`color{green}("Electrochemistry") :` It is the study of production of electricity from energy released during spontaneous chemical reactions and the use of electrical energy to bring about non-spontaneous chemical transformations.

`=>` A large number of metals, sodium hydroxide, chlorine, fluorine and many other chemicals are produced by electrochemical methods.

`=>` Batteries and fuel cells convert chemical energy into electrical energy and are used on a large scale in various instruments and devices.

`=>` The reactions carried out electrochemically can be energy efficient and less polluting.

`=>` The transmission of sensory signals through cells to brain and vice versa and communication between the cells are known to have electrochemical origin.

In Class XI (Unit 8) we had studied the construction and functioning of Daniell cell (Fig).

`=>` This cell converts the chemical energy liberated during the redox reaction

`color{red}(Zn(s) +Cu^(2+)(aq) → Zn^(2+) (aq) + Cu (s))`.

to electrical energy and has an electrical potential equal to `1.1 V` when concentration of `Zn^(2+)` and `Cu^(2+)` ions is unity `(1 mol dm^(-3))^(**)`. Such a device is called a galvanic or a voltaic cell.

`=>` If an external opposite potential is applied and increased slowly, we find that the reaction continues to take place till the opposing voltage reaches the value 1.1 V when, the reaction stops altogether and no current flows through the cell.

`=>`Any further increase in the external potential again starts the reaction but in the opposite direction. It now functions as an electrolytic cell, a device for using electrical energy to carry non-spontaneous chemical reactions. Both types of cells are quite important.

`color{green}("Definition") :` A galvanic cell is an electrochemical cell that converts the chemical energy of a spontaneous redox reaction into electrical energy.

`=>` In this device the Gibbs energy of the spontaneous redox reaction is converted into electrical work which may be used for running a motor or other electrical gadgets like heater, fan, geyser, etc.

Daniell cell is one such cell in which the following redox reaction occurs.
`Zn-CuSO_4` cell= It consists of two beakers containing `ZnSO_4` solution in one and `CuSO_4` in the other. A `Zn`-rod is immersed in `ZnSO_4` solution and `Cu`-rod in `CuSO_4`. These are known as electrodes. The two electrodes are connected with the help of `Cu`-wire through an ammeter(voltmeter). The two solutions are connected internally by the salt bridge.When circuit is completed by inserting key the voltmeter shows a deflection indicating the flow of electric current through the circuit. The following observations are made:
a)Zn rod gradually loses weight and the concentration of `Zn^(2+)` in solution increases. It is because the Zn-electrode loses electrons and changes into Zn^(2+) i.e. oxidation takes place at zinc,hence termed as anode `color{red}(Zn(s) → Zn^(2+) +2e^-) `
b) The electrons from the Zn rod flows to the Cu-rod through the copper wire and get captured by Cu^(2+) ions of the solution. They change into Cu-metal and get deposited on Cu-electrode.Since reduction occurs here,hence termed as cathode. `color{red}(Cu^(2+) +2e^(-) → Cu(s))`
So the flow of electrons takes place from Zn-rod to Cu-rod and the electric current starts flowing. Conventionally the flow of electric current is taken opposite to the flow of electrons. Hence current flows from copper to zinc.

`color{red}(Zn(s) +Cu^(2+) (aq) → Zn^(2+) (aq) + Cu(s))`.

`color{red}("This reaction is a combination of two half reactions whose addition gives the overall cell reaction :")`

(i) `color{red}(Cu^(2+) +2e^(-) → Cu(s))` (reduction half reaction) .......................(1).
(ii) `color{red}(Zn(s) → Zn^(2+) +2e^(-) )` (oxidation half reaction) ......................(2).

`=>` These reactions occur in two different portions of the Daniell cell.

`=>` The reduction half reaction occurs on the copper electrode while the oxidation half reaction occurs on the zinc electrode.

`=>` These two portions of the cell are also called half-cells or redox couples.

`=>` The copper electrode is called the reduction half cell and the zinc electrode, the oxidation half-cell.

`=>` We can construct innumerable number of galvanic cells on the pattern of Daniell cell by taking combinations of different half-cells.

`=>` Each half-cell consists of a metallic electrode dipped into an electrolyte.

`=>` The two half-cells are connected by a metallic wire through a voltmeter and a switch externally.

`=>` The electrolytes of the two half-cells are connected internally through a salt bridge as shown. Sometimes, both the electrodes dip in the same electrolyte solution and in such cases we don’t require a salt bridge.

`=>` At each electrode-electrolyte interface there is a tendency of metal ions from the solution to deposit on the metal electrode trying to make it positively charged.

`=>` At the same time, metal atoms of the electrode have a tendency to go into the solution as ions and leave behind the electrons at the electrode trying to make it negatively charged.

