Galvanic Cells and Redox Reactions

Learning Objectives

• Understand spontaneous chemical changes through redox reactions.
• Describe the function and components of galvanic cells.
• Use cell notation to represent galvanic cells.

Spontaneous Redox Reaction Example

• A demonstration is shown by immersing a coiled copper wire into a silver nitrate solution, resulting in a change in color.
• Observations:
  • Colorless solution turns blue due to formation of $Cu^{2+}(aq)$.
  • Copper wire surface becomes covered with gray solid silver $Ag(s)$.
• Chemical reactions involved:
  • Overall reaction: $Cu(s) + 2Ag^+(aq) <br /> ightarrow Cu^{2+}(aq) + 2Ag(s)$
  • Oxidation half-reaction: $Cu(s) <br /> ightarrow Cu^{2+}(aq) + 2e^-$
  • Reduction half-reaction: $2Ag^+(aq) + 2e^- <br /> ightarrow 2Ag(s)$

Galvanic Cells

• Defined as electrochemical cells where spontaneous redox reactions occur without direct contact between reactants.
• Composition of a galvanic cell:
  • Two half-cells, each containing a redox pair (reactant and product).
  • Left Half-Cell: Contains $Cu(0)/Cu(II)$ couple with solid copper and copper nitrate solution.
  • Right Half-Cell: Contains $Ag(I)/Ag(0)$ couple with solid silver and silver nitrate solution.
  • Electrodes: Copper (anode) where oxidation occurs; Silver (cathode) where reduction occurs.
• Charge balance is maintained using a salt bridge, allowing the flow of inert ions.

Cell Reaction Dynamics

• Copper ions ($Cu^{2+}$) are generated at the anode while silver ions ($Ag^+$) are consumed at the cathode.
• The influx of ions from the salt bridge maintains electrical neutrality as concentrations change.

Cell Notation

• Cell notation provides a symbolic representation of galvanic cells:
  • Components are written using chemical formulas.
  • Interfaces are shown with vertical lines; separate phases are denoted by commas.
  • Example Notation: $Cu(s) | 1M ext{ }Cu(NO<em>3)</em>2(aq) ext{ } || ext{ } 1M ext{ }AgNO_3(aq) | Ag(s)$

Example of Magnesium and Iron Galvanic Cell

• A system with solid magnesium and aqueous iron(III) ions:
  • Overall Cell Reaction: $Mg(s) + 2Fe^{3+}(aq) <br /> ightarrow Mg^{2+}(aq) + 2Fe^{2+}(aq)$
  • Oxidation Half-Reaction: $Mg(s) <br /> ightarrow Mg^{2+}(aq) + 2e^-$
  • Reduction Half-Reaction: $2Fe^{3+}(aq) + 2e^- <br /> ightarrow 2Fe^{2+}(aq)$
• Notation would be:
  $Mg(s) | 0.1M ext{ }MgCl<em>2(aq) || 0.2M ext{ }FeCl</em>3(aq), 0.3M ext{ }FeCl_2(aq) | Pt(s)$

Practical Example

• Given a galvanic cell with Chromium and Copper:
  • Schematic: $Cr(s)|1M ext{ }CrCl<em>3(aq) || 1M ext{ }CuCl</em>2(aq) | Cu(s)$
  • Half Reactions:
  • Oxidation (Anode): $Cr(s) <br /> ightarrow Cr^{3+}(aq) + 3e^-$
  • Reduction (Cathode): $Cu^{2+}(aq) + 2e^- <br /> ightarrow Cu(s)$
  • Adjust electrons for balancing, leading to:
  • Overall Reaction: $2Cr(s) + 3Cu^{2+}(aq) <br /> ightarrow 2Cr^{3+}(aq) + 3Cu(s)$

Learning Check

• Write a schematic for a galvanic cell reaction of:
  $Sn^{4+}(aq) + Zn(s) <br /> ightarrow Sn^{2+}(aq) + Zn^{2+}(aq)$
• Schematic Answer: $Zn(s) | Zn^{2+}(aq) || Sn^{4+}(aq), Sn^{2+}(aq) | Pt(s)$