Common Cell

Common cell:

(i) Leclanche dry cell:

Construction:

Zinc (Using anode filled with a mixture of ammonium chloride and MnO₂, filled with a graphite cathode in the center.

Cell reaction:

At cathode:

$2NH{{}_4^{+}}_{(aq)}+2e^{-}\to 2N{H}_{3(g)}+H_{2(g)}$

These gases are removed by depolariser MnO₂.

$2Mn{O}_{2(s)}+H_{2(g)}\to M{n}_2{O}_{3(s)}+H_2{O}_{(\ell )}$

Ammonia is absorbed by zinc ions.

$N{H}_{3(g)}+Z{n^{++}}_{(aq)}+2C{l^{-}}_{(aq)}\to Zn{(N{H}_3)}_{2}C{l}_{2(s)}$

At anode:

$Z{n}_{(s)}\to Z{n^{++}}_{(aq)}+2e^{-}$

Cell reaction:

$2Mn{O_2}_{(s)}+2N{H}_{4}C{l}_{(s)}+Z{n}_{(s)}\to M{n}_2{O}_{3(s)}+H_2{O}_{(\ell )}+Zn{(N{H}_3)}_{2}C{l}_{2(s)}$

Voltage $=1.5V$

Disadvantage:

(i) Rapid withdrawl of current causes gases to collect at cathode and voltage suddenly drops.

(ii) Slow but spontaneous reaction of zinc with ammonium ions results in poor shell life.

Alkaline Battery:

No gases are formed, hence it is an improvement over a previous model.

Cell reaction:

At cathode:

$2Mn{O}_{2(s)}+H_2{O}_{(\ell )}+2e^{-}\to M{n}_2{O}_{3(s)}+2O{H^{-}}_{(aq)}$

At anode:

$Z{n}_{(s)}+2O{H^{-}}_{(aq)}\to Zn{O}_{(s)}+H_2{O}_{(\ell )}+2e^{-}$

Voltage = 1.54 V

Buffer cells:

Construction:

Anode is metallic zinc while cathode is HgO. Both in the form at tightly compacted powders separated by a moist paper of HgO containing some NaOH / KOH. Moistened paper serves as salt bridge.

Cell reaction:

At cathode:

$Hg{O}_{(s)}+2H_2{O}_{(\ell )}+2e^{-}\to H{g}_{(\ell )}+2O{H}^{-}$

At anode:

$Z{n}_{(s)}+2O{H^{-}}_{(aq)}\to Zn{O}_{(s)}+H_2{O}_{(\ell )}+2e^{-}$

Storage Batteries: (discharging)

A rechargeable array of many cells.

(i) Lead storage Battery:

Construction: Alternate plates of porous lead and lead deposited with PbO₂. Electrolyte is aqueous H₂SO₄.

Cell reaction:

At anode:

$P{b}_{(s)}+SO{{}_4^{--}}_{(aq)}\to PbS{O}_{4(s)}+2e^{-}$

At cathode:

$Pb{O}_{2(s)}+4H_3{O^{+}}_{(aq)}+SO{{}_4^{--}}_{(aq)}+2e^{-}\to PbS{O}_{4(s)}+6H_2{O}_{(\ell )}$

Overall reaction:

$P{b}_{(s)}+Pb{O}_{2(s)}+2H_{2}S{O}_{4(aq)}\to 2PbS{O}_{4(s)}+2H_2{O}_{(\ell )}$

Cell reaction (charging): For this cell is connected to an external source of current.

At anode: $PbS{O}_4+2e^{-}\to Pb+S{O}_4^{--}$

(From electrical source)

At cathode: $\underset{\left(s\right)}{PbS{O}_4}+\underset{\left(\ell \right)}{6H_{2}O}\to Pb{O}_2+4H_3{O}^{+}+S{O}_4^{--}+2e^{-}$

Over all reaction:

$2\underset{\left(s\right)}{PbS{O}_4}+\underset{\left(\ell \right)}{2H_{2}O}\to Pb\left(s\right)+Pb{O}_2\left(s\right)+2H_{2}S{O}_4\left(aq\right)$

Reactions for discharging are enthalpy opposite to that of charging.

Density of solution decreasing (due to consumption of H₂SO₄) when it is discharged and density increases (due to formation of H₂SO₄) when it is charged.

Fuel cells

The main disadvantages of a primary cell is that it can deliver current for a short period of time due to limited amounts of oxidising agents.

A fuel cell is a device which converts chemical energy into electrical energy as long as the outside supply fuel is maintained and example is H₂ - O₂ fuel cell.

The over all reactions are

$Anode{H}_2+2O{H}^{\_}\left(aq\right)\to 2H_{2}O\left(\ell \right)+2e^{-}$

$Cathode{O}_2+2H_{2}O\left(\ell \right)+4e^{-}\to 4O{H}^{-}\left(aq\right)$

$Overall2H_2\left(g\right)+O_2\left(g\right)\to 2H_{2}O\left(\ell \right)$

The fuel cells are efficient (upto 95%) and are pollution free.

The main disadvantage of fuel cell is high cost of catalyst and difficult storage of H₂ and O₂ gas.

Corrosion:

The conversion of a metal into undesirable compounds by the action of air, moisture is called corrosion, In case of iron corrosion it is called rusting.

Corrosion is a redox reaction by which metal gets oxidised by air in presence of moisture. Small cathodic and anodic cells are set up on metal. The area of metal in contact with water acts as anode. Metal loses electrons to form cations.

$Fe\to F{e}^{2+}+2e^{-}$

The electrons flow along the metal into the cathode area. Where O2 is reduced to hydroxyl ions in presence of H2O.

Cathode reaction $O_2+2H_{2}O+4e^{-}\to 4O{H}^{-}$

The over all reaction is

$2Fe+O_2+2H_{2}O\to 2Fe{\left(OH\right)}_2$

Fe(OH)₂ may dehydrate to form its oxide FeO, which than further gets oxidised to Fe₂O₃. Which is further hydrated to form rust?

$F{e}_2{O}_3+n{H}_{2}O\to F{e}_2{O}_3.n{H}_{2}O\left(rust\right)$

Corrosion of a metal weakness the metal and hence is undesirable. The following methods are used to protect a metal from corrosion.

(i) Applying a protective film such as metal oxide.

(ii) Plating the metal with a more electropositive metal, which loses electrons in preference to the less electropositive metal e.g. Zn is coated on steel. (Galvanization)

(iii) By connecting the metal with more electropositive metal (Electrical protection)