Welcome to Topic 5.6: Electrode Potentials
Welcome to one of the most practical and scoring topics in CCEA A2 Chemistry! In this unit, we explore how chemical reactions can generate electricity and how we can use standard electrode potentials to predict whether a redox reaction will occur spontaneously. Don't worry if electrochemistry felt daunting at AS Level—we will build every idea step-by-step with clear analogies, standard conventions, and exam-focused tips.
1. Understanding Half-Cells and Standard Conditions
What is a Half-Cell?
When a piece of metal is placed into a solution of its own ions (for example, a strip of zinc metal in a solution of zinc sulfate, \(\text{Zn}^{2+}_{(aq)}\)), an equilibrium is established at the surface of the metal:
\(\text{Zn}^{2+}_{(aq)} + 2\text{e}^- \rightleftharpoons \text{Zn}_{(s)}\)
Some zinc atoms lose electrons to form aqueous ions and dissolve into the solution, while some zinc ions gain electrons from the metal to become solid atoms. This separation of charge between the solid metal and the surrounding solution sets up an electrical potential known as an electrode potential.
Standard Conditions
Because the position of equilibrium changes if we alter temperature, pressure, or concentration, scientists measure electrode potentials under strictly controlled standard conditions:
• Temperature: \(298\text{ K}\) (\(25^\circ\text{C}\))
• Pressure: \(100\text{ kPa}\) (or \(1\text{ atmosphere}\)) for any gases involved
• Concentration: \(1.00\text{ mol dm}^{-3}\) for all aqueous ions
Standard Electrode Potential (\(E^\theta\))
Key Definition: The Standard Electrode Potential (\(E^\theta\)) of a half-cell is the electromotive force (EMF) generated by a standard half-cell compared to a standard hydrogen electrode (SHE) under standard conditions of \(298\text{ K}\), \(100\text{ kPa}\), and \(1.00\text{ mol dm}^{-3}\) aqueous ion concentration.
Key Takeaway: We cannot measure the absolute potential of an isolated half-cell alone; we can only measure the difference in potential when it is connected to a reference standard.
2. The Standard Hydrogen Electrode (SHE)
The Zero Reference Standard
To compare different half-cells, chemists chose a universal baseline: the Standard Hydrogen Electrode (SHE). By international agreement, the standard electrode potential of the SHE is assigned a value of exactly \(0.00\text{ V}\) (\(E^\theta = 0.00\text{ V}\)).
Components of the SHE
The SHE consists of the following components set up under standard conditions:
1. Hydrogen gas (\(\text{H}_{2(g)}\)): Bubbled into the solution at a constant pressure of \(100\text{ kPa}\).
2. Acid solution: Containing hydrogen ions with a concentration of \([\text{H}^+_{(aq)}] = 1.00\text{ mol dm}^{-3}\) (e.g., \(1.00\text{ mol dm}^{-3}\text{ HCl}_{(aq)}\) or \(0.50\text{ mol dm}^{-3}\text{ H}_2\text{SO}_{4(aq)}\)).
3. Platinum electrode: A piece of inert platinum foil coated with finely divided platinum black.
Equilibrium Half-Equation:
\(2\text{H}^+_{(aq)} + 2\text{e}^- \rightleftharpoons \text{H}_{2(g)}\) \(E^\theta = 0.00\text{ V}\)
Why Use Platinum?
Platinum is used because it is chemically inert (it will not react with the acid or take part in the redox reaction) and acts as an electrical conductor. The platinum black coating provides a high surface area to act as a catalyst for the rapid transfer of electrons between hydrogen gas and aqueous hydrogen ions.
Key Takeaway: The SHE is the universal reference electrode against which all other standard electrode potentials are measured.
3. Types of Half-Cells
In CCEA A2 Chemistry, you need to recognize and draw three main types of half-cells:
Type 1: Metal in Contact with its Aqueous Ions
Consists of a solid metal strip dipping into a \(1.00\text{ mol dm}^{-3}\) solution of its own ions.
Example: A strip of zinc dipping into aqueous zinc sulfate.
Equilibrium: \(\text{Zn}^{2+}_{(aq)} + 2\text{e}^- \rightleftharpoons \text{Zn}_{(s)}\)
Type 2: Gas in Contact with Aqueous Ions
Consists of a gas bubbled over an inert platinum electrode dipping into a \(1.00\text{ mol dm}^{-3}\) solution of the non-metal ions.
Example: Chlorine gas bubbled over a platinum electrode in a solution of chloride ions.
Equilibrium: \(\text{Cl}_{2(g)} + 2\text{e}^- \rightleftharpoons 2\text{Cl}^-_{(aq)}\)
Type 3: Solution of Ions of the Same Element in Two Different Oxidation States
Consists of an inert platinum electrode dipping into an equimolar mixture containing \(1.00\text{ mol dm}^{-3}\) of both ions.
