Haan Research Lab

Electrochemistry

Electrochemical methods utilize the controlled flow of electrons to: produce electricity from chemical reactions, drive chemical reactions, and perform analytical measurements of chemical phenomena.

Energy production

The combination of an oxidation half reaction and a reduction half reaction can produce energy spontaneously.  If the overall reaction is carefully controlled, the energy can be used on demand (e.g., galvanic cell, battery, fuel cell).

chemical reactions

A non-spontaneous electrochemical reaction can be driven by forced movement of electrons induced by an applied voltage.  Aluminum production via such an electrolysis method (reduction of molten aluminum oxide) consumes nearly 5% of the electricity that is generated in the United States.

analytical measurements

The flow of electrons can be carefully controlled to study an individual half reaction in an electrochemical cell.  Techniques such as voltammetry, amperometry, and potentiometry, can be used to study the interaction between catalyst electrode surfaces and electrochemically active species in solution.  A fundamental understanding of these interactions leads to rational catalyst design for efficient and stable energy production and fuel synthesis.

Energy Production

The combination of an oxidation half reaction and a reduction half reaction can produce energy spontaneously.

galvanic cell

When redox reactions are combined to a favorable overall standard potential, a spontaneous reaction occurs as in the galvanic cell shown above where zinc solid is oxidized and copper ions are reduced.  The salt bridge keeps the solutions separated while allowing ion migration to maintain charge neutrality.

The following video explains the concepts behind galvanic cells:

battery

A dry cell battery (e.g., AAA, AA, C, D) is used to provide energy for numerous portable devices.  Like the galvanic cell, the cathode and anode are separated to prevent reactant mixing, but the paste between them allows ion (i.e., electrolyte) migration.  The dry cell is so named because it contains no free-flowing solutions.

An explanation of the dry cell and comparison to a wet cell (e.g., car battery) is shown here.

A lithium-ion battery is not a dry cell since it contains liquid electrolytes.  If the separator between these flammable liquids fails, the results can be spectacular.

fuel cell

Fuel cells combine an anode fuel (e.g., hydrogen, ethanol, formate) with a cathode oxidant (typically oxygen from air) to produce energy spontaneously .  While a dry cell battery must be discarded when the reactants are consumed, the fuel compartment in a fuel cell is easily replenished with fuel and the fuel cell continues to produce power.

The Department of Energy gives a summary and diagram of a hydrogen fuel cell here.

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Chemical Reactions

A non-spontaneous electrochemical reaction is driven by forced movement of electrons induced by an applied voltage.

Electrolysis

When two half reactions are combined that do not lead to spontaneous production of energy, electric potential must be applied for the overall redox reaction to proceed.  In the example electrolytic cell above, molten NaCl is converted into chlorine gas and sodium metal via electrolysis.  

The following video compares and contrasts galvanic cells and electrolytic cells:

Production of aluminum

Approximately 5% of electricity consumption in the U.S. is devoted to electrolytic production of aluminum via reduction of alumina (aluminum oxide, Al2O3) at the electrolytic cathode and oxidation of the carbon anode.

The following video demonstrates the electrolytic production of aluminum:

conversion of co2 to fuels

Many spontaneous reactions convert fossil fuels into CO2 (e.g., driving a car).  But CO2 is a stable greenhouse gas molecule  present at the highest concentration in the atmosphere now than anytime in the past 800,000 years.  Conversion of CO2 into fuel would limit growth of greenhouse gas emissions and produce fuels renewably.

The following video shows one of many electrochemical reduction processes presently under consideration to convert CO2 to fuel:

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Analytical Measurements

Reduction and oxidation half reactions can be isolated in an electrochemical cell for more fundamental analysis.  Specifically, one can probe the interactions between catalyst electrode surfaces and electrochemically active species in solution in order to improve the efficiency of energy production or electrolysis processes.

voltammetry

In voltammetry, the flow of current is measured as the applied potential is varied.  When current and potential are plotted, the size and shape of the peaks indicates information about the electrode surface and/or the electrochemically active solution species.  For example, the graph above shows the deactivation of a Pd surface by oxygen (E > 0.5 V) during the oxidation of formic acid.  The current is also directly proportional to analyte concentration, so measurement of current at a specific potential can be used for standard calibration.

The University of Cambridge gives an excellent overview of voltammetry here.

amperometry

In amperometry, an applied potential is maintained constant for a period of time and current is measured for the duration of this time period.  The rate of current decay provides information about the electrode surface efficiency and stability as well as the diffusion limitations of the solution.  The amperometric current is also directly proportional to analyte concentration, so the magnitude of current at a specific time can be used for standard calibration.

Cambridge also gives an excellent overview of amperometry here (they call it "potential step voltammetry").

potentiometry

In potentiometry, no current flows, but the voltage is measured as a potential for current to flow.  For example, a pH meter uses this potential for current flow (also known as open circuit potential).  In a pH meter, the potential for hydrogen ions to flow across a glass membrane between pH electrode and analyte solution is measured to determine the solution concentration of hydrogen ions against a standard calibration (i.e., pH buffers).

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