Electrochemistry : Electrochemical Cells

Electrochemical or Galvanic cell

Electrochemical cell or Galvanic cell is a device in which a spontaneous redox reaction is used to convert chemical energy into electrical energy i.e. electricity can be obtained with the help of oxidation and reduction reaction”.

(1)     Characteristics of electrochemical cell :  Following are the important characteristics of electrochemical cell,

(i)     Electrochemical cell consists of two vessels, two electrodes, two electrolytic solutions and a salt bridge.

(ii)    The two electrodes taken are made of different materials and usually set up in two separate vessels.

(iii)   The electrolytes are taken in the two different vessels called as half – cells.

(iv)   The two vessels are connected by a salt bridge/porous pot.

(v)    The electrode on which oxidation takes place is called the anode (or – ve pole) and the electrode on which reduction takes place is called the cathode (or + ve pole).

(vi)   In electrochemical cell, ions are discharged only on the cathode.

(vii)  Like electrolytic cell, in electrochemical cell, from outside the electrolytes electrons flow from anode to cathode and current flow from cathode to anode.

(viii) For electrochemical cell,

            Ecell = +ve, DG = −ve.

(ix)   In a electrochemical cell, cell reaction is exothermic.

(2)     Salt bridge and its significance

(i)     Salt bridge is U– shaped glass tube filled with a gelly like substance, agar – agar (plant gel) mixed with an electrolyte like KCl, KNO3, NH4NO3 etc.

(ii)    The electrolytes of the two half-cells should be inert and should not react chemically with each other.

(iii)   The cation as well as anion of the electrolyte should have same ionic mobility and almost same transport number, viz. KCl, KNO3, NH4NO3 etc.

(iv)   The following are the functions of the salt bridge,

(a)     It connects the solutions of two half – cells and completes the cell circuit.

(b)     It prevent transference or diffusion of the solutions fromone half cell to the other.

(c)      It keeps the solution of two half – cells electrically neutral.

(d)     It prevents liquid – liquid junction potential i.e. the potential difference which arises between two solutions when they contact with each other.

Note : Salt bridge can be replaced by a porous partition which allows the migration of ions without intermixing of solution.  

 KCl aq) cannot be used as a salt bridge for the cell, Cu(s)|CuSO4(aq)| |AgNO3(aq)|Ag(s)..

Because AgCl is precipitated as follows,  \[AgN{{O}_{3}}+KCl\xrightarrow{{}}\underset{(\text{ppt}\text{.})}{\mathop{AgCl}}\,\downarrow +KN{{O}_{3}}

 

(3)     Representation of an electrochemical cell

(i)     The interfaces across which a potential difference exists are shown by a semicolon (;) or a single vertical line (|). For example, the two half- cells of the following electrochemical cell can be represented as follows,

          Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) ; Zn ; Zn2+ or

          Zn | Zn2+ and Cu2+; Cu or Cu2+ | Cu

These indicate that potential difference exists at the Zn and Zn2+ ions interface, and similarly at the Cu2+ and Cu interface. Sometimes coma or plus signs is observed in the formulation of half-cells. For example,

Ag. AgCl | Cl or Ag + AgCl | Cl.  These indicate that Agand AgCl together constitute the electrode.

(ii)    The contact between two solutions by means of a salt bridge is indicated by double vertical line (||) between them e.g.; Cu2+|| Zn2+

(iii)   The anode half-cell (or the oxidation half cell) is always written on the left hand side and the cathode half cell (or the reduction half-cell) on the right hand side, with the respective metal electrode on the outside extremes, e.g., Zn ; Zn2+ || Cu2+ ; Cu or   Zn|Zn2+ ||Cu2+|Cu. Sometimes negative and positive sings are put on the electrodes to show that they are negative (anode) and positive (cathode) electrodes.

                              e.g.   (–) Zn ; Zn2+ ||Cu2+ ; Cu(+)

(iv)   An arrow when drawn below the cell formulation gives the direction of the current inside the cell while an arrow drawn above the formulation gives the direction of the electrons flown in the outer circuit.

                             e.g.,   Zn ; Zn2+ ||Cu2+ ; Cu            

(v)    The potential difference between the electrodes (i.e., the cell potential or EMF or emf) is stated in volts along with the temperature at which it is applicable.  For example, E25°C = 1.130 volt or E298K =1.130 volt

With the usual above conventions, the emf of the cell will have a positive value. However, when the cell is formulated in the reverse order, the emf will have a negative value. In other words, the negative value of emf indicates that the oxidation process expected on the left hand electrode will not occur spontaneously and in case oxidation must be made to occur on the left hand electrode, an emf of a value of somewhat larger then that of the specified cell potential will be required from an external source.

