Electrochemistry

The One Thing

Electrochemistry links oxidation–reduction reactions with electrical energy: galvanic cells generate electricity from spontaneous reactions, whereas electrolytic cells use electricity to drive non-spontaneous reactions. Cell potential, ionic conductance and electrolysis are quantified by the Nernst equation, conductivity relationships and Faraday’s laws.

Reactions, Processes and Experiments

Key Terms

• Electrochemical cell: A system in which oxidation and reduction reactions are separated or controlled to produce electrical energy.
• Galvanic or voltaic cell: A cell that converts the energy of a spontaneous redox reaction into electrical energy.
• Electrolytic cell: A cell in which external electrical energy drives a non-spontaneous chemical reaction.
• Anode: The electrode at which oxidation occurs. It is negative in a galvanic cell and positive in an electrolytic cell.
• Cathode: The electrode at which reduction occurs. It is positive in a galvanic cell and negative in an electrolytic cell.
• Salt bridge: A device containing an inert electrolyte that completes the electrical circuit, maintains electrical neutrality and reduces liquid junction potential.
• Electrode potential: The tendency of an electrode to undergo oxidation or reduction when it is in contact with its ionic solution.
• Standard electrode potential: The electrode potential measured under standard conditions, usually with solute concentration of 1 mol L−1, gas pressure of 1 bar and specified temperature, commonly 298 K.
• Cell potential or EMF: The potential difference between the cathode and anode when no current is drawn from the cell.
• Nernst equation: An equation that relates electrode or cell potential to the standard potential and the reaction quotient.
• Conductance: The ability of a solution to conduct electric current; it is the reciprocal of resistance.
• Specific conductance or conductivity: The conductance of a solution contained between electrodes of unit area separated by unit distance.
• Molar conductivity: The conductance of the volume of solution containing one mole of an electrolyte, placed between electrodes separated by unit distance.
• Kohlrausch’s law: At infinite dilution, each ion contributes independently to the total molar conductivity of an electrolyte.
• Electrolysis: The chemical decomposition of an electrolyte by passing direct electric current through its molten state or aqueous solution.
• Faraday’s first law: The mass of a substance deposited or liberated at an electrode is directly proportional to the quantity of electricity passed.
• Faraday’s second law: When the same quantity of electricity passes through different electrolytes, the masses deposited are proportional to their equivalent masses.
• Primary cell: A battery that cannot be conveniently recharged because its chemical reaction is essentially irreversible.
• Secondary cell: A rechargeable battery in which the cell reaction can be reversed by passing an external current.
• Fuel cell: An electrochemical cell that continuously converts the chemical energy of externally supplied fuels and oxidants into electrical energy.
• Reaction quotient, Q: For a complete cell reaction, the activities or concentrations of products divided by those of reactants, excluding pure solids and liquids.
• Conductivity, κ: The reciprocal of resistivity and a measure of the conducting ability of an electrolyte solution.
• Cell constant: The ratio $l/A$, where $l$ is the distance between electrodes and $A$ is their area.
• Equivalent mass: The quantity used in Faraday’s second law to compare the masses deposited by the same quantity of electricity.
• Faraday constant, F: The charge associated with one mole of electrons, approximately 96485 C mol−1.

Easily Confused

• Anode and cathode: Oxidation always occurs at the anode and reduction always occurs at the cathode; their signs differ between galvanic and electrolytic cells.
• Galvanic and electrolytic cells: A galvanic cell converts a spontaneous redox reaction into electrical energy, whereas an electrolytic cell uses electrical energy to drive a non-spontaneous reaction.
• Conductance and conductivity: Conductance is $G=1/R$ and depends on the particular cell, whereas conductivity is $\kappa =G(l/A)$ and refers to a standardized geometry.
• Conductivity and molar conductivity: Conductivity generally decreases on dilution because there are fewer ions per unit volume, whereas molar conductivity generally increases because ionization increases and ionic interactions decrease.
• Strong and weak electrolytes: Strong electrolytes show a slow increase in molar conductivity on dilution, whereas weak electrolytes show a much larger increase because dilution increases their degree of ionization.
• Primary and secondary cells: A primary cell is not conveniently rechargeable because its reaction is essentially irreversible; a secondary cell can be recharged because its reaction can be reversed.
• Cell potential and standard cell potential: Cell potential depends on the actual reaction conditions, whereas standard cell potential is measured under specified standard conditions.
• Faraday’s first and second laws: The first law relates deposited mass to charge passed, whereas the second compares deposited masses for different electrolytes using equivalent masses.

What Gets Asked

• Calculate cell potential using the Nernst equation: $E=E°-(RT/nF)\ln Q$ or, at 298 K, $E=E°-(0.0591/n)\log Q$. Marks are lost by constructing $Q$ incorrectly or including pure solids and liquids.
• Calculate the potential of a metal-ion electrode using $E=E°+(0.0591/n)\log [Mn+]$. The activity of the solid metal is taken as unity.
• Determine spontaneity, Gibbs energy or the equilibrium constant using $E°cell=E°cathode-E°anode$, $\Delta G°=-nFE°cell$ and $\log K=nE°cell/0.0591$. A positive $E°cell$ corresponds to a spontaneous cell reaction under standard conditions.
• Calculate resistance, conductance, conductivity or cell constant using $V=IR$, $G=1/R$, $R=\rho l/A$, $\kappa =1/\rho =l/(RA)=G(l/A)$ and $\kappa =G\times$ cell constant. The common error is confusing conductance with conductivity.
• Explain the effect of dilution on conductivity and molar conductivity, distinguishing the slow increase for strong electrolytes from the much larger increase for weak electrolytes.
• Apply Faraday’s laws to calculate deposited mass using $m=ZIt$, $m_1/m_2=E_1/E_2$ or $m=(MIt)/(nF)$. Marks are commonly lost by using the wrong electron number $n$, equivalent mass or quantity of electricity $Q=It$.