Example: A platinum electrode dipping into a solution containing both \(\text{Fe}^{2+}_{(aq)}\) and \(\text{Fe}^{3+}_{(aq)}\).
Equilibrium: \(\text{Fe}^{3+}_{(aq)} + \text{e}^- \rightleftharpoons \text{Fe}^{2+}_{(aq)}\)
Exam Hint: Whenever a half-cell does not contain a solid conducting metal (such as gas systems or ion-ion solutions), you must include an inert platinum electrode in your description and diagram!
4. Building an Electrochemical Cell and Cell Notation
Constructing a Complete Electrochemical Cell
To construct a working electrochemical cell that measures standard electrode potentials, two half-cells are connected together using two key components:
1. High-Resistance Voltmeter: Connected in the external circuit via metal wires. The high resistance prevents current from flowing, ensuring that the maximum potential difference (electromotive force / EMF) is measured without disturbing the equilibrium positions.
2. Salt Bridge: Typically a strip of filter paper soaked in an inert aqueous electrolyte such as concentrated potassium nitrate (\(\text{KNO}_{3(aq)}\)) or potassium chloride (\(\text{KCl}_{(aq)}\)).
• Function of the Salt Bridge: It completes the electrical circuit and allows the movement of ions between the two half-cells to maintain electrical neutrality. The electrolyte used must not react with any of the ions present in either half-cell.
Crucial Distinction: Electrons move only through the external wires and electrodes; ions move only through the salt bridge and solutions!
IUPAC Conventional Cell Representation
To avoid drawing full apparatus diagrams every time, chemists use standard shorthand notation governed by IUPAC rules:
• Left-hand side (LHS): By convention, the half-cell where oxidation occurs (the more negative/less positive \(E^\theta\)) is placed on the left.
• Right-hand side (RHS): The half-cell where reduction occurs (the more positive \(E^\theta\)) is placed on the right.
• Single vertical line (\(|\)): Represents a phase boundary (e.g., between a solid electrode and an aqueous solution).
• Double vertical line (\(\parallel\)): Represents the salt bridge.
• Comma (,\()): Separates species that are in the same phase. • Position rule: The species in their most oxidised states are placed closest to the salt bridge.
Example: The Standard Zinc-Copper Cell \)\text{Zn}^{2+}_{(aq)} + 2\text{e}^- \rightleftharpoons \text{Zn}_{(s)}\) \(E^\theta = -0.76\text{ V}\) (Oxidation / LHS)
\(\text{Cu}^{2+}_{(aq)} + 2\text{e}^- \rightleftharpoons \text{Cu}_{(s)}\) \(E^\theta = +0.34\text{ V}\) (Reduction / RHS)
Cell notation:
\(\text{Zn}_{(s)} \mid \text{Zn}^{2+}_{(aq)} \parallel \text{Cu}^{2+}_{(aq)} \mid \text{Cu}_{(s)}\)
Example: Cell with Platinum Electrodes
\(\text{Pt}_{(s)} \mid \text{H}_{2(g)} \mid \text{H}^+_{(aq)} \parallel \text{Fe}^{3+}_{(aq)}, \text{Fe}^{2+}_{(aq)} \mid \text{Pt}_{(s)}\)
5. Electromotive Force (EMF) & Predicting Redox Feasibility
Calculating Cell EMF (\(E^\theta_{\text{cell}}\))
The standard cell potential (electromotive force) is calculated using:
\(E^\theta_{\text{cell}} = E^\theta_{\text{RHS}} - E^\theta_{\text{LHS}}\)
or
\(E^\theta_{\text{cell}} = E^\theta_{\text{reduction}} - E^\theta_{\text{oxidation}}\)
or
\(E^\theta_{\text{cell}} = E^\theta_{\text{more positive}} - E^\theta_{\text{less positive}}\)
Worked Example:
For the Zinc-Copper cell:
\(E^\theta_{\text{cell}} = E^\theta(\text{Cu}^{2+}/\text{Cu}) - E^\theta(\text{Zn}^{2+}/\text{Zn})\)
\(E^\theta_{\text{cell}} = (+0.34\text{ V}) - (-0.76\text{ V}) = +0.34\text{ V} + 0.76\text{ V} = +1.10\text{ V}\)
The Electrochemical Series and Redox Trends
All standard electrode potentials are written as reduction equilibria (\(\text{Oxidised form} + n\text{e}^- \rightleftharpoons \text{Reduced form}\)):
• More positive (or less negative) \(E^\theta\): Stronger tendency to gain electrons. The species on the left-hand side of the half-equation is a stronger oxidising agent and will undergo reduction (forward reaction favoured).
• More negative (or less positive) \(E^\theta\): Stronger tendency to lose electrons. The species on the right-hand side of the half-equation is a stronger reducing agent and will undergo oxidation (reverse reaction favoured).
Thermodynamic Feasibility Rule
A proposed redox reaction is thermodynamically feasible under standard conditions if the overall standard cell potential is positive:
\(E^\theta_{\text{cell}} > 0\text{ V}\)
Limitations of \(E^\theta\) Predictions
Even if a calculation shows \(E^\theta_{\text{cell}} > 0\text{ V}\), a reaction might not be observed in the laboratory due to two major limitations:
1. Kinetic Limitation (High Activation Energy):
Thermodynamic feasibility tells us only that a reaction can happen energetically, not how fast it will occur. If the reaction has a very high activation energy (\(E_a\)), the rate of reaction at \(298\text{ K}\) will be imperceptibly slow.
2. Non-Standard Conditions:
If concentrations, pressures, or temperatures differ from standard conditions (\(1.00\text{ mol dm}^{-3}\), \(100\text{ kPa}\), \(298\text{ K}\)), the position of equilibrium shifts in accordance with Le Chatelier's Principle, altering the electrode potential value away from \(E^\theta\). A reaction that is unfeasible under standard conditions might become feasible under non-standard conditions.
Example of Concentration Effect:
For \(\text{Cu}^{2+}_{(aq)} + 2\text{e}^- \rightleftharpoons \text{Cu}_{(s)}\) (\(E^\theta = +0.34\text{ V}\)):
If \([\text{Cu}^{2+}]\) is decreased below \(1.00\text{ mol dm}^{-3}\), the equilibrium shifts to the left to replace the lost ions. Releasing electrons makes the electrode potential less positive.
6. Commercial Storage Cells: Lithium-ion Cells
Electrode potentials find widespread everyday use in portable energy storage, most notably in rechargeable lithium-ion cells (found in mobile phones, laptops, and electric vehicles).
Reactions During Discharge (CCEA 5.6.5)
When a lithium-ion cell is discharging and supplying power to an external circuit:
• Negative electrode (anode on discharge / oxidation):
Lithium atoms release electrons:
\(\text{Li} \rightarrow \text{Li}^+ + \text{e}^-\)
(Often represented in intercalated graphite as: \(\text{LiC}_6 \rightarrow \text{C}_6 + \text{Li}^+ + \text{e}^-\))
• Positive electrode (cathode on discharge / reduction):
Lithium ions are incorporated into cobalt(IV) oxide alongside incoming electrons:
\(\text{Li}^+ + \text{CoO}_2 + \text{e}^- \rightarrow \text{LiCoO}_2\)
• Overall cell discharge reaction:
\(\text{Li} + \text{CoO}_2 \rightarrow \text{LiCoO}_2\)
During recharging, an external power supply forces current in the opposite direction, reversing these half-equations and restoring the original reactants.
7. Pitfalls and Examiner Cautions
Keep these common exam traps in mind when tackling CCEA A2 questions:
• Salt Bridge Traps: Never say that electrons travel through the salt bridge. Always state clearly that ions move through the salt bridge to maintain electrical neutrality and complete the circuit.
• Double Negative Arithmetic Errors: When calculating \(E^\theta_{\text{cell}} = E^\theta_{\text{red}} - E^\theta_{\text{ox}}\), subtract the oxidation value directly. Do not change the sign before plugging it into the subtraction formula.
• Stoichiometric Multipliers: If you multiply a half-equation by 2 or 3 to balance electrons when writing an overall redox equation, do not multiply the \(E^\theta\) value. Electrode potential is an intensive property and does not depend on the number of moles reacted.
• Feasibility vs. Rate: If an examiner asks why a feasible reaction (\(E^\theta_{\text{cell}} > 0\)) fails to occur in practice at room temperature, the answer is always a high activation energy (\(E_a\)), meaning the rate is too slow to detect.
• Missing Platinum: Always include \(\text{Pt}_{(s)}\) on the outer ends of cell diagrams and cell notation when the half-cell consists purely of solutions of ions (e.g., \(\text{Fe}^{3+}/\text{Fe}^{2+}\)) or gases (e.g., \(\text{H}_2/\text{H}^+\)).
Quick Summary Review
• \(E^\theta\) is measured under \(298\text{ K}\), \(100\text{ kPa}\), and \(1.00\text{ mol dm}^{-3}\) relative to the SHE (\(E^\theta = 0.00\text{ V}\)).
• More positive \(E^\theta\) = stronger oxidising agent = undergoes reduction (on the RHS).
• More negative \(E^\theta\) = stronger reducing agent = undergoes oxidation (on the LHS).
• \(E^\theta_{\text{cell}} = E^\theta_{\text{RHS}} - E^\theta_{\text{LHS}}\). A positive \(E^\theta_{\text{cell}}\) means the reaction is thermodynamically feasible.
• Reactions may still not happen due to high activation energy or non-standard conditions.
• Lithium-ion cell discharge reaction: \(\text{Li} + \text{CoO}_2 \rightarrow \text{LiCoO}_2\).