Non-Rechargeable (Primary) Galvanic Cells

Introduction

Galvanic cells (also known as voltaic cells) are electrochemical cells that convert chemical energy into electrical energy through spontaneous redox reactions. Primary cells are non-rechargeable, meaning the chemical reaction proceeds in one direction only, and once the reactants are depleted, the cell cannot be restored to its original state. Understanding their design and operation is crucial.

Common Design Features

• Two Half-Cells: A galvanic cell consists of two separate half-cells, each containing an electrode immersed in an electrolyte solution. These half-cells are connected to allow for the flow of ions and electrons.
• Electrodes:
  • Anode: The electrode where oxidation occurs (loss of electrons). It is the negative terminal in a galvanic cell.
  • Cathode: The electrode where reduction occurs (gain of electrons). It is the positive terminal in a galvanic cell.
  • Electrodes can be reactive (participating in the redox reaction) or inert (providing a surface for the reaction without being consumed).
• Electrolyte Solutions: Each half-cell contains an electrolyte solution, which is a substance that conducts electricity through the movement of ions. The electrolyte must contain ions that can participate in the half-cell reactions or facilitate charge transfer.
• Salt Bridge: A salt bridge connects the two half-cells, allowing for the migration of ions to maintain electrical neutrality in the half-cells. It typically contains an inert electrolyte (e.g., $KNO_3$, $KCl$) that does not interfere with the cell reactions. Without a salt bridge, charge would build up in the half-cells, quickly stopping the reaction.
• External Circuit: An external circuit (e.g., a wire) connects the two electrodes, allowing electrons to flow from the anode to the cathode, creating an electric current. A voltmeter is typically included in the external circuit to measure the cell potential (voltage).

General Operating Principles

1. Redox Reactions: The core of a galvanic cell is a spontaneous redox reaction, where one species is oxidized (loses electrons) and another is reduced (gains electrons).
2. Separation of Half-Reactions: The oxidation and reduction half-reactions are physically separated into the two half-cells. This separation forces the electrons released during oxidation to flow through the external circuit to reach the cathode, generating an electric current.
3. Electron Flow: Electrons flow from the anode (where oxidation occurs) through the external circuit to the cathode (where reduction occurs).
4. Ion Flow: Ions flow through the salt bridge to maintain charge neutrality in the half-cells. Anions migrate from the cathode half-cell to the anode half-cell, and cations migrate from the anode half-cell to the cathode half-cell.
5. Electrode Polarities:
   • Anode: Negative (-), as it is the source of electrons.
   • Cathode: Positive (+), as it receives electrons.
6. Energy Conversion: Chemical energy stored in the reactants is converted into electrical energy as the redox reaction proceeds. This electrical energy can then be used to power external devices.

Role of Electrodes and Electrolytes

• Reactive Electrodes: Reactive electrodes directly participate in the redox reaction. For example, in a zinc-copper cell, the zinc electrode is oxidized to $Zn^{2+}$ ions, and the copper electrode is the site where $Cu^{2+}$ ions are reduced to solid copper. The mass of the anode decreases, and the mass of the cathode increases (if a metal).

• Inert Electrodes: Inert electrodes (e.g., platinum, carbon) do not participate directly in the redox reaction but provide a surface for electron transfer. They are used when the redox reaction involves species in solution (e.g., oxidation of $Fe^{2+}$ to $Fe^{3+}$).

• Electrolyte Solutions: The electrolyte solutions provide the ions necessary for the half-cell reactions and facilitate charge transport. The choice of electrolyte depends on the specific redox reaction taking place.

Key Terms

• Electrode: A conductor through which electricity enters or leaves an object, substance, or device.
• Electrolyte: A substance that produces an electrically conducting solution when dissolved in a polar solvent, such as water. Electrolytes carry electric current through the movement of ions.
• Anode: The electrode at which oxidation occurs.
• Cathode: The electrode at which reduction occurs.
• Salt Bridge: A connection containing an electrolyte between the oxidation and reduction half-cells, which maintains electrical neutrality by allowing the flow of ions.
• Half-Cell: One of the two compartments of a galvanic cell, containing an electrode and an electrolyte where either oxidation or reduction occurs.

Limitations

• Non-Rechargeable: Once the reactants are consumed, primary cells cannot be recharged and must be discarded.
• Voltage Decrease: As the cell operates, the concentrations of reactants decrease, leading to a gradual decrease in the cell voltage.

Example: Zinc-Carbon Dry Cell (Leclanché Cell)

This is a common type of primary cell.

• Anode: Zinc casing (oxidation of zinc)
• Cathode: Carbon rod (reduction of manganese dioxide)
• Electrolyte: Paste of ammonium chloride ($NH_4Cl$), zinc chloride ($ZnCl_2$), and manganese dioxide ($MnO_2$)

Reactions:

• Anode (Oxidation): $Zn(s)
  ightarrow Zn^{2+}(aq) + 2e^-$
• Cathode (Reduction): $2MnO_2(s) + 2NH_4^+(aq) + 2e^-
  ightarrow Mn_2O_3(s) + 2NH_3(aq) + H_2O(l)$

Galvanic Cell Diagram (Description)

Imagine a diagram with two beakers. The left beaker contains a zinc electrode in a solution of zinc sulfate ($ZnSO_4$). The right beaker contains a copper electrode in a solution of copper sulfate ($CuSO_4$). A wire connects the zinc and copper electrodes, with a voltmeter in the circuit to measure the voltage. A salt bridge, filled with potassium nitrate ($KNO_3$), connects the two beakers. The zinc electrode is labeled as the anode (-), and the copper electrode is labeled as the cathode (+). Arrows show the flow of electrons from the zinc electrode to the copper electrode through the wire, and the flow of nitrate ions ($NO_3^−$) from the salt bridge into the zinc sulfate solution, and potassium ions ($K^+$) into the copper sulfate solution.