must therefore have a potential of at least 4.07 volts.

Electrolytic Reaction

Electrode Potential/ Half Cell Potential/ Single Electrode Potential/Redox Potential/ Oxidation-reduction potential

Definition of Electrode Potential                                                         

The magnitude of electrode potential of a metal (when it is in contact with a solution of its own ions) is a measure of relative tendency of a metallic electrode to lose or gain electrons or to undergo redox reaction i.e. it is a measure of the relative tendency of the metal to undergo oxidation (loss of electrons) or reduction (gain of electrons).

The potential of the electrodes is also called redox potential or oxidation-reduction potential or half-cell potential.

The difference of potential (difference of electron concentration) set up between a metal and solution of its own salt (i.e. its own ions) in equilibrium is called Electrode Potential of the metal (or electrode) or half-cell potential or single electrode potential denoted as E. By convention, the electrode (redox) potentials are always quoted as Reduction Potential. It is the tendency of a metal to get oxidized or reduced when it is placed in a solution of its own salt. It is expressed in volts.

According to IUPAC convention, the reduction potential alone be called as the electrode potential unless it is specifically mentioned.

When a metal (electrode) is placed in aqueous solution of its own ions in a half cell, it has a tendency to lose or gain electrons and acquires either negative or positive charge with respect to solution and potential is developed between metal and solution which is named by electrode potential of metal. It is the capability or power of a metal electrode to undergo oxidation or reduction. 

Nature of Electrode Potential

Electrode potential is an intrinsic or intensive property and does not depend on the size of metallic rod. Its value does not change when half-cell reaction is multiplied or divided by a co-efficient. It is independent of the amount of species in the reaction

 

Factors affecting Electrode potential

1.         Nature of electrode

2.         Concentration of ions in solution

3.         Temperature

Cell Potential or E°_cell

The force that causes the flow of electrons from one electrode to another electrode is called e.m.f of the cell also called cell potential.

When two electrodes of different electrode potentials are connected in the same circuit then electrons (charge carrier) start experiencing force due to potential difference between electrodes and start moving from high potential to low potential terminal. This force experienced by electrons is named by electromotive force and measured by potential difference. A Galvanic cell consists of an oxidizing agent in one compartment that pulls electrons through a wire from a reducing agent in the other compartment. The ‘pull or driving force’ on the electrons is called the cell potential (E_cell). The unit of electrical potential is the volt which is defined as1 joule work per coulomb of charge transferred.

The strength of this force is equal to the difference of the reduction potential of two electrodes i.e. reduced and oxidized. It is also equal to the sum of the oxidation potential and reduction potential of the electrodes which are oxidized and reduced respectively. 

Oxidation Potential and Reduction Potential

1.         Both oxidation and reduction potentials are equal in magnitude but opposite in sign.

2.         It is not a thermodynamic property, so values of E are not additive.

 

Oxidation potential/ 

Standard Oxidation potential (SOP)

Metal atoms of electrode have tendency to go into the solution as positive ion and leave behind the electrons at the electrode trying make it negatively charged. The electrode potential developed so called oxidation potential.

Oxidation potential is the potential developed at the anode by oxidation. It is a measure of electron losing tendency of metal (to get oxidize). [In Galvanic cell, species or substances having smaller reduction potential is oxidized and act as anode]. Thus, oxidation potential of the half-cell undergoing oxidation (M ⇌  M⁺  +  1e^–) is represented by:

For example; in Zn-Cu Galvanic cell, Zn acts as anode while Cu acts as cathode. Thus Zn oxidizes into Zn²⁺ ions by losing electrons going into the solution surrounding the anode (i.e. ZnSO₄). The oxidation potential of Zn anode is +0.76 V. However, in Mg-Zn Galvanic cell where Zn acts as cathode, its reduction potential is – 0.76 V. 

Reduction Potential/

Standard Reduction Potential (SRP)

Metal ions present in the solution have tendency to deposit on the electrode by accepting electrons from it making the electrode positively charged. The electrode potential developed so called reduction potential.

