3.1.11.1 Epot and Ecells improved

Redox reactions occur in electrochemical cells; electrons transfer from reducing agent to oxidizing agent via an external circuit.
A potential difference is created, enabling electrical work.
Electrochemical cells are crucial for powering portable electronic devices (e.g., mobile phones, tablets, laptops) and vehicles.

IUPAC Convention: Understand how to write half-equations for electrode reactions.
Cell Representation: Familiarity with the conventional representation of cells.
Measuring Electrode Potentials: Cells measure electrode potentials based on the standard hydrogen electrode.
Condition Importance: Recognition of conditions affecting the measurement of electrode potential (Nernst equation not required).
Standard Electrode Potential (EѲ): Defined under specific conditions: 298K, 100 kPa, and 1.00 mol dm−3.
Electrochemical Series: Standard electrode potentials can be organized into a series.

Predict Redox Reaction Direction: Utilize EѲ values to forecast simple redox reactions.
Calculate Cell EMF: Ability to compute electromotive force (EMF) of a cell.
Cell Representation Application: Write and implement the conventional representation of a cell.

Electron Transfer in Zinc Reaction:
- Reaction: Zn(s) → Zn2+(aq) + 2e- (Zinc ions leave the metal strip).
- Electrons remain on the metal surface where:
  - Zn2+(aq) + 2e- → Zn(s) (Metal ions recombine with electrons).
Equilibrium Establishment:
- An equilibrium is established by the reaction: Zn2+(aq) + 2e- ⇌ Zn(s).
- Forward Reaction: Zinc ions reduced.
- Backward Reaction: Zinc atoms oxidised.

Electrons on Metal Strip: Negatively charged electrons create a potential difference (voltage) between metal and solution.
Tendency and Voltage: The greater the metal's tendency to produce ions, the higher the potential difference at equilibrium.

Two metal strips:
- Zinc strip in zinc ion solution (Zn2+(aq)).
- Copper strip in copper ion solution (Cu2+(aq)).
Connection: Metal strips connected via a voltmeter.
Salt Bridge Role: Allows ion movement, completing the circuit (potassium nitrate solution).

The potential difference of a cell (EӨcell) can be computed:
Assemble half-reactions to find oxidation and reduction components:
- Example:
  - Half-Reactions:
    - Zn2+(aq) + 2e- ⇌ Zn(s) (E° = -0.76V)
    - Cu2+(aq) + 2e- ⇌ Cu(s) (E° = +0.34V)
- By convention, write reactions with the more negative potential on the left side (oxidization).
Calculation Example:
- For the Daniell cell: Ecell = Eright - Eleft = +0.34 - (-0.76)V = +1.10V.

When reactions occur in the Daniell cell:
- Zinc electrode diminishes; copper electrode sees deposition of copper solid.

Represents the reference half-cell for standard electrode potentials.
Standard conditions: 298K, 100 kPa, 1.0 mol dm−3 H+ ions.
Defined reaction:
- 2H+(aq) + 2e- ⇌ H2(g)
Construction of SHE:
- Platinum electrode in acid; continuous flow of hydrogen gas allows observation of potential measurements.

EѲ indicates the potential difference between an electrode at standard conditions and SHE.
Determining if reactions occur by connecting half-cells (e.g., Ag+/Ag and Cu2+/Cu) and calculating Ecell:
- Assemble half-reactions and apply cell conventions:
  - Ecell = 0.80V (Ag) - 0.34V (Cu) = +0.46V (spontaneous reaction).
Real-world implications: Copper reduces silver ions, forming solid silver.

Discuss the definition and examples of oxidizing agents.
Identify and analyze half-cell reactions and derive overall cell reactions from standard electrode potentials.
Reaction predictions based on electrode potential values; assess observable changes in solutions and deposits.