Chemistry > Electrochemistry > 17.0 Battery
Electrochemistry
1.0 Introduction
2.0 Conductors and Non-Conductors
3.0 Electrochemical Cells
4.0 Electrolysis and electrode Reactions
5.0 Electrochemical Cell
6.0 Electrode Potential
7.0 Nature of Electrodes
8.0 IUPAC Cell Representation and Convention
9.0 Standard Cell EMF and Standard Reduction Potential
10.0 Electropositive Character of Metals
11.0 Difference between EMF and potential difference
12.0 Nernst Equation
13.0 Laws of Electrolysis
14.0 Electromotive Force
15.0 Thermodynamics of the Cells
16.0 Concentration Cells
17.0 Battery
18.0 Fuel Cell
17.1 Primary Battery or Voltaic Cells (Dry Cells):
These are non-rechargeable batteries so these cannot be used for a longer period of time.In this cell , once all the chemicals are used then reaction won't occur and hence no electrical energy will be produced. Once it gets exhausted, it becomes dead.
It cannot be regenerated by reversing the current flow through the cell using the external direct current source of electrical energy.
Examples are dry cells and mercury cells.
Dry Cell:
These are used in toys, torches and other electronic devices, used for portable electrical devices. Leclanche is the form of dry cell.
Features :
$(i)$ Container of dry cell is made of zinc. This acts as one of the electrode. Container acts as anode. Zinc contained is lined with porous paper.
$(ii)$ The other electrode is of carbon which is in the centre of the cell. Graphite rod is surrounded by black paste of manganese dioxide.
$(iii)$ Moist mixture of ammonium chloride acts as a electrolyte which is kept next to zinc anode. The remaining space between the electrolyte and carbon cathode is filled up with a paste of ammonium chloride and manganese dioxide, where manganese dioxide is used ad depolariser. In some cells, ammonium chloride is replaced by zinc chloride.
$(iv)$ The cell is sealed with wax.
$(v)$ Now when cell is in operation, following reaction occurs:
Zinc gets oxidises to $Z{n^{2 + }}$ so reaction at anode: $$Zn \to Z{n^{2 + }} + 2{e^ - }$$
$(vi)$ The electrons get utilized at carbon(cathode) as ammonium ions are reduced. So cathodic reaction: $$2N{H_4}^ + + 2{e^ - } \to 2N{H_3} + {H_2}$$
$(vii)$ Overall reaction: $$Zn + 2N{H_4}^ + \to 2N{H_3} + Z{n^{2 + }} + {H_2}$$
$(viii)$ Hydrogen is oxidized by $Mn{O_2}$ as $$2Mn{O_2} + {H_2} \to 2MnO(OH)$$
$(ix)$ Ammonia produced at cathode will combine with zinc ions to form a complex ion $$Z{n^{2 + }} + 4N{H_3} \to {\left[ {Zn{{(N{H_3})}_4}} \right]_{2 + }}$$
Alkaline dry cell is same as ordinary dry cell, in this potassium hydroxide is used.
Anode: $$Zn + 2O{H^ - } \to Zn{(OH)_2} + 2{e^ - }$$
Cathode: $$2Mn{O_2} + 2{H_2}O + 2{e^ - } \to 2MnO(OH) + 2O{H^ - }$$
Overall Reaction: $$Zn + 2Mn{O_2} + 2{H_2}O \to 2MnO(OH) + Zn{(OH)_2}$$
Since ammonium chloride in aqueous medium is slightly acidic so $Zn$ will get corrode and will stop the functioning of the cell. Due to corrosion and consumption of $Zn$ in this process, it cannot be used again.
Mercury Cell:
It is also called mercury oxide battery. It is non-rechargable, electrochemical primary battery cell. In these batteries, a reaction occurs between mercuric oxide and zinc electrodes in an alkaline electrolyte. The voltage remains practically constant which is about $1.35 volts$ during discharging. Capacity of these batteries are much more than the zinc carbon batteries. These were used in the shape of button and hence also called button cells for watches, hearing aids, cameras and calculators etc.
Features:
$(i)$ These batteries use either pure mercury $(II)$ oxide $(HgO)$ also called mercuric oxide, or a mixture of $HgO$ and Manganese dioxide which acts as cathode. Since $HgO$ is a non-conductor, so graphite is mixed with it. Graphite is mixed to prevent the collection of mercury into large droplets.
So reaction at cathode: $$HgO + {H_2}O + 2{e^ - } \to Hg + 2O{H^ - }$$
$(ii)$ Anode is made of zinc (Zn) and is separated from cathode with a layer of paper or other porous material soaked with electrolyte, which is called salt bridge. Reaction at anode: $$Zn + 4O{H^ - } \to Zn{(OH)_4}^{2 - } + 2{e^ - }$$
It is followed by chemical reaction which is as follows:$$Zn{(OH)_4}^{2 - } \to ZnO + 2O{H^ - } + {H_2}O$$
So overall anodic reaction is $$Zn + 2O{H^ - } \to ZnO + 2{e^ - } + {H_2}O$$
$(iii)$ Overall reaction :$$Zn + HgO \to ZnO + Hg$$
During discharge, zinc loses electrons (gets oxidized) to become zinc oxide while mercuric oxide gets reduced (gains electrons) to form elemental mercury. A little mercury is kept in the cell to prevent hydrogen gas evolution at the end.
$(iv)$ Sodium hydroxide or potassium hydroxide is used as electrolyte.
$NaOH$ cells have nearly constant voltage at low discharge currents making them suitable to use in calculators, torches etc.
$KOH$ cells have nearly constant voltage at higher currents, and are useful in applications like photographic cameras with flash, and watch with backlight. $KOH$ cells give better performances at lower temperatures.
A different form of mercury cells use mercuric oxide and cadmium, which has lower terminal voltage at lower energy density.