Reduction Potential is the potential developed at the cathode by reduction. It is a measure of the tendency of metal for reduction to occur. [In Galvanic cell, species or substances having greater reduction potential is reduced and act as cathode].Thus, reduction potential of the half-cell undergoing reduction, (M⁺ +1e^– ⇌ M) is represented by:

For example; in Zn-Cu Galvanic cell, Zn acts as anode while Cu acts as cathode. Thus Cu²⁺ ions reduce at cathode by gaining electrons from the cathode and convert into Cu atoms which deposit on the cathode. The reduction potential of Cu cathode is +0.34 V. However, in Cu-Ag Galvanic cell where Cu acts as anode, its oxidation potential is  – 0.34 V.

According to IUPAC convention, the reduction potential alone be called as the electrode potential unless it is specifically mentioned.

The oxidation potential and the reduction potential of an electrode have the same value but differ in signs [Oxidation potential are obtained by reversing the sign of reduction potential].

Electromotive Force (e.m.f) 

It is the difference between the electrode potentials of two half-cells and cause flow of current from electrode at higher potential to electrode at lower potential. It is also the measure of free energy change. Standard emf of a cell is given by,

Standard Electrode Potential (SEP)

The potential of an electrode (half-cell) determined by comparing it with reference electrode called Standard Hydrogen Electrode (S.H.E.) at standard conditions (at 298 K, 1 atm. pressure and 1M concentration) is called Standard Electrode Potential denoted by E^o. It is the potential difference developed between metal electrode and solution of ions of unit molarity (1M) at 1 atm pressure and 25°C (298 K) is called standard electrode potential.

It is a measure of tendency of a metallic electrode to lose or gain electrons when it is in contact with a solution of its own ions of 1 M concentration at 25^oC and 1 atm. pressure.

Since absolute electrode potential cannot be measured, so the relative electrode potential is determined. The potential of a single electrode can be measured by coupling (comparing) it with a standard half-cell whose electrode (reduction) potential has been assigned a constant value. For this purpose a reference electrode called Standard Hydrogen Electrode (SHE) is used whose potential is arbitrarily chosen as 0.000 volt.

When an electrode is combined with S.H.E., it either undergoes oxidation or reduction. At the electrodes having greater reduction potential than 0.000, reduction takes place and SHE undergoes oxidation and acts as anode. [e.g. SHE acts as anode when it is combined with copper electrode. In this case, oxidation will take place at SHE]. At the electrodes having smaller reduction potential than 0.000, oxidation takes place and SHE undergoes reduction and acts as cathode [e.g. SHE acts as cathode when it is combined with zinc electrode.  In this case reduction takes place at SHE].

Reference Electrode

The electrode of known potential is called reference electrode. It may be primary reference electrode like hydrogen electrode or secondary reference electrode like calomel electrode.

Primary Reference Electrodes

The electrode potential is found out by coupling the electrode with a primary reference electrode, the potential of which is arbitrarily fixed as zero. The important primary reference electrode used is a standard hydrogen electrode or normal hydrogen electrode.

Secondary Reference electrode

The use of SHE is difficult, because it is difficult to maintain 1M H+ ion concentration and the pressure of the gas at one atmosphere. Also the electrode will easily get poisoned in case of traces of impurities in the gas and hence other reference electrode called secondary reference electrodes are used. e.g. Saturated calomel electrode (saturated KCl)

Standard Hydrogen Electrode (SHE)/Normal Hydrogen Electrode (NHE)

In order to determine the electrode potential of an unknown electrode, some standard electrodes are used like

(i)         Standard Electrode

(ii)        Calomel electrode                        

It is impossible to measure the absolute electrode potential of a single electrode, so the relative electrode potential is determined by connecting an electrode (whose electrode potential is to be measured) with a standard half-cell whose electrode (reduction) potential has been assigned a constant value. For this purpose a reference electrode called Standard Hydrogen Half-Cell or Standard Hydrogen Electrode (SHE) is selected whose potential is arbitrary chosen as 0.000 volt for coupling with unknown half-cell.

