Electrochemistry P1

 Galvanic cell

 1. Introduction to Galvanic Cells

 

 Definition: A galvanic cell is a device that generates an electric current from a spontaneous chemical reaction.

 Historical Context: Named after Luigi Galvani (who observed muscle contractions in frogs due to electrical stimulation) and Alessandro Volta (who invented the first battery, the voltaic pile).

 Key Principle: The separation of oxidation and reduction half-reactions, allowing electron transfer to occur through an external circuit.

 

 2. Components of a Galvanic Cell

 

A typical galvanic cell consists of the following essential components:

 

 Anode:

     The electrode where oxidation occurs (loss of electrons).

     It is the negative terminal of the galvanic cell.

     Electrons flow from the anode.

 Cathode:

     The electrode where reduction occurs (gain of electrons).

     It is the positive terminal of the galvanic cell.

     Electrons flow to the cathode.

 Electrolytes:

     Solutions containing ions that can conduct electricity.

     Each electrode is immersed in an electrolyte containing ions of the same metal (or other suitable species).

 Salt Bridge:

    A U-shaped tube containing an inert electrolyte (e.g., KNO3, KCl) in a gel.

     Function:

         Maintains electrical neutrality in the half-cells by allowing the migration of ions.

         Prevents charge build-up, which would stop the flow of electrons.

         Completes the electrical circuit.

 External Circuit (Wire):

     Connects the anode and cathode, providing a path for electrons to flow from the anode to the cathode.

     Often includes a voltmeter or ammeter to measure potential difference or current.

 

 3. How a Galvanic Cell Works (The Daniell Cell Example)

 

Let's consider the Daniell cell, a classic example of a galvanic cell, which uses zinc and copper electrodes.

 

 Setup:

     Anode compartment: Zinc electrode immersed in a ZnSO4 solution.

     Cathode compartment: Copper electrode immersed in a CuSO4 solution.

     A salt bridge connects the two solutions.

     A wire connects the zinc and copper electrodes.

 

 Half-Reactions:

     At the Anode (Oxidation):

         Zinc metal loses two electrons and forms zinc ions, which dissolve into the solution.

       Zn(s)      Zn²⁺ (aq) + 2e^-

         The zinc electrode gradually decreases in mass.

     At the Cathode (Reduction):

         Copper ions from the solution gain two electrons and deposit as copper metal onto the electrode.

         Cu²⁺ (aq) + 2e^-       Cu(s)

         The copper electrode gradually increases in mass.

 

 Overall Cell Reaction:

     Summing the half-reactions, the electrons cancel out:

          Zn(s) + Cu²⁺ (aq)    Zn²⁺ (aq)  + Cu(s)

 

 Electron Flow:

     Electrons generated at the zinc anode flow through the external wire to the copper cathode.

 

 Ion Movement (Salt Bridge):

     As  Zn²⁺ (aq)  ions are produced at the anode, anions (e.g., NO- from the salt bridge) migrate into the anode compartment to neutralize the excess positive charge.

     As Cu²⁺ (aq) ions are consumed at the cathode, cations (e.g., K+ from the salt bridge) migrate into the cathode compartment to neutralize the excess negative charge left by the departing Cu²⁺ (aq) ions (or to compensate for the buildup of anions like SO4(-2).

 

 4. Cell Notation (Shorthand Representation)

 

Galvanic cells can be represented concisely using cell notation:

 

 Anode | Anode Electrolyte || Cathode Electrolyte | Cathode

 

 Rules:

     A single vertical line (|) represents a phase boundary (e.g., solid electrode in liquid electrolyte).

     A double vertical line (||) represents the salt bridge.

     The anode is always written on the left, and the cathode on the right.

     States of matter (s), (l), (g), (aq) are usually included.

 

 Example (Daniell Cell):

     Zn(s) | ZnSO4(aq) || CuSO4 (aq) | Cu(s)

     Or more generally: Zn(s) |  Zn²⁺ (aq)  || Cu²⁺ (aq) | Cu(s)

 

 5. Cell Potential (Electromotive Force, EMF)

 

 Definition: The cell potential (E{cell}) is the potential difference between the cathode and the anode, representing the driving force for the spontaneous redox reaction. It is measured in volts (V).

 Standard Cell Potential :Ecell⁰

     The cell potential measured under standard conditions:

         1 M concentration for all aqueous solutions.

         1 atm partial pressure for all gases.

         25 degree celcius (298 K).

 Calculation of Cell Potential:

    Ecell⁰ = E_cathode - E_anode (using standard reduction potentials)

     Standard reduction potentials are tabulated values, usually relative to the Standard Hydrogen Electrode (SHE).

 

 6. Applications of Galvanic Cells

 

Galvanic cells are the basis for many practical devices:

 

 Batteries:

     Primary Batteries (Non-rechargeable): Dry cells (e.g., zinc-carbon, alkaline), button cells. Reactions proceed until reactants are depleted.

     Secondary Batteries (Rechargeable): Lead-acid batteries (car batteries), lithium-ion batteries (smartphones, electric vehicles), nickel-cadmium batteries. Can be recharged by applying an external voltage to reverse the reaction.

 Fuel Cells:

     Galvanic cells that continuously convert the chemical energy of a fuel (e.g., hydrogen, methane) and an oxidant (e.g., oxygen) into electrical energy.

     Highly efficient and produce minimal pollution (e.g., hydrogen fuel cells produce only water).

 Corrosion Prevention:

     Sacrificial Anodes: A more reactive metal (e.g., magnesium, zinc) is connected to a less reactive metal (e.g., iron pipe) to protect it from corrosion. The more reactive metal acts as the anode and corrodes preferentially.

 

 7. Key Takeaways

 

 Galvanic cells convert chemical energy into electrical energy via spontaneous redox reactions.

 The anode is where oxidation occurs (negative terminal), and the cathode is where reduction occurs (positive terminal).

 The salt bridge maintains electrical neutrality.

 Electron flow is from anode to cathode through the external circuit.

 Cell potential Ecell is a measure of the driving force of the reaction.

 The Nernst equation allows for calculation of cell potential under non-standard conditions.

 Galvanic cells are the foundation of batteries and fuel cells.

 

This concludes our lecture on galvanic cells. Understanding these principles is crucial for comprehending a wide range of electrochemical phenomena and technologies.