1. Electrochemical cells

Chapter 3
electrochemistry

Introduction

Batteries are everywhere in modern societies. They provide the electric current to start our autombiles and to power a host of products such as pocket caculators, digital watches, heart pacemaker, radio, and tape recorders.

Electrochemistry is the area of chemistry concerned with the interconversion of chemical and electrical.A battery is a an electrochemical cell, a device for interconverting chemical and electrical energy. A battery takes the energy relased by a spontaneous chemical reaction and uses it to produce electricity.

Electrochemical cell

It is device for converting chemical energy in to electrical energy.

The two types of cells are therefore reverse of each other   

Construction/ Working principle   

When ever an metal strip is put in an electrolyte the process of oxidation and reduction takes place simultaneously within the system. Due to this there is a potential difference between the metal phase and the liquid phase.

On joining the metal strips through a wire (of negligible resistence)  the current flows as long as the potential difference exists between the metal phase and the liquid phase.

I   Anode :

Some metals (which are reactive) are found to have tendency to go into the solution phase when these are placed in contact  with their ions or their salt solutions.

For example :  Zn rod is placed in ZnSO₄solution .

The Zn atom or metal atoms will move in the solution to form Zn⁺². After some time following equilibrium will be established.  

Zn(s)   Zn²⁺ +2e^–

There will be accumulation of sufficient negative charge on the rod which will not allow extra zinc ions to move in the solution. i.e. solution will be saturated with Zn⁺² ions. 

The positive charge will be more concentrated nearly the rod.

The extra positive charge of the solution will be more concentrated around the negatively charged rod.  An electrical double layer is developed in the system and hence a potential difference is created between the rod and the solution which is known as electrode potential

This particular electrode is known as anode :

• On anode oxidation will take place. (release of electron).
• To act as source of electrons.
• It is of negative polarity.
• The electrode potential is represented by E_{Zn(s) / Zn²⁺ (aq)} 

II Cathode :

Some metals(Cu, Ag, Au etc.,) are found to have the opposite tendency i.e. when placed in contact with their aqueous ions, the ions from the solution will get deposited on the metal rod.

The following equilibrium will be established :  

Cu²⁺ +^2e–  Cu(s).

So rod will have deficiency of electron (positive charge).Extra negative charge will surround this positively charged rod and form double layer. An electrical double layer is developed in the system and hence a potential difference is created between the rod and the solution which is known as electrode potential. This will be known as cathode.

• At cathode reduction will take place.(gain of e– will take place)
• To act as sink of electron.
• Positive polarity will be developed.
• Their electrode potential can be represented by : E_{Cu²⁺(aq)/Cu(s)}

Anode :

Cathode :

Construction of Cell :

It has  two half–cells,each having a beaker containing a metal strip that dips in its aqueous solution.

The metal strips are  called electrodes and are connected by an conducting wire.

Two solutions are connected by a salt bridge.

The oxidation and reduction half reactions occur at a separate electrodes and electric current flows   through the wire.

Selection of electrolyte for Salt Bridge :

The electrolyte in salt bridge should be such that speed of it's cation equals speed of it's anion in electrical field.

For that charge and sign of the ions should be almost equal.

Transport number of cation = Transport number of anion

                or

Mobility of cation = Mobility of anion

KCl is generally preffered but KNO₃ or NH₄NO₃ can also be used.

If Ag⁺, Hg²⁺, Pb²⁺, Tl⁺ ions are present in a cell then in salt bridge KCl is not used because there can be formation of  precipitate of AgCl, Hg₂Cl₂, PbCl₂ or TlCl at mouth of tube which will prevent the migration of ions and its functioning will stop.

Functions of Salt Bridge :

A salt bridge is a U–shaped inverted tube that contains a gel permeated with an inert electrolyte.

It connects the solution of two half cell to complete the circuit.

It minimise the liquid junction potential. The potential difference between the junction of two liquids.

It maintains the electhical neutrality of the solution in order to give continious flow or generation of current.

" The simultaneous electrical neutrality of the anodic oxidation chamber and cathodic reduction chamber is due to same mobility or velocity of K⁺ and NO₃^– ions taken into salt bridge.

If the salt bridge is removed then voltage drops to zero.

The ions of the inert electrolyte do not react with other ion in the solution and the ions are not oxidised or reduced at the electrodes.

Generally tube is filled with a paste of agar-agar powder with a natural electrolyte/generally not common to anionic/cathodic compartment with porous plugs at each mouth of tube.

It prevents mechanical mixing of two electrolytic solution.

Electrode Potential :

The driving force that pushes the negative charge electrons away from the anode and pulls them towards the cathode is an electrical potential called electromotive force also known as cell potential or the cell voltage. Its unit is volt 

The potential difference devepoled between metal electrode and its ions in solution in known as electrode potential.