Standard Hydrogen Electrode (SHE) is used as reference electrode to determine the reduction potential of different electrodes. To define an electrochemical ‘Sea-Level’, chemists have chosen a reference standard half-cell called the standard hydrogen electrode (SHE).

 

An SHE consists of a platinum plate (inert electrode) coated with a layer of finely divided platinum black electrolytically (to give it a large surface area) immersed in one molar (1 M) HCl/H₂SO₄ (H⁺ ions) solution encased in a glass hood (tube) so that H₂ gas at standard state conditions (i.e. 1 atm pressure and 25°C/ 298 K temperature) can bubble over the platinum electrode. The platinum acts as an electrical conductor and also facilitates the attainment of equilibrium between the gas and its ion in solution. [The platinum adsorbs H₂ gas on its surface and the platinum coated with hydrogen behaves as if it was made entirely of hydrogen. An equilibrium is established between adsorbed layer of hydrogen and H⁺ ion in the solution].

Standard hydrogen electrode (SHE) Standard hydrogen electrode (SHE) also known as normal hydrogen electrode (NHE), consists of platinum wire, carrying platinum foil coated with finely divided platinum black. The wire is sealed into a glass tube placed in beaker containing 1 M HCl. The hydrogen gas at 1 atm pressure is bubbled through the solution at 298K. Half-cell is pt H₂ (1 atm) H⁺ (1 M)

The standard electrode potential (E_oxidation or E_reduction) of the SHE is arbitrary taken as zero volt at all temperature. The possible SHE half reactions are:

SHE can acts as cathode as well as anode depends on the other electrode. 

function of platinised platinum

1.   It provides support surface for establishing an equilibrium between H₂ gas and H⁺ ions as soon as possible.

                   H₂ (g)  ⇌   2H⁺(aq)   + 2e^–

2.    It allows H₂ gas to be adsorbed on its surface (the platinum coated with hydrogen behaves as if it were made entirely of hydrogen).

3.    It acts as inert metal connection to H⁺/H₂ system but it does not take part in the chemical reaction (i.e. it has no tendency to form ions itself).

4.         It also provides the necessary electrical connection to carry away or supply electrons.

Drawbacks

Its main drawbacks are

1.  It is difficult to maintain 1 atm pressure of H₂ gas.

2.  It is difficult to maintain H⁺ ion concentration 1 M.

3.  The platinum electrode is easily poisoned by traces of impurities.

Hence, calomel electrodes are conveniently used as reference electrodes, It consists of mercury in contact with Hg₂ Cl₂ (calomel) paste in a solution of KCl.

Measurement of Electrode Potential of Zinc - SHE Cell

Construction of Galvanic Cell

The potential of a single electrode (like Zn) can be measured by connecting it with SHE. For this purpose, a Galvanic cell consisting of two half-cells (each containing an electrode dipped in solution of its own ions) is constructed.

1).  Anode Half Cell;   zinc metal dipped in 1M ZnSO₄ solution.

2).  Cathode Half Cell, SHE (inert Pt electrode immersed in 1M HCl solution in a current of H₂ gas at 1 atm)

3).  Salt bridge;          inert electrolyte (KCl jelly) completes the circuit and prevents mixing of solutions

4).  voltmeter;             Both electrodes are connected with voltmeter giving positive cell potential (E°).

1).  Anode Half Cell

On the left is placed a half cell consisting of zinc metal (whose electrode potential is to be measured) dipped in 1M ZnSO₄ solution. 

2).  Cathode Half Cell

On the right, is arranged SHE consisting of an inert platinum electrode immersed in 1M HCl solution through which H₂ gas is bubbled at 1 atm. pressure. 

