Question

Describe a lead-acid storage battery. Include equations for the anode and cathode reactions.

Short answer

A lead-acid battery consists of cells with lead anodes and lead dioxide cathodes in sulfuric acid. Anode reaction: \( Pb(s) + SO4^{2-}(aq) \rightarrow PbSO4(s) + 2e^- \). Cathode reaction: \( PbO2(s) + SO4^{2-}(aq) + 4H^+(aq) + 2e^- \rightarrow PbSO4(s) + 2H2O(l) \).

Step-by-step solution

Describing the Lead-Acid Battery Components

A lead-acid storage battery consists of a series of cells, each containing a lead (Pb) anode and a lead dioxide (PbO2) cathode, immersed in a sulfuric acid (H2SO4) electrolyte solution. During discharge, the sulfuric acid reacts with the anode and cathode to create an electric current.

Anode Reaction Equation

At the anode, the lead (Pb) reacts with the sulfate ions (SO4^2-) from the sulfuric acid to produce lead sulfate (PbSO4) and electrons (e^-). The anode reaction can be represented by the following half-equation: \[Pb(s) + SO4^{2-}(aq) \rightarrow PbSO4(s) + 2e^-\]

Cathode Reaction Equation

At the cathode, lead dioxide (PbO2) reacts with the sulfate ions (SO4^2-) and hydrogen ions (H+) produced by the sulfuric acid, along with electrons from the external circuit, to produce lead sulfate (PbSO4), water (H2O), and electrical energy. The cathode reaction is represented by the half-equation: \[PbO2(s) + SO4^{2-}(aq) + 4H^+(aq) + 2e^- \rightarrow PbSO4(s) + 2H2O(l)\]

Overall Cell Reaction

The overall cell reaction for the lead-acid battery during discharge is obtained by combining the anode and cathode reactions. The two half-equations from Steps 2 and 3 can be summed up to give: \[Pb(s) + PbO2(s) + 2SO4^{2-}(aq) + 4H^+(aq) \rightarrow 2PbSO4(s) + 2H2O(l)\] During charging, the reverse reactions occur at each electrode.

Key concepts

Anode and Cathode Reactions in Lead-Acid Batteries

In a lead-acid battery, essential electrochemical reactions occur at the anode and cathode that allow the battery to store and release energy.

At the anode, lead (Pb) reacts with sulfate (SO4^{2-}) from the sulfuric acid (H2SO4) electrolyte to form lead sulfate (PbSO4) and release electrons in the process. This is a crucial oxidation reaction where lead loses electrons, represented by the equation: \[Pb(s) + SO4^{2-}(aq) \rightarrow PbSO4(s) + 2e^-\].

Conversely, at the cathode, lead dioxide (PbO2) accepts electrons from the external circuit and reacts with sulfate ions (SO4^{2-}) and hydrogen ions (H^+) from the acid. This reduction reaction produces lead sulfate (PbSO4) and water (H2O), as shown by the equation: \[PbO2(s) + SO4^{2-}(aq) + 4H^+(aq) + 2e^- \rightarrow PbSO4(s) + 2H2O(l)\].

These reactions are reversible, which allows the battery to be recharged by applying an external voltage, driving the reverse reactions at each electrode.

Electrochemistry Fundamentals

Electrochemistry is the branch of chemistry that deals with the relationship between electricity and chemical changes. It revolves mainly around the concept of redox reactions—where oxidation (loss of electrons) and reduction (gain of electrons) occur.

In the context of a lead-acid battery, electrochemistry explains how electrical energy is converted to chemical energy and vice versa. The electrolyte (sulfuric acid in this case) facilitates the flow of ions, while the electrodes (lead and lead dioxide) are sites where the redox reactions happen, allowing electrons to move through an external circuit to do work. Understanding these principles is essential for grasping how batteries, as well as other electrochemical cells, operate.

Half-Reactions in Electrochemical Processes

Half-reactions are a way of representing the two-part process of redox reactions—one showing oxidation and the other showing reduction. In a lead-acid battery, each electrode hosts one half-reaction.

The anode half-reaction is the oxidation process where lead becomes lead sulfate while releasing electrons: \[Pb(s) + SO4^{2-}(aq) \rightarrow PbSO4(s) + 2e^-\].

The cathode half-reaction is the reduction process where lead dioxide becomes lead sulfate as it takes up electrons: \[PbO2(s) + SO4^{2-}(aq) + 4H^+(aq) + 2e^- \rightarrow PbSO4(s) + 2H2O(l)\].

These half-reactions illustrate the changes that occur at each electrode separately, and when combined, they yield the overall cell reaction. This split into half-reactions is a very useful method for analyzing and balancing complex redox reactions occurring in electrochemical cells.

Galvanic Cells and Battery Operation

Galvanic cells, like the cells in a lead-acid battery, are devices that generate electricity through spontaneous redox reactions. They consist of two electrodes—the anode where oxidation occurs, and the cathode where reduction takes place—connected by an electrolyte that facilitates ion exchange.

For the galvanic cell to function, there must be a flow of ions within the cell and a flow of electrons through an external circuit, which is what happens during the discharge of a lead-acid battery. The redox reactions at the anode and cathode push electrons through the external circuit while the internal ionic flow maintains electrical neutrality. This simultaneous movement is the essence of how galvanic cells like lead-acid batteries produce electrical energy efficiently from chemical processes.