Reactivity of metals in electrochemical cells
Energy changes • Chemical and fuel cells
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Definition of an electrochemical cell
An electrochemical cell consists of two different metal electrodes in contact with electrolyte solutions that contain their ions. Electrons move through the external circuit from the electrode that undergoes oxidation to the electrode that undergoes reduction. The cell produces an electromotive force (EMF) measured as the cell voltage between the two electrodes. Cell notation lists the anode on the left and the cathode on the right. Oxidation occurs at the anode and reduction occurs at the cathode. The sign and magnitude of measured voltage reflect the difference in the electrodes' tendencies to gain or lose electrons.
Reactivity series and oxidation tendency
The reactivity series ranks metals by their chemical reactivity, especially their tendency to lose electrons and form positive ions. Metals higher in the series oxidise more easily and thus act as better reducing agents. Metals lower in the series resist oxidation and act as better oxidising agents in metal–ion comparisons. Cause → effect: A more reactive metal loses electrons more readily → that metal becomes the anode in a cell. The more reactive metal produces positive ions in the solution, increasing the oxidation half-reaction rate relative to a less reactive metal.
Standard electrode potentials as quantitative reactivity data
Standard electrode potentials (E°) provide a numerical measure of each half-cell's tendency to gain electrons under standard conditions (1 mol dm⁻³, 298 K, 1 atm). More negative E° values indicate a stronger tendency to be oxidised (more reactive metal), while more positive E° values indicate a stronger tendency to be reduced. Cause → effect: A metal with a lower (more negative) E° value loses electrons more readily → pairing that metal with one having a higher E° produces a positive cell EMF equal to E°(cathode) − E°(anode).
Predicting direction of electron flow and cell EMF
Electron flow direction follows from comparing electrode potentials: electrons flow from the electrode with the lower (more negative) potential to the electrode with the higher (more positive) potential. The cell EMF equals the difference between the two standard potentials: E°cell = E°cathode − E°anode. Cause → effect: A larger difference between electrode potentials → a larger magnitude of measured cell voltage. A negative calculated E°cell indicates the reverse spontaneous direction under standard conditions.
Interpretation of experimental data
Measured cell voltages, combined with known electrode potentials or reactivity rankings, allow determination of which metal is more reactive. Direct comparisons of voltmeter readings from metal pairs permit ranking of metals by relative reactivity if conditions remain constant across experiments. Limiting factors: Differences in ion concentrations, temperature, electrode surface area and solution impurities alter measured voltages and can change apparent rankings unless corrected or controlled.
Practical factors and limiting conditions
Standard electrode potentials assume standard conditions. Deviations in concentration, temperature or pressure shift potentials according to the Nernst equation. Electrode coatings or slow kinetics can reduce observed voltages and mask true reactivity order. Cause → effect: Lower ion concentration or reaction rate → smaller measured EMF and slower attainment of equilibrium. Experimental interpretation requires consideration of these limiting factors to avoid incorrect conclusions about relative metal reactivity.
Key notes
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