3).  Salt bridge

The two half cells are connected by salt bridge made of KCl jelly (inert electrolyte) completes the circuit between two half cells and prevents mixing of solutions. 

4).  voltmeter

Both electrodes are connected with the terminals of a voltmeter in such a way that measured cell potential (E°) cell be positive. [The electrode connected with –ive terminal of voltmeter would be negative electrode i.e. anode whereas the electrode connected with positive terminal of voltmeter would be positive electrode i.e. cathode].

 

Principle of Working

1).  Anode Reaction

As the cell operates, the mass of zinc electrode decreases and concentration of Zn²⁺ ions increase around the zinc electrode. 

2).  Cathode Reaction

The H⁺ ions concentration decreases in SHE half-cell and H₂ is produced. 

3).  Function of Salt bridge

Due to movement of ions (i.e. Zn²⁺ and H⁺ ions), charge imbalance creates in half cells which ceases electrode reactions thereby halting electron flow through the wire. Thus ions of inert electrolyte from the salt bridge migrate to neutralize the charge in the cell compartments. Anions move towards the anode and cations move towards the cathode.

4).  Cell Potential

The voltage or cell potential or E° cell given by voltmeter is found to be +0.76 volt, when zinc electrode is connected with the negative terminal of voltmeter and SHE with positive terminal of voltmeter. [This shows that zinc electrode acts as anode i.e. oxidation takes place at zinc electrode while SHE acts as cathode i.e. at SHE reduction takes place.  Thus zinc electrode is negative with respect to SHE].

Cell Reactions

As the reduction potential is the reverse of oxidation potential, therefore, reduction potential of zinc would be –0.76 volt. The negative sign signifies that actually the reaction at zinc electrode occurs in opposite direction i.e. it is the oxidation rather than reduction which occurs at zinc.

Cell Diagram

In this cell zinc reduces H⁺ ion to H₂ gas and itself oxidizes to Zn²⁺ ions.  Thus flow of electrons takes place from zinc to SHE.

Summary of Zinc-SHE cell

SHE may act as a cathode or anode depends on other electrode. If Zn half-cell is connected with SHE, Zn will oxidize into Zn²⁺ ions thereby acting as anode while SHE get reduced into H₂ gas acting as cathode. The reading in the voltmeter (+0.76 V) will be oxidation potential of the other electrode (i.e. Zn) as the electrode (reduction) potential of SHE is 0.000 V (The reduction potential value is same as oxidation potential but with opposite sign i.e. -0.76 V. The – sign implies that the reaction at Zn electrode is oxidation rather than reduction). The reactions occur are:

Measurement of Electrode Potential of Copper (SHE-Copper Cell)

Construction of Galvanic Cell

The potential of a single electrode (like Cu) can be measured by connecting it with SHE. For this purpose, a Galvanic cell consisting of two half-cells (each containing an electrode dipped in solution of its own ions) is constructed.

1).  Anode Half Cell;    SHE (inert Pt electrode immersed in 1M HCl solution in a current of H₂ gas at 1 atm)

2).  Cathode Half Cell, Copper metal dipped in 1M CuSO₄ solution.

3).  Salt bridge;           inert electrolyte (KCl jelly) completes the circuit and prevents mixing of solutions

4).  voltmeter;             Both electrodes are connected with voltmeter giving positive cell potential (E°).

1).  Anode Half Cell

On the left, is arranged SHE consisting of an inert platinum electrode immersed in 1M HCl solution through which H₂ gas is bubbled at 1 atm. pressure. 

2).  Cathode Half Cell

On the right is placed a half cell consisting of copper metal (whose electrode potential is to be measured) dipped in 1M CuSO₄ solution. 

3).  Salt bridge

The two half cells are connected by salt bridge made of KCl jelly (inert electrolyte) completes the circuit between two half cells and prevents mixing of solutions. 