Electrode potential depends upon :

• Concentration of the solution.
• Nature of the metal.
• Nature of the electrolyte.
• Pressure temperature coditions.

The potential difference developed between metal electrodes and the solution of its ions at 1 M concentration at 1 bar pressure and 298 K is known as standard electrode potential.

Reference electrode  :

The potential of a singal electode cannot be determined what were the potential difference between two electrodes can be accurately measured using a reference electrode.

An electrode is chosen as a reference with respect to which all other electrodes are valued. 

Standard Hydrogen Electrode (SHE) is taken as standard reference electrode. Its electrode potential is arbitrarily assumed to be 0.00 volt.    

Standard Hydrogen Electrode (SHE)  consists of a platinum electrode in contact with H2 gas and aqueous H+ ions at standard state conditions (1 atm H₂ gas, 1 M H⁺ (aq), 25°C).

  2H+ (aq, 1M) + 2e^– →  H₂ (g, 1 atm)  [E° = 0V]

H₂(g, 1atm) →  2H⁺ (aq, 1M) + 2e^–    [E° = 0V]

Cell potential :  

The difference in electrode potentials of the two half cell reactions (oxidation half cell and reduction half cell) is known as emf of the cell or cell potential.

The emf of the cell or cell potential can be calculated from the values of electrode potential of the two half cell constituning the cell. The following three methode are in use :

When oxidation potential of anode and reduction potential of cathode are taken into account :

E°_cell = oxidation potential of anode + reduction potential of cathode

E°_ox (anode) + E°_red(cathode)

When reduction potential of both electrodes are taken into account :

E°_cell = Reduction potential of cathode – Reduction potential of anode

 = E°_cathode – E°_anode  C both are reduction potential.

When oxidation potential of both electrodes are taken into account :

 E°_cell = oxidation potential of anode – Oxidation potential of cathode

  = E°_ox (anode) – E°_ox (cathode)

The standard cell potential E° is the cell potential when both reactants and products are in their standard  states – solutes at 1 M concentration, gases at a potential pressure of 1 atm, solids and liquids in pure from, with all at a specified temperature, usually 25° C. 

E°_cell is intensive property so on multiplying/Dividing cell reaction reaction by any number, the E°_cell value would not change.

Free energy changes for cell reaction :

The free energy changeDG (a thermochemical quantity) and the cell potential E(an electrochemical quantity) both measure the  driving force of a chemical reaction.

The values ofDG and E are directly proportional and are related by the equation.

 DG = –nFE

where 

n = Number of moles of electron transfered in the reaction.

 F = Faraday constant = 96485 C/mole e^–  96500  C/mole e^–

Shorthand Notation for Galvanic Cells

We require two half cells to produce an electrochemical cell, which can be represented by follwing few rules;

The anode half-cell is always written on the left followed on the right by cathode half cell.

The separation of two phases (state of matter) is shown by a vertical line.

The various materials present in the same phase are shown together using commas.

The salt bridge is represented by a double slash (||).

The significant features of the substance viz. pressure of a gas, concentration of ions etc. are indicated in brackets immediately after writing the substance.

For a gas electrode, the gas is indicated after the electrode for anode and before the electrode in case of cathode. (i.e  Pt H₂ / H⁺ or H⁺ /H₂ Pt)

For SHE reference potential is taken to be zero at all temperature.

SOP = – SRP = 0 for SHE.

To calculate standard potential of any other electrode a cell is coupled with standard hydrogen electrode (SHE) and it's potential is measured that gives the value of electrode potential of that electrode.

Anode  :  Zinc electrode                                                

Cathode :  SHE 

Cell :   Zinc electrode || SHE                  

Cell potential  :

Ecell =  – E°_Zn²⁺/Zn      

  = 0.76 V (at 298 K experimentaly)                   

So, E⁰_Zn²⁺/Zn = – 0.76 V (SRP) 

 E⁰ _Zn/Zn²⁺(aq) = 0.76 V(SOP)

So, w.r.t. H2, Zn has greater

tendency to get oxidised. In similar manner reduction potentials (SRP) at 298 K for many other electrodes are calculated and are arranged in a series increasing order known as electro chemical series. 

Electrochemical Series :

Calculation of Electrode Potential of unknown electrode with the help of given (two) electrode.

Obtain the reaction of the 3rd electrode with the help of some algebraic operations on reactions of the given electrodes.

Then calculate DG of the 3rd reaction with the help of some algebaric operations of DG⁰ of 1st and 2nd reactions.

Use  DG⁰ = –nF E⁰_elec. to calculate unknown  E.P.

is intensive property so if we multiply/Devide electrode reaction by any number the  value would not changed

i.e.    Zn²⁺ + 2^e– ®  Zn(s)   [E° = – 0.76 V[

Multiply by 2

2Zn²⁺ + 4e^– ®  2Zn(s)  [E° = – 0.76 V (remain same)]