Electrochemistry deals with the relationship between electrical and chemical energy. Standard states provide a consistent framework for measuring and comparing cell potentials, allowing us to predict the direction of electron flow and reaction spontaneity in electrochemical cells.
The extends our understanding beyond . It lets us calculate cell potentials under various concentrations, helping predict how real-world factors affect electrochemical reactions and when they reach equilibrium.
Standard States and Cell Potentials
Standard states in electrochemistry
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Standard state conditions for electrochemical systems require:
Temperature maintained at 25°C (298 K)
Gases at a pressure of 1 atm
Solutes in aqueous solutions at a concentration of 1 M
Pure substances in their most stable form at the specified temperature and pressure (25°C and 1 atm)
Standard reduction potentials (E0) measured under these well-defined standard state conditions
Reduction potentials expressed relative to the (SHE) which has an assigned potential of 0.00 V by convention
Examples of standard reduction potentials: E\ceCu2+/Cu0=+0.34 V, E\ceZn2+/Zn0=−0.76 V
Calculation of standard cell potentials
Standard cell potential (Ecell0) calculated as the difference between the standard reduction potentials of the cathode (Ecathode0) and anode (Eanode0)
Standard reduction potentials typically listed in a table in order of increasing reduction potential
Species with the most positive E0 acts as the strongest oxidizing agent and undergoes reduction at the cathode (gains electrons)
Species with the least positive (or most negative) E0 acts as the strongest reducing agent and undergoes oxidation at the anode (loses electrons)
Example: In a cell with \ceCu2+/Cu (E0=+0.34 V) and \ceZn2+/Zn (E0=−0.76 V), \ceCu2+ reduces at the cathode while \ceZn oxidizes at the anode
Electron flow in electrochemical cells
Electrons flow from the anode (site of oxidation) to the cathode (site of reduction) in an electrochemical cell
Species with the least positive (or most negative) E0 undergoes oxidation at the anode
Anode is the electrode where oxidation occurs and electrons are released
Species with the most positive E0 undergoes reduction at the cathode
Cathode is the electrode where reduction occurs and electrons are consumed
Spontaneity of the cell reaction determined by the sign of Ecell0
Positive Ecell0 indicates a spontaneous cell reaction as written ()
Negative Ecell0 indicates the reverse reaction is spontaneous ()
Examples: In a \ceZn/Cu cell, \ceZn (E0=−0.76 V) oxidizes at the anode and \ceCu2+ (E0=+0.34 V) reduces at the cathode; Ecell0=+1.10 V (spontaneous)
Nernst Equation and Non-Standard Conditions
Nernst equation for non-standard conditions
Nernst equation relates the cell potential (Ecell) to the standard cell potential (Ecell0) and the concentrations (or partial pressures) of reactants and products
Mathematical expression: Ecell=Ecell0−nFRTlnQ
R: universal gas constant (8.314 J/mol·K)
T: temperature in Kelvin (K)
n: number of electrons transferred in the balanced
F: Faraday's constant (96,485 C/mol)
Q: reaction quotient (ratio of product concentrations to reactant concentrations, each raised to their stoichiometric coefficients)
Nernst equation used to calculate cell potential at any given concentration (or partial pressure) of reactants and products
As reaction progresses, reactant concentrations decrease and product concentrations increase, causing cell potential to decrease
Example: For the cell reaction \ceZn+Cu2+−>Zn2++Cu, Ecell=Ecell0−2FRTln[\ceCu2+][\ceZn2+]
Nernst equation also used to determine concentration (or partial pressure) of a reactant or product at equilibrium when Ecell=0
Example: For the cell reaction \ceZn+Cu2+−>Zn2++Cu, at equilibrium Ecell=0 and [\ceCu2+][\ceZn2+]=eRT2FEcell0