(vi)   The concentration of solutions, pressure of gases and physical state of solids and liquids involved, are indicated in the cell formation. For example, Pt, H2(0.9  atm); H+(a = 0.1) || Cu2+ (a = 0.1) ; Cu       

Note Sometimes we get confused in the nomenclature of electrodes. As a memory aid keep in mind the alphabetical order of the first e.g. A (anode) comes before C (cathode). The cell may be written by arranging each of the pair left – right, anode – cathode, oxidation – reduction, negative and positive in the alphabetical order as,

(4)     Reversible and irreversible cells : A cell is said to be reversible if the following two conditions are fulfilled

(i)     The chemical reaction of the cell stops when an exactly equal external emf is applied.

(ii)    The chemical reaction of the cell is reversed and the current flows in opposite direction when the external emf is slightly higher than that of the cell. Any other cell, which does not obey the above two conditions, is termed as irreversible. Daniell cell is reversible but Zn |H2SO4| Ag cell is irreversible in nature

(5)     Types of electrochemical cells : Two main types of electrochemical cells have been reported, these are,

(i)     Chemical cells : The cells in which electrical energy is produced from the energy change accompanying a chemical reaction or a physical process are known as chemical cells. Chemical cells are of two types,

(a)     Chemical cells without transference : In this type of chemical cells, the liquid junction potential is neglected or the transference number is not taken into consideration. In these cells, one electrode is reversible to cations while the other is reversible to the anions of the electrolyte.

(b)     Chemical cells with transference : In this type of chemical cells, the liquid-liquid junction potential or diffusion potential is developed across the boundary between the two solutions. This potential develops due to the difference in mobilities of +ne and −ne ions of the electrolytes.

(6)     Concentration cells : “A cell in which electrical energy is produced by the transference of a substance from a system of high concentration to one at low concentration is known as concentration cells”. Concentration cells are of two types.            

(i)     Electrode concentration cells : In these cells, the potential difference is developed between two electrodes at different concentrations dipped in the same solution of the electrolyte. For example, two hydrogen electrodes at different gaseous pressures in the same solution of hydrogen ions constitute a cell of this type.

\[\frac{Pt,\,\,\,{{H}_{2}}(\text{pressure}\,\,{{p}_{1}})}{\text{Anode}}\,\,|{{H}^{+}}|\,\,\,\frac{{{H}_{2}}\,\,\text{(pressure}\,\,{{p}_{2}})\,\,Pt}{\text{Cathode}}

\[{{E}_{\text{cell}}}=\frac{0.0591}{2}\log \frac{({{p}_{1}})}{({{p}_{2}})} at 25°C If p1 > p2, oxidation occurs at L. H. S. electrode and reduction occurs at R. H. S. electrode.

In the amalgam cells, two amalgams of the same metal at two different concentrations are immersed in the same electrolytic solution. M(HgC1) |Mn+ | Zn(Hg C2)

The emf of the cell is given by the expression, \[{{E}_{\text{cell}}}=\frac{0.0591}{n}\log \frac{{{C}_{1}}}{{{C}_{2}}}   at 25°C

(ii)    Electrolyte concentration cells : In these cells, electrodes are identical but these are immersed in solutions of the same electrolyte of different concentrations. The source of electrical energy in the cell is the tendency of the electrolyte to diffuse from a solution of higher concentration to that of lower concentration. With the expiry of time, the two concentrations tend to become equal. Thus, at the start the emf of the cell is maximum and it gradually falls to zero. Such a cell is represented in the following manner (C2 is greater then C1).

M|M+n(C1)||Mn+(C2)|M or  \[\frac{Zn|Z{{n}^{2+}}({{C}_{1}})}{\text{Anode}}\,\,||\,\,\frac{Z{{n}^{2+}}({{C}_{2}})|Zn}{\text{Cathode}}     

The emf of the cell is given by the following expression,

\[{{E}_{\text{cell}}}=\frac{0.0591}{n}\log \frac{{{C}_{2(R.H.S)}}}{{{C}_{1(L.H.S.)}}}e\ at 25o C

The concentration cells are used to determine the solubility of sparingly soluble salts, valency of the cation of the electrolyte and transition point of the two allotropic forms of metal used as electrodes, etc.

Note :      In concentration cell net redox change is zero and the decrease in free energy during transfer of substance from one concentration to other is responsible for production of electrical energy.

 

(7)     Heat of reaction in an electrochemical cell :  Let n Faraday charge  flows out of a cell of emf E, then

−DG = nFE                                               …….(i)

Gibbs – Helmholtz equation from thermodynamics may be given as

\[\Delta G=\Delta H+T{{\left( \frac{\partial \Delta G}{\partial T} \right)}_{P}}                                                     …….(ii)

From equation (i)and (ii) we get,

\[-nFE=\Delta H+T{{\left[ \frac{\partial (-nFE)}{\partial T} \right]}_{P}}=\Delta H-nFT{{\left( \frac{\partial E}{\partial T} \right)}_{P}} ; \[\Delta H=-nFE+nFT{{\left( \frac{\partial E}{\partial T} \right)}_{P}}

Where \[{{\left( \frac{\partial E}{\partial T} \right)}_{P}} = Temperature coefficient of cell

Case I: When \[{{\left( \frac{\partial E}{\partial T} \right)}_{P}}\ = 0, then DH = −nFE

Case II: When \[\left( \frac{\partial E}{\partial T} \right)\ > 0, then nFE > DH, i.e. process inside the cell is endothermic.    

Case III: When \[\left( \frac{\partial E}{\partial T} \right)\ < 0, then nFE < DH, i.e., process inside the cell is exothermic.  

 

Some Commercial cells (Batteries).

One of the main use of galvanic cells is the generation of portable electrical energy. These cells are also popularly known as batteries. The term battery is generally used for two or more Galvanic cells connected in series. Thus, a battery is an arrangement of electrochemical cells used as an energy source. The basis of an electrochemical cell is an oxidation – reduction reaction. However, for practical purposes there are some limitations to the use of redox reactions. A useful battery should also fulfil the following requirements; 

It should be light and compact so that it can be easily transported.

It should have reasonably long life both when it is being used and when it is not used.

The voltage of the battery should not vary appreciably during its use.

Types of commercial cells : There are mainly two types of commercial cells,

  • Primary cells : In these cells, the electrode reactions cannot be reversed by an external electric energy source. In these cells, reactions occur only once and after use they become dead. Therefore, they are not chargeable. Some common example are, dry cell, mercury cell, Daniell cell and alkaline dry cell.

 

(i) Voltaic cell

 

Cathode : Cu  rod

Anode : Zn rod

Electrolyte : dil. H2SO4

Emf : 1.08 V

At cathode : Cu2+ + 2e → Cu

At Anode : Zn + Zn2+ + 2e

Over all reaction :Zn + Cu2+ → Zn2+ + Cu

 

(ii)  Daniall cell

Cathode : Cu rodAnode : Zn rod

Electrolyte : dil H2SO4

Emf : 1.1 V

At Cathode : Cu2+ + 2e → Cu

At Anode : Zn → Zn2+ + 2e

Over all reaction : Zn + Cu2+ → Zn2+ + Cu

(iii) Lechlanche cell (Dry cell)

 

Cathode : Graphite rodAnode : Zn pot

Electrolyte : Paste of NH4Cl + ZnCl2

Emf : 1.2 V to 1.5 V

At Cathode : \[NH_{4}^{+}+Mn{{O}_{2}}+2{{e}^{-}}\to MnO{{(OH)}^{-}}+N{{H}_{3}}

At Anode : Zn → Zn2+ + 2e

Over all reaction : \[Zn+NH_{4}^{+}+Mn{{O}_{2}}\to Z{{n}^{2+}}+MnO{{(OH)}^{-}}+N{{H}_{3}}

(iv) Mercury cell

 

 

 

Cathode : Mercury (II) oxide

Anode : Zn rod

Electrolyte : Paste of KOH + ZnO

Emf : 1.35 V

At Cathode : \[Hg{{O}_{(s)}}+{{H}_{2}}{{O}_{(l)}}+2{{e}^{-}}\to H{{g}_{(l)}}+2OH_{(aq)}^{-}

At Anode : \[\underset{(\text{amalgam})}{\mathop{Z{{n}_{(s)}}}}\,+20H_{(aq)}^{-}\to Zn{{O}_{(s)}}+{{H}_{2}}{{O}_{(l)}}+2{{e}^{-}}

Over all reaction : 

Zn(s) + HgO(s) → ZnO(s) + Hg(l)

 

Note    In a dry cell ZnCl2 combines with NH3 produced to form the complex [Zn(NH3)2Cl2], otherwise the pressure developed due to NH3 would crack the seal of the cell.

Mercury cell give a constant voltage throughout its life because the electrolyte KOH is not consumed in the reaction.

(2)     Secondary cells: In the secondary cells, the reactions can be reversed by an external electrical energy source. Therefore, these cells can be recharged by passing electric current and used again and again. These are also celled storage cells. Examples of secondary cells are, lead storage battery and nickel – cadmium storage cell.

 

In charged Lead storage cell Alkali cell
 

 

 

 

Positive electrode Perforated lead plates coated with PbO2 Perforated steel plate coated with Ni(OH)4
Negative electrode Perforated lead plates coated with pure lead Perforated steel plate coated with Fe
Electrolyte dil. H2SO4 20% solution of KOH + 1% LiOH
During charging Chemical reaction

At cathode : PbSO4 + 2H+ + 2e– → Pb + H2SO4

At anode :PbSO4 + SO4– – + 2H2O – 2e–  → PbO2 + 2H2SO4

Specific gravity of H2SO4 increases and when specific gravity becomes 1.25 the cell is fully charged.

Emf of cell: When cell is fully charged then E = 2.2 volt

Chemical reaction

At cathode : Ni (OH)2 + 2OH+ – 2e– → Ni(OH)4

At anode :  Fe(OH)2 + 2K+ + 2e–  → Fe + 2KOH

Emf of cell : When cell is fully charged then E = 1.36 volt

 

 

During discharging

 

Chemical reaction

At cathodePb + SO4– – – 2ePbSO4

At anode : PbO2 + 2H+ – 2e + H2SO4  → PbSO4 + 2H2O

Specific gravity of H2SO4 decreases and when specific gravity falls below 1.18 the cell requires recharging.

Emf of cell : When emf of cell falls below 1.9 volt the cell requires recharging.

 

Chemical reaction

At cathode : Fe + 2OH – 2eFe(OH)2

At anode : Ni(OH)4 + 2K+ + 2e– → Ni(OH)2 + 2KOH

Emf of cell : When emf of cell falls below 1.1 V it requires charging.

Efficiency 80% 60%


Fuel cells.

These are Voltaic cells in which the reactants are continuously supplied to the electrodes. These are designed to convert the energy from the combustion of fuels such as H2, CO, CH4, etc. directly into electrical energy. The common example is hydrogen-oxygen fuel cell as described below, 

In this cell, hydrogen and oxygen are bubbled through a porous carbon electrode into concentrated aqueous sodium hydroxide or potassium hydroxide. Hydrogen (the fuel) is fed into the anode compartment where it is oxidised. The oxygen is fed into cathode compartment where it is reduced. The diffusion rates of the gases into the cell are carefully regulated to get maximum efficiency. The net reaction is the same as burning of hydrogen and oxygen to form water. The reactions are

At anode :  2[H2(g) + 2OH(aq) → 2H2O(l) + 2e

At cathode : O2(g) + 2H2O(l) + 4e → 4OH(aq)

Overall reaction : 2H2(g) + O2(g) → 2H2O(l)

Each electrode is made of porous compressed carbon containing a small amount of catalyst (Pt. Ag or CoO). This cell runs continuously as long as the reactants are fed. These fuel cells are more efficient than conventionally used methods of generating electricity on a large scale by burning hydrogen, carbon, fuels because these fuel cells convert the energy of the fuel directly into electricity. This cell has been used for electric power in the Apollo space program. Fuel cells offer great promises for energy conversion in future. The important advantages of fuel cells over ordinary batteries are

(1)     High efficiency : The fuel cells convert the energy of a fuel directly into electricity and therefore, they are more efficient than the conventional methods of generating electricity on a large scale by burning hydrogen, carbon fuels. Though we expect 100 % efficiency in fuel cells, so far 60 – 70% efficiency has been attained. The conventional methods of production of electrical energy involve combustion of a fuel to liberate heat which is then used to produce electricity. The efficiency of these methods is only about 40%.

(2)     Continuous source of energy : There is no electrode material to be replaced as in ordinary battery. The fuel can be fed continuously to produce power. For this reason, H2 – O2 fuel cells have been used in space crafts.

(3)     Pollution free working : There are no objectionable byproducts and, therefore, they do not cause pollution problems. Since fuel cells are efficient and free from pollution, attempts are being made to get better commercially practical fuel cells.

 

Thermodynamic characteristics of some Fuel cells

Fuel cells Cell reaction ΔGo (kJ mol−1) ΔHo (kJ mol−1) E (volt)
H2 – O2 2H2 + O2 → 2H2O – 237.2 – 258.9 1.23
C – O2 C + O2 → CO2 – 137.3 – 110.5 0.71
CH4 – O2 CH4 + 2O2 → CO2 + H2O – 818.0 – 890.4 1.060
CH3OH – O2 CH3OH + 3/2O2 → CO2 + 2H2O – 706.9 – 764.0 1.22