`=>` At equilibrium, there is a separation of charges and depending on the tendencies of the two opposing reactions, the electrode may be positively or negatively charged with respect to the solution.

`=>` A potential difference develops between the electrode and the electrolyte which is called electrode potential.

`=>` When the concentrations of all the species involved in a half-cell is unity then the electrode potential is known as standard electrode potential.

`=>` According to IUPAC convention, standard reduction potentials are now called standard electrode potentials.

`=>` In a galvanic cell, the half-cell in which oxidation takes place is called anode and it has a negative potential with respect to the solution.

`=>` The other half-cell in which reduction takes place is called cathode and it has a positive potential with respect to the solution.

`=>` So, there exists a potential difference between the two electrodes and as soon as the switch is in the on position the electrons flow from negative electrode to positive electrode. The direction of current flow is opposite to that of electron flow.

`=>` The potential difference between the two electrodes of a galvanic cell is called the cell potential and is measured in volts.

`=>` The cell potential is the difference between the electrode potentials (reduction potentials) of the cathode and anode. It is called the cell electromotive force (emf) of the cell when no current is drawn through the cell.

`=>` We keep the anode on the left and the cathode on the right while representing the galvanic cell.

`color{green}("Representation of Galvanic Cell") :` A galvanic cell is generally represented by putting a vertical line between metal and electrolyte solution and putting a double vertical line between the two electrolytes connected by a salt bridge. Under this convention the emf of the cell is positive and is given by the potential of the half-cell on the right hand side minus the potential of the half-cell on the left hand side i.e.

`color{red}(E_text(cell) = E_text(right) - E_text(left))`

`color{red}("This is illustrated by the following example:")`

`color{green}("Cell reaction :")` `color{red}(Cu(s) +2Ag^(+) (aq) → Cu^(2+) (aq) +2 Ag (s))` ............................(3).

`color{green}("Half-cell reactions :")`

`color{green}("Cathode (reduction)")` : `color{red}(2Ag^+ (aq) +2e^(-) → 2Ag(s))` ...............................(4).

`color{green}("Anode (oxidation)")` : `color{red}(Cu(s) → Cu^(2+) (aq) +2e^(-))` .......................................(5).

e. The cell can be represented as :

`color{red}(Cu(s) | Cu^(2+) (aq) || Ag^(+) (aq) | Ag(s))`

And we have `color{red}(E_text(cell) = E_text(right) - E_text(left) = E_(Ag^+ | Ag) - E_(Cu^(2+) | Cu))` ..................(6).

The potential of individual half-cell cannot be measured. We can measure only the difference between the two half-cell potentials that gives the emf of the cell. It can be measured by using some electrode as the reference electrode. If we arbitrarily choose the potential of one electrode (half- cell) then that of the other can be determined with respect to this.

`color{green}("Standard Hydrogen Electrode :")`

According to convention, a half-cell called standard hydrogen electrode represented by

`color{red}(Pt(s) | H_2(g) | H^(+) (aq))`,

is assigned a zero potential at all temperatures corresponding to the reaction `color{red}(H^(+) (aq) +e^(-) → 1/2 H_2 (g))`.

`=>` The standard hydrogen electrode consists of a platinum electrode coated with platinum black.

`=>` The electrode is dipped in an acidic solution and pure hydrogen gas is bubbled through it.

`=>` The concentration of both the reduced and oxidised forms of hydrogen is maintained at unity . This implies that the pressure of hydrogen gas is one bar and the concentration of hydrogen ion in the solution is one molar.

At 298 K the emf of the cell, standard hydrogen electrode | | second half-cell constructed by taking standard hydrogen
electrode as anode (reference half-cell) and the other half-cell as cathode, gives the reduction potential of the other half-cell. If
the concentrations of the oxidised and the reduced forms of the species in the right hand half-cell are unity, then the cell potential is equal to standard electrode potential `color{red}(E_R^(⊖) - E_L^(⊖))` As `color{red}(E_L^(⊖))` for standard hydrogen electrode is zero.

`color{red}(E^(⊖) = E_R^(⊖) - 0 = E_R^(⊖))`

`color{green}("The measured emf of the cell")` : `color{red}(Pt(s) | H_2 (g , 1 text(bar) | H^(+) (aq , 1 M) | | Cu^(2+) (aq , 1 M) | Cu)`

is `0.34 V` and it is also the value for the standard electrode potential of the half-cell corresponding to the reaction :

`color{red}(Cu^(2+) (aq , 1 M) +2 e^(-) → Cu(s))`.