4).  voltmeter

Both electrodes are connected with the terminals of a voltmeter in such a way that measured cell potential (E°) cell be positive. [The electrode connected with –ive terminal of voltmeter would be negative electrode i.e. anode whereas the electrode connected with positive terminal of voltmeter would be positive electrode i.e. cathode].

Principle of Working

 

1).  Anode Reaction

As the cell operates, the H⁺ ions concentration in SHE half-cell increases and H₂ gas is consumed.

2).  Cathode Reaction

The concentration of Cu²⁺ ions increase around the copper electrode. 

3).  Function of Salt bridge

Due to movement of ions (i.e. Cu²⁺ and H⁺ ions), charge imbalance creates in half-cells which ceases electrode reactions thereby halting electron flow through the wire. Thus ions of inert electrolyte from the salt bridge migrate to neutralize the charge in the cell compartments. Anions move towards the anode and cations move towards the cathode.

4).  Cell Potential

The voltage or cell potential or E° cell given by voltmeter is found to be +0.34 volt, when SHE is connected with the negative terminal of voltmeter and copper electrode with positive terminal of voltmeter. [This shows that SHE acts as anode i.e. oxidation takes place at SHE while copper electrode acts as cathode i.e. at copper electrode reduction takes place.  Thus copper electrode is positive with respect to SHE].

Cell Reactions

As the oxidation potential is the reverse of reduction potential, therefore, oxidation potential of copper would be –0.34 volt. The negative sign signifies that actually the reaction at copper electrode occurs in opposite direction i.e. it is the reduction rather than oxidation which occurs at copper.

Cell Diagram

In this cell, SHE (H₂) reduces Cu²⁺ ion to copper metal and itself oxidizes to H⁺ ions.  Thus flow of electrons takes place from SHE to copper.

Summary of SHE-Cu cell

SHE may act as a cathode or anode depends on other electrode. If Cu half-cell is connected with SHE, Cu²⁺ ions are reduced into Cu thereby Cu acting as cathode while SHE gets oxidized into H⁺ ion acting as anode. The reading in the voltmeter (+0.34 V) will be the reduction potential of the other electrode (i.e. Cu) as the electrode (oxidation) potential of SHE is 0.000 V (The oxidation potential value is same as reduction potential but with opposite sign i.e. – 0.34 V. The – sign implies that the reaction at Cu electrode is reduction rather than oxidation). The reactions occur are:

Electrochemical Series (E.C.S)

 

Definition of E.C.S

The list of substances (elements or ions) arranged in order of increasing standard reduction potential (SRP) is called electrochemical series (E.C.S).

Consequently, the strongest reducing agents are located in the upper right of the table (Li, Cs, K, Ba, Sr, Ca, Na, Mg, Be, Al, Zn etc.) and the strongest oxidizing agents are found in the lower left of the table (F₂, O₃, MnO₄^‒, Cr₂O₇^2‒ etc.). All metals above hydrogen have negative reduction potentials. 

Features of E.C.S

Different elements have different tendencies to lose or gain electrons i.e. to be oxidized or reduced. For determination of standard reduction potential of elements, electrode potential of SHE is used as a reference point.

1.      In the series, all the elements above SHE have negative reduction potential and positive  oxidation potential.

2.      In the series, all the elements below SHE have positive reduction potential and negative   oxidation potential.

3.  Oxidation potential of electrode is numerically equal to reduction potential but with opposite sign.

4.  Metals placed above hydrogen in E.C.S. have lower (negative) reduction potential acts as anode (i.e. they are oxidized) while metals (or non-metals) placed below hydrogen in E.C.S. have greater (positive) reduction potential acts as cathode (i.e. they are reduced).

5.  From top to bottom in E.C.S, the reducing strength decreases and oxidizing strength increases.  Thus Li is the strongest reducing agent and F₂ is the strongest oxidizing agent.

6. Position of metals in E.C.S shows the order of their reactivity. Metal displaces metal lying below them in the series from its solution of their salts.      

Applications of E.C.S

E.C.S is used: