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ChemistryEdu Logo Electrochemistry | Batteries and Corrosion#

Battery#

Any battery or cell that we use as a source of electrical energy is basically a galvanic cell where the chemical energy is converted into electrical energy.

There are two types of batteries:

  • Primary batteries
  • Secondary batteries

Primary Batteries#

In primary batteries, the reaction occurs only once and after use over a period of time, battery becomes dead and cannot be used again.

Examples of primary batteries are:

  • Dry cell or Leclanche cell
  • Mercury cell

Dry cell or Leclanche cell#

  • It consists of zinc container that also acts as anode and the cathode is a carbon (graphite) rod surrounded by powdered manganese dioxide and carbon.

  • The space between the electrodes is filled by a moist paste of ammonium chloride (NH4Cl) and zinc chloride (ZnCl2). The electrode reactions are complex but they can be written approximately as:

\[Anode:\ Zn_{(s)} → Zn^{2+} + 2e^-\]
\[Cathode:\ MnO_2 + NH_4^+ + e^- → MnO(OH) + NH_3\]

  • In the reaction at cathode, manganese is reduced from +4 oxidation state to +3 oxidation state.

  • Ammonia produced in the reaction forms a complex with Zn2+ to give [Zn(NH3)4]4+.

  • The cell has a potential of nearly 1.5 V.

Mercury Cell#

  • Mercury cell, suitable for low current devices like hearing aids, watches etc consists of zinc in mercury amalgam as anode and a paste of HgO and carbon as cathode.

  • The electrolyte is a paste of KOH and ZnO. The electrode reactions for the cell are given below:

\[Anode:\ Zn(Hg) + 2OH^- → ZnO + H_2O + 2e^-\]
\[Cathode:\ HgO + H_2O + 2e^- → Hg_{(l)} + 2OH^-\]
  • The overall reaction is represented by:
\[Zn(Hg) + HgO_{(s)} → ZnO_{(s)} + Hg_{(l)}\]
  • The cell potential is approximately 1.35 V and remains constant during its life as the overall reaction does not involve any ion in solution whose concentration can change during its lifetime.

Secondary Batteries#

A secondary cell after use can be recharged by passing current through it in the opposite direction so that it can be used again.

Examples of secondary batteries are:

  • Lead storage battery
  • Nickel-Cadmium cell

Lead storage battery#

  • It consists of a lead anode and a grid of lead packed with lead dioxide (PbO2) as cathode.

  • A 38% solution of sulphuric acid is used as an electrolyte.

  • The cell reactions are given as:

\[Anode:\ Pb_{(s)} + {SO_4^{2-}}_{(aq)} → PbSO_{4(s)} + 2e^-\]
\[Cathode:\ PbO_{2(s)} + {SO_4^{2-}}_{(aq)} + 4H^+_{(aq)} + 2e^- → PbSO_{4(s)} + 2H_2O_{(l)}\]
  • The overall cell reaction is given by:
\[Pb_{(s)} + PbO_{2(s)} + 2H_2SO_{4(aq)} → 2PbSO_4{(s)} + 2H_2O_{(l)}\]
  • On charging the battery, the reaction is reversed and PbSO4 on anode and cathode is converted into Pb and PbO2 respectively.

Nickel-Cadmium cell#

  • It has longer life than the lead storage battery but it is more expensive to manufacture.

  • The overall reaction during discharge is:

\[Cd_{(s)} + 2Ni(OH)_{3(s)} → CdO_{(s)} + 2Ni(OH)_{2(s)} + H_2O_{(l)}\]

Fuel Cells#

Galvanic cells that are designed to convert the energy of combustion of fuels like hydrogen, methane, methanol etc directly into electrical energy are called fuel cells.

Fuel cell using H2 and O2 produces electricity#

  • In this cell, hydrogen and oxygen are bubbled through porous carbon electrodes into concentrated aqueous sodium hydroxide solution.

  • Catalysts like finely divided platinum or palladium metal are incorporated into the electrodes for increasing the rate of electrode reactions.

  • The cathode and anode reactions are given below:

\[Cathode:\ O_{2(g)} + 2H_2O_{(l)} + 4e^- → 4OH^-_{(aq)}\]
\[Anode:\ 2H_2O_{(g)} + 4OH^{-}_{(aq)} → 4H_2O_{(l)} + 4e^-\]
  • The overall reaction is given by:
\[2H_{2(g)} + O_{2(g)} → 2H_2O_{(l)}\]

Advantages of fuel cells#

  • Fuel cells produce electricity with an efficiency of about 60% compared to thermal plants whose efficiency is about 40%.
  • Fuel cells are pollution free.

Suggest two materials other than hydrogen that can be used as fuels in fuel cells.

Methane (CH4) and methanol (CH3OH) are the two materials other than hydrogen that can be used as fuels in fuel cells.

Corrosion#

  • Corrosion slowly coats the surfaces of metallic objects with oxides or other salts of the metal.
  • The rusting of iron, tarnishing of silver, development of green coating on copper and bronze are some of the examples of corrosion.
  • It causes enormous damage to buildings, bridges, ships and to all objects made of metal especially that of iron.
  • In corrosion, a metal is oxidized by loss of electron to oxygen and formation of oxides.
  • Corrosion of iron (rusting) occurs in the presence of water and air.

Corrosion is an electrochemical phenomenon#

  • At a particular spot of an object made of iron, oxidation takes place and that spot behaves as anode.
\[Anode:\ 2Fe_{(s)} → 2Fe^+_{(aq)} + 4e^-,\ E^0_{(Fe^{2+}/Fe)} = -0.44V\]
  • Electrons released at anodic spot move through the metal and go to another spot on the metal and reduce oxygen in the presence of H+. This spot behaves as cathode with the reaction:
\[Cathode:\ O_2{(g)} + 4H^+_{(aq)} + 4e^- → 2H_2O_{(l)},\ E^0_{H^+/O_2/H_2O} = 1.23V\]
  • The overall reaction is:
\[2Fe_{(s)} + O_{2((g)} + 4H^+_{(aq)} → 2Fe^+_{(aq)} + 2H_2O,\ E^0_{cell} = 1.67V\]
  • The ferrous ions are further oxidized by atmospheric oxygen to ferric ions which come out as rust in the form of hydrated ferric oxide (Fe2O3.xH2O) and with further production of hydrogen ions.

Prevention of Corrosion#

  • By painting and greasing.
  • By coating the surface by some inert metals (Sn, Zn etc).
  • By alloying.