`color{green}("Similarly, the measured emf of the cell :")` `color{red}(Pt(s) | H_2 (g , 1 text(bar) ) | H^(+) (aq , 1M) | | Zn^(2+) (aq , 1 M) | Zn)` is -0.76 V corresponding to the standard electrode potential of the half-cell reaction :

`color{red}(Zn^(2+) (aq , 1M) +2e^(-) → Zn(s))`.

`=>` The positive value of the standard electrode potential in the first case indicates that `Cu^(2+)` ions get reduced more easily than `H^+` ions. The reverse process cannot occur, that is, hydrogen ions cannot oxidise `Cu` (or alternatively we can say that hydrogen gas can reduce copper ion) under the standard conditions described above. Thus, `Cu` does not dissolve in `HCl`. In nitric acid it is oxidised by nitrate ion and not by hydrogen ion.

`=>` The negative value of the standard electrode potential in the second case indicates that hydrogen ions can oxidise zinc (or zinc can reduce hydrogen ions).

In view of this convention, the half reaction for the Daniell cell in Fig can be written as :

`color{green}("Left electrode")` : `color{red}(Zn (s) → Zn^(2+) (aq , 1M) +2e^(-))`.

`color{green}("Right electrode")` : `color{red}(Cu^(2+) (aq , 1M) +2 e^(-) → Cu(s))`.

The overall reaction of the cell is the sum of above two reactions and we obtain the equation :

`color{red}(Zn(s) +Cu^(2+) (aq) → Zn^(2+) (aq) + Cu(s))`

Emf of the cell `color{red}( = E_text(cell)^0 = E_R^0 - E_L^0)`

`color{red}( = 0.34V - (-0.76) V = 1.10 V)`

`color{green}("Note") :` Sometimes metals like platinum or gold are used as inert electrodes. They do not participate in the reaction but provide their surface for oxidation or reduction reactions and for the conduction of electrons.

For example, `Pt` is used in the following half-cells :

`color{green}("Hydrogen electrode")` : `color{red}(Pt(s) | H_2 (g) | H^(+) (aq))`.

`color{green}("With half-cell reaction")` : `color{red}(H^(+) (aq) + e^(-) → 1/2 H_2 (g))`.

`color{green}("Bromine electrode ")` : `color{red}(Pt(s) | Br_2(aq) | Br^(-) (aq))`.

`color{green}("With half-cell reaction")` : `color{red}(1/2 Br_2(aq) + e^(-) → Br^(-) (aq))`.

The standard electrode potentials are very important and we can extract a lot of useful information from them. The values of standard electrode potentials for some selected half-cell reduction reactions are given in Table.

`=>` If the standard electrode potential of an electrode is greater than zero then its reduced form is more stable compared to hydrogen gas.

`=>` If the standard electrode potential is negative then hydrogen gas is more stable than the reduced form of the species.

`=>` It can be seen that the standard electrode potential for fluorine is the highest in the Table indicating that fluorine gas `(F_2)` has the maximum tendency to get reduced to fluoride ions `(F^–)` and therefore fluorine gas is the strongest oxidising agent and fluoride ion is the weakest reducing agent.

`=>` Lithium has the lowest electrode potential indicating that lithium ion is the weakest oxidising agent while lithium metal is the most
powerful reducing agent in an aqueous solution.

`=>` It may be seen that as we go from top to bottom in Table 3.1 the standard electrode potential decreases and with this, decreases the oxidising power of the species on the left and increases the reducing power of the species on the right hand side of the reaction.

`=>` Electrochemical cells are extensively used for determining the pH of solutions, solubility product, equilibrium constant and other thermodynamic properties and for potentiometric titrations.

The standard reduction potentials of a large number of electrodes have been measured by using the SHE. The arrangement of elements in order of increasing or decreasing order of standard reduction potential values is called the electrochemical series or activity series.
`=>` Applications of electrochemical series:-
1 To find the relative oxidizing and reducing electrodes: In a cell the electrode with high value of reduction potential will be reduced and act as cathode, while the electrode with high value of oxidation potential will be oxidized and act as anode.
2 To calculate the EMF of cell:- The standard e.m.f. of cell is calculated as:
STANDARD EMF OF CELL=[STD. OXIDATION POTENTIAL OF OXIDATION HALF CELL]+[STD. REDUCTION POTENTIAL OF REDUCTION HALF
CELL]
=[STD. REDUCTION POT. OF REDUCTION HALF CELL]-[STD. REDUCTION POT. OF OXIDATION HALF CELL]
3 To predict whether a metal give `H_2` gas with acid:- If the emf comes out to be positive the metal will liberate hydrogen gas otherwise
not.
4 To predict the spontaneity of a reaction:- If emf comes out to be positive the reaction will occur and if comes out to be negative the reaction will not occur.