⚡ Standard Electrode Potential
According to IUPAC, electrode potential is defined as the electromotive force (EMF) of a galvanic cell constructed using two electrodes. It is represented by the symbol \(E\). The absolute value of a single electrode potential cannot be measured directly; only the potential difference between two electrodes can be determined experimentally.
To measure the electrode potential of an unknown electrode, it is paired with a reference electrode of known potential in a galvanic cell. The cell potential obtained is the sum of both electrode potentials. Using a voltmeter, \(E_{\text{cell}}\) is measured, and since the reference potential is known, the unknown can be calculated.
🧪 The Standard Hydrogen Electrode (SHE)
The absolute value of a half‑cell potential cannot be measured; only the potential difference between two half‑cells is accessible. A standard reference electrode is needed to compare potentials reliably. The standard hydrogen electrode (SHE) serves as the universal reference and consists of:
- Hydrogen gas at 100 kPa (1 atm) pressure
- H⁺ ions at 1.00 mol dm⁻³ (1 M)
- An inert platinum electrode in contact with both H₂ gas and H⁺ ions
By convention, the SHE is assigned an arbitrary value of 0.00 volts. When connected to another half‑cell, the standard electrode potential of that half‑cell is read directly from a high‑resistance voltmeter.
🔬 DIAGRAM 1: Standard Hydrogen Electrode (SHE) – platinum electrode in 1 M H⁺ solution with H₂ gas bubbling at 1 atm, maintained at 298 K.
[Illustration: glass vessel with Pt wire, H₂ inlet, and salt bridge connection]
📐 Measuring Electrode Potentials Using SHE
If a hydrogen electrode is used to measure the electrode potentials of zinc and copper half‑reactions, the conventional cell diagrams are:
The hydrogen electrode is always written on the left‑hand side by convention. The polarity of the other half‑cell is measured relative to SHE. Standard electrode potential tables list standard reduction potentials (all half‑reactions written as reductions) measured at 298.15 K (25 °C).
📊 Standard Electrode Potential (E°) – Definition & Conditions
Standard electrode potential (E°) of a half‑cell is the potential measured against SHE under standard conditions:
- Temperature: 298 K (25 °C)
- Pressure: 1 atmosphere (1 atm)
- Concentration: 1 mol dm⁻³ (1 M) for all aqueous species
The standard reduction potential and oxidation potential of SHE are 0.00 V at 298 K. This zero value forms the basis of the thermodynamic scale for redox potentials.
📋 Table of Standard Reduction Potentials (298 K)
| Oxidised species + ne⁻ | ⇌ Reduced species | \(E^\theta\) (V) |
|---|---|---|
| Li⁺(aq) + e⁻ | ⇌ Li(s) | –3.04 |
| Al³⁺(aq) + 3e⁻ | ⇌ Al(s) | –1.66 |
| Pb²⁺(aq) + 2e⁻ | ⇌ Pb(s) | –0.13 |
| Fe³⁺(aq) + e⁻ | ⇌ Fe²⁺(aq) | +0.77 |
| F₂(g) + 2e⁻ | ⇌ 2F⁻(aq) | +2.87 |
🔍 Interpreting Electrode Potential Values
The electrode potential of a half‑cell indicates how easily the species undergo reduction or oxidation. Consider the half‑reaction:
Cu²⁺ is reduced (gains electrons), while Cu(s) is oxidised (loses electrons). The sign and magnitude of \(E^\theta\) reflect the tendency toward reduction.
- Greater tendency to lose electrons (oxidise)
- Equilibrium lies toward the left (favours oxidised form)
- Stronger reducing agent
- Greater tendency to gain electrons (reduce)
- Equilibrium lies toward the right (favours reduced form)
- Stronger oxidising agent
⚡ Electrochemical Series & Applications
The electrochemical series is a list of half‑reactions arranged in order of decreasing (or increasing) standard electrode potentials. It allows us to:
- Predict the spontaneity of redox reactions: a reaction is spontaneous if \(E^\theta_{\text{cell}} > 0\).
- Determine the direction of electron flow in a galvanic cell.
- Compare relative strengths of oxidising and reducing agents.
A positive \(E^\theta_{\text{cell}}\) indicates a spontaneous reaction under standard conditions.
🧮 Standard Cell Potential Calculator
🧪 Limitations & Practical Considerations
In practice, the SHE is rarely used due to several limitations:
- The electrode reaction is slow.
- The setup is not easily portable.
- Maintaining a constant pressure of hydrogen gas can be challenging.
However, once the standard electrode potential of a half‑cell is established relative to SHE, that electrode (e.g., calomel electrode or Ag/AgCl) can serve as a secondary reference. This allows other electrode potentials to be measured more conveniently using a known reference.
🔋 DIAGRAM 2: Galvanic cell – Zn (anode) and Cu (cathode) connected through a salt bridge. Voltmeter shows cell potential.
[Zn electrode in Zn²⁺ solution, Cu electrode in Cu²⁺ solution, electrons flow from Zn to Cu]
📚 Real‑World Applications
- Batteries: Design of voltaic cells (e.g., lead‑acid, lithium‑ion) relies on selecting half‑cells with appropriate E° values.
- Corrosion prevention: More negative E° metals (like Zn) are used as sacrificial anodes to protect iron (E° = –0.44 V for Fe²⁺/Fe).
- Electrolysis: Predicting which species will be oxidised or reduced at electrodes.
- Analytical chemistry: Potentiometric titrations and ion‑selective electrodes (e.g., pH meter).
📘 English lecture available – toggle above to watch.
- Standard electrode potential (E°) is measured relative to SHE (0.00 V).
- More negative E° → stronger reducing agent; more positive E° → stronger oxidising agent.
- Cell potential \(E^\theta_{\text{cell}} = E^\theta_{\text{cathode}} – E^\theta_{\text{anode}}\).
- Electrochemical series predicts spontaneity and redox behaviour.
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