Galvanic Cells and Batteries

Very few of the cells obtained by combining the electrodes in Table 1 in Electromotive Force of Galvanic Cells are suitable for everyday use as a source of electrical energy. The chief reason for this is that most of them can only deliver a very small current per unit area of electrode and need to be made very large before they become useful. A less serious difficulty is that they all involve solutions, many of them poisonous or corrosive which must be stored in a robust splash-proof container. Neither of these two difficulties is encountered with the common flashlight battery, known as the Leclanché cell or dry cell.

Despite the name dry cell, this battery does contain an electrolyte solution but only in the form of a thick paste. A saturated solution of ammonium chloride also containing zinc chloride is mixed with ammonium chloride crystals, as well as some inert filler like diatomaceous earth. As shown in Fig. \(\PageIndex{1}\), the center of the dry cell is occupied by a graphite rod which functions purely as a conductor in much the same way as a platinum wire in laboratory cells described in the previous section.

Figure \(\PageIndex{1}\) The Leclanché or dry cell.

This rod is surrounded by a powdered mixture of graphite and manganese dioxide, MnO₂. This powder in turn is surrounded by the NH₄Cl–ZnCl₂ paste, and the whole is encased in a thin zinc cylinder which acts as the second electrode and also as the container for the cell. In some cases a further cylinder of steel is wrapped around the zinc for added mechanical strength. In shorthand notation the dry cell corresponds to

Zn│Zn²⁺, NH₄⁺│MnO₂(s), C

while the electrode half-equations can be represented as

\[\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + \text{2}e^{-}\]

and \[\text{2MnO}_2(s) + \text{2NH}_4^{+} + 2e^{-} \rightarrow \text{2MnO(OH)}(s) + \text{2NH}_3 (aq)\]

in which the manganese is reduced from Mn(IV) in MnO₂ to Mn(III) in MnO(OH). Note that this cell contains no salt bridge or other device for separating the two electrodes. This is possible because both the oxidizing agent (MnO₂) and the reducing agent (Zn) are solids and so cannot diffuse toward each other and react. Nevertheless dry cells have a limited life. The zinc electrode is eventually eaten away by the slightly acidic ammonium chloride solution, which can also ruin a flashlight if old batteries are left inside. A recent improvement to the dry cell uses sodium hydroxide in place of ammonium chloride. Such an alkaline battery delivers a larger quantity of electrical energy but is also more expensive because of the necessity for a stronger casing to prevent leakage. Both types of batteries deliver a potential difference of about 1.5 V.

Another commonly used modem battery, found in electric and electronic watches, hearing aids, and light meters is the mercury cell. This cell has the form

Zn(Hg), ZnO(s)│OH^–(8 M)│HgO(s), Hg(l)

with the following electrode reactions:

\[\text{Zn} + \text{2OH}^{-}(aq) \rightarrow \text{ZnO}(s) + \text{H}_2\text{O} + 2e^{-}\]

and \[\text{HgO}(s) + \text{H}_2\text{O} + 2e^{-} \rightarrow \text{Hg}(l) + \text{2OH}^{-}\]

Again the use of a solid reducing agent and a solid oxidizing agent obviates the necessity of separating the contents of the two electrodes. A mercury, cell delivers a potential difference of about 1.34 V.

From ChemPRIME: 17.7: Galvanic Cells

Contributors and Attributions

• Ed Vitz (Kutztown University), John W. Moore (UW-Madison), Justin Shorb (Hope College), Xavier Prat-Resina (University of Minnesota Rochester), Tim Wendorff, and Adam Hahn.