Batteries — Primary & Secondary

What a Battery Really Is

Any battery that we use as a source of electrical energy is, at its core, a galvanic cell in which the chemical energy of a redox reaction is converted into electrical energy. The word battery strictly denotes one or more such cells connected in series, so that their potentials add up to a useful working voltage. The same Daniell-type machinery that produces $\ce{1.1\,V}$ from $\ce{Zn(s) + Cu^2+(aq) -> Zn^2+(aq) + Cu(s)}$ underlies every commercial cell — only the chemistry is chosen for portability and stability.

For a battery to be of practical use, NCERT lists three demands: it should be reasonably light, compact, and its voltage should not vary appreciably during use. The third condition is subtle. By the Nernst equation, the cell potential depends on the concentrations of dissolved ions; a cell whose net reaction changes ion concentrations will sag in voltage as it discharges. The best cell chemistries are those that keep the dissolved-ion activity nearly fixed — a theme that explains why the mercury cell holds its voltage so well.

On this design logic, all commercial cells fall into three families: primary cells (single-use), secondary cells (rechargeable), and fuel cells (continuously fed). This page covers the first two; fuel cells are treated separately as a sibling subtopic.

Primary vs Secondary Cells

The dividing line is whether the cell reaction can be reversed. In a primary cell the reaction occurs only once: after the active materials are consumed over a period of use, the cell becomes dead and cannot be reused. In a secondary cell, passing current through the cell in the opposite direction reverses the reaction and regenerates the original reactants, so the cell can undergo a large number of discharging and charging cycles. During recharge the secondary cell stops being a galvanic cell and behaves as an electrolytic cell, with an external source forcing a non-spontaneous reaction.

Keep this taxonomy firmly in mind: NEET frequently tests it as a single-line discriminator — for example, asking which of four named cells cannot be recharged, where the dry cell and mercury cell are the correct picks.

The Dry (Leclanché) Cell

The most familiar primary cell is the dry cell, known as the Leclanché cell after its discoverer, used in transistors, clocks and torches. It consists of a zinc container that also acts as the anode, and a carbon (graphite) rod as the cathode, the latter surrounded by powdered manganese dioxide ($\ce{MnO2}$) and carbon. The space between the electrodes is filled by a moist paste of ammonium chloride ($\ce{NH4Cl}$) and zinc chloride ($\ce{ZnCl2}$) which serves as the electrolyte. There is no free-flowing liquid — hence "dry" cell.

The actual electrode reactions are complex, but NCERT writes them approximately as:

Anode (oxidation): $$\ce{Zn(s) -> Zn^2+ + 2e^-}$$ Cathode (reduction): $$\ce{MnO2 + NH4+ + e^- -> MnO(OH) + NH3}$$

At the cathode, manganese is reduced from the +4 oxidation state to the +3 state. The ammonia produced is not released as gas; it forms a complex with $\ce{Zn^2+}$, namely $\ce{[Zn(NH3)4]^2+}$. The cell has a potential of nearly 1.5 V. Because the active materials are consumed and not regenerated, the dry cell is strictly a primary, single-use cell.

The Mercury Cell

The mercury cell is a primary cell suited to low-current devices such as hearing aids and watches. It uses a zinc–mercury amalgam, $\ce{Zn(Hg)}$, as the anode and a paste of $\ce{HgO}$ and carbon as the cathode. The electrolyte is a moist paste of $\ce{KOH}$ and $\ce{ZnO}$. The electrode reactions are:

Anode: $$\ce{Zn(Hg) + 2OH^- -> ZnO(s) + H2O + 2e^-}$$ Cathode: $$\ce{HgO + H2O + 2e^- -> Hg(l) + 2OH^-}$$ Overall: $$\ce{Zn(Hg) + HgO(s) -> ZnO(s) + Hg(l)}$$

The cell potential is approximately 1.35 V and — crucially — it remains constant during the cell's life. The reason is structural: the overall reaction involves only solids ($\ce{ZnO}$, $\ce{HgO}$) and liquid mercury, and no ion in solution whose concentration changes during the cell's lifetime. With no shifting ion activity, the Nernst term stays fixed and the voltage holds steady — exactly the stability that makes it ideal for instruments needing a reliable reference voltage.

The Lead Storage Battery

The most important secondary cell is the lead storage battery, commonly used in automobiles and inverters. It consists of a lead anode and a grid of lead packed with lead dioxide ($\ce{PbO2}$) as the cathode. A 38% solution of sulphuric acid (about 3.7 M) is used as the electrolyte. A practical car battery stacks several such cells in series to reach roughly 12 V.

Discharge reactions

When the battery is in use (discharging), the reactions are:

Anode: $$\ce{Pb(s) + SO4^2-(aq) -> PbSO4(s) + 2e^-}$$ Cathode: $$\ce{PbO2(s) + SO4^2-(aq) + 4H+(aq) + 2e^- -> PbSO4(s) + 2H2O(l)}$$ Overall: $$\ce{Pb(s) + PbO2(s) + 2H2SO4(aq) -> 2PbSO4(s) + 2H2O(l)}$$

Note two consequences. First, both electrodes become coated with insoluble $\ce{PbSO4}$ as the cell discharges. Second, sulphuric acid is consumed and water is produced, so the density of the electrolyte falls — which is exactly why a hydrometer reading on the acid tells a mechanic the state of charge of a car battery.

Charging reactions

On charging, an external source drives current through the cell in the opposite direction; the cell now behaves as an electrolytic cell and the discharge reaction is reversed. The $\ce{PbSO4}$ on the anode is converted back to $\ce{Pb}$ and the $\ce{PbSO4}$ on the cathode back to $\ce{PbO2}$, regenerating the original reactants:

Overall (charging): $$\ce{2PbSO4(s) + 2H2O(l) -> Pb(s) + PbO2(s) + 2H2SO4(aq)}$$

The Nickel-Cadmium Cell

Another important secondary cell is the nickel-cadmium (Ni-Cd) cell, built in a "jelly roll" arrangement with electrodes separated by a layer soaked in moist sodium or potassium hydroxide. It has a longer life than the lead storage cell but is more expensive to manufacture. NCERT does not require the individual electrode half-reactions; it states only the overall reaction during discharge:

Overall (discharge): $$\ce{Cd(s) + 2Ni(OH)3(s) -> CdO(s) + 2Ni(OH)2(s) + H2O(l)}$$

Here cadmium is oxidised (it is the reducing agent) and nickel is reduced from the +3 to the +2 state. Like all secondary cells, the Ni-Cd cell is recharged by reversing this reaction with an external current. Its alkaline electrolyte and solid-state chemistry give it a steady output and a long cycle life, which is why it has been favoured in cordless power tools, emergency lighting and older portable electronics.

Comparison, Voltages & Uses

The four standard cells of the NEET syllabus can be compressed into one reference table. Memorise the anode, cathode, electrolyte and voltage column by column; examiners mix and match these entries to build distractors.

One unifying thread runs through all four: the anode is always the more easily oxidised, electropositive metal — zinc, lead or cadmium — while the cathode hosts a higher-oxidation-state species ($\ce{MnO2}$, $\ce{HgO}$, $\ce{PbO2}$, $\ce{Ni(OH)3}$) that is reduced. Whether the cell is rechargeable depends solely on whether the products are stable solids that can be electrolytically converted back to the starting materials — which is precisely the case for the lead and nickel-cadmium systems but not for the dry or mercury cells.