An electrochemical cell is created by connecting two different half-cells with an external circuit and a salt bridge. The potential difference generated by this combined setup forces electrons to flow from one half-cell to the other. To standardise how we describe and diagram these cells, we use specific rules for diagrams and shorthand notation.
A pathway containing an unreactive electrolyte solution that connects the two half-cells of an electrochemical cell. It completes the electrical circuit by allowing the flow of ions to maintain charge neutrality in each half-cell.
🔑 Key Principle
By IUPAC convention, we write and draw electrochemical cells with the half-cell where oxidation occurs (the negative electrode, anode) on the left, and the half-cell where reduction occurs (the positive electrode, cathode) on the right. Under this convention, electrons always flow from left to right through the external wire.
Cell Structure and Diagram
A standard zinc-copper cell consists of a zinc half-cell and a copper half-cell. The zinc electrode has a standard potential of -0.76 V, and the copper electrode has +0.34 V. The zinc electrode is therefore placed on the left, and the copper electrode is placed on the right:
A shorthand written code that summarizes the structure of an electrochemical cell. It uses vertical lines to represent phase boundaries and a double vertical line to represent the salt bridge.
How the Salt Bridge Works
During the operation of the cell, chemical changes occur in the half-cells:
- At the left electrode, zinc atoms are oxidized: \( \text{Zn(s)} \rightarrow \text{Zn}^{2+}\text{(aq)} + 2\text{e}^- \). This adds positive zinc ions to the left solution.
- At the right electrode, copper ions are reduced: \( \text{Cu}^{2+}\text{(aq)} + 2\text{e}^- \rightarrow \text{Cu(s)} \). This removes positive copper ions, leaving behind negative sulphate ions.
If these charge changes were not balanced, a large charge imbalance would build up instantly, stopping electron flow. The salt bridge solves this problem. It is typically a piece of filter paper soaked in potassium nitrate (\( \text{KNO}_3 \)) solution. Ions migrate out of the salt bridge into the half-cells: negative nitrate ions (\( \text{NO}_3^- \)) flow into the zinc half-cell, and positive potassium ions (\( \text{K}^+ \)) flow into the copper half-cell. This maintains electrical neutrality without mixing the solutions.
When choosing a substance for a salt bridge in an exam, always choose an inert electrolyte like potassium nitrate. A common question asks why potassium chloride (\( \text{KCl} \)) is unsuitable if a half-cell contains silver ions (\( \text{Ag}^+ \)) or lead ions (\( \text{Pb}^{2+} \)). The chloride ions from the salt bridge will react with silver ions to form insoluble silver chloride (\( \text{AgCl} \)). This precipitate blocks the pore paths of the bridge, preventing ion flow and stopping the cell from working!
Rules for Standard IUPAC Cell Notation
To avoid drawing complex diagrams, we write cells using a standardized shorthand representation. The rules are:
- The half-cell with the more negative potential (undergoing oxidation) is written on the left.
- The half-cell with the more positive potential (undergoing reduction) is written on the right.
- A single vertical line \( | \) represents a phase boundary (e.g. between a solid electrode and an aqueous solution).
- A double vertical line \( || \) represents the salt bridge.
- Within each half-cell, the species in the lowest oxidation state is placed closest to the phase boundary with the solid electrode (i.e. on the outside of the notation).
- If a half-cell does not contain a solid reactant (e.g. it contains two aqueous ions like \( \text{Fe}^{3+} \) and \( \text{Fe}^{2+} \)), we must include an inert platinum electrode (\( \text{Pt} \)) on the outer edge, separating the aqueous ions with a comma instead of a vertical line.
For example, the standard zinc-copper cell is written as:
\( \text{Zn(s)} | \text{Zn}^{2+}\text{(aq)} || \text{Cu}^{2+}\text{(aq)} | \text{Cu(s)} \)
Step 1: Determine left and right placement:
Zinc has the more negative potential (-0.76 V), so it undergoes oxidation and is placed on the left. Silver has the more positive potential (+0.80 V), so it undergoes reduction and is placed on the right.
Step 2: Write the left half-cell notation:
Solid zinc is separated from zinc ions by a phase boundary:
\( \text{Zn(s)} | \text{Zn}^{2+}\text{(aq)} \)
Step 3: Write the right half-cell notation:
Silver ions are separated from solid silver by a phase boundary. The solution species is placed closest to the salt bridge:
\( \text{Ag}^+\text{(aq)} | \text{Ag(s)} \)
Step 4: Combine with the salt bridge:
\( \text{Zn(s)} | \text{Zn}^{2+}\text{(aq)} || \text{Ag}^+\text{(aq)} | \text{Ag(s)} \)
Step 5: Calculate standard cell potential:
\[ E^\theta_{\text{cell}} = E^\theta_{\text{positive}} - E^\theta_{\text{negative}} \]
\[ E^\theta_{\text{cell}} = (+0.80\text{ V}) - (-0.76\text{ V}) = +1.56\text{ V} \]
Step 1: Determine left and right placement:
Zinc is more negative (-0.76 V) and goes on the left. The iron half-cell is more positive (+0.77 V) and goes on the right.
Step 2: Write the left half-cell:
\( \text{Zn(s)} | \text{Zn}^{2+}\text{(aq)} \)
Step 3: Write the right half-cell:
The iron half-cell contains only aqueous ions (\( \text{Fe}^{3+} \) and \( \text{Fe}^{2+} \)). There is no solid reactant to act as an electrode, so an inert platinum electrode must be used. We separate the two ions in the same phase with a comma, and place the platinum electrode on the outside:
\( \text{Fe}^{3+}\text{(aq)}, \text{Fe}^{2+}\text{(aq)} | \text{Pt(s)} \)
Step 4: Combine with the salt bridge:
\( \text{Zn(s)} | \text{Zn}^{2+}\text{(aq)} || \text{Fe}^{3+}\text{(aq)}, \text{Fe}^{2+}\text{(aq)} | \text{Pt(s)} \)
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