standard reduction potential

Examples of standard reduction potential in the following topics:

• Standard Reduction Potentials

  • Standard reduction potentials provide a systematic measurement for different molecules' tendency to be reduced.
  • The standard reduction potential is defined relative to a standard hydrogen electrode (SHE) reference electrode, which is arbitrarily given a potential of 0.00 volts.
  • Since the reduction potential measures the intrinsic tendency for a species to undergo reduction, comparing standard reduction potential for two processes can be useful for determining how a reaction will proceed.
  • Historically, many countries, including the United States and Canada, used standard oxidation potentials rather than reduction potentials in their calculations.
  • These are simply the negative of standard reduction potentials, so it is not a difficult conversion in practice.

• Thermodynamics of Redox Reactions

  • The thermodynamics of redox reactions can be determined using their standard reduction potentials and the Nernst equation.
  • The Nernst equation allows the reduction potential to be calculated at any temperature and concentration of reactants and products; the standard reaction potential must be measured at 298K and with each solution at 1M.
  • This equation allows the equilibrium constant to be calculated just from the standard reduction potential and the number of electrons transferred in the reaction.
  • The relationship between the Gibbs free energy change and the standard reaction potential is:
  • Translate between the equilibrium constant/reaction quotient, the standard reduction potential, and the Gibbs free energy change for a given redox reaction

• Free Energy and Cell Potential

  • Electricity is generated due to the electric potential difference between two electrodes.
  • In electrochemistry, the standard electrode potential, abbreviated E°, is the measure of the individual potential of a reversible electrode at standard state, which is with solutes at an effective concentration of 1 M, and gases at a pressure of 1 atm.
  • Since the standard electrode potentials are given in their ability to be reduced, the bigger the standard reduction potentials, the easier they are to be reduced; in other words, they are simply better oxidizing agents.
  • For example, F2 has a potential of 2.87 V and Li+ has a potential of -3.05 V.
  • In the example of Zn2+, whose standard reduction potential is -0.76 V, it can be oxidized by any other electrode whose standard reduction potential is greater than -0.76 V and can be reduced by any electrode with standard reduction potential less than -0.76 V.

• Predicting if a Metal Will Dissolve in Acid

  • A metal is soluble in acid if it displaces H2 from solution, which is determined by the metal's standard reduction potential.
  • These values can be determined using standard reduction potentials, which can often be looked up.
  • Using the standard reduction potentials of a reaction, one can determine how likely a given metal is to accept or donate electrons.
  • Set up the oxidation and reduction half-reactions with their cell potential:
  • Predict whether a metal will dissolve in acid, given its reduction potential

• Electrolytic Properties

  • In order to determine which species in solution will be oxidized and which will be reduced, the standard electrode potential of each species may be obtained from a table of standard reduction potentials, a small sampling of which is shown here:
  • Historically, oxidation potentials were tabulated and used in calculations, but the current standard is to only record the reduction potential in tables.
  • If a problem demands use of oxidation potential, it may be interpreted as the negative of the recorded reduction potential.
  • This is the standard reduction potential for the reaction shown, measured in volts.
  • Use a table of standard reduction potentials to determine which species in solution will be reduced or oxidized.

• Equilibrium Constant and Cell Potential

  • In electrochemistry, the Nernst equation can be used, in conjunction with other information, to determine the equilibrium reduction potential of a half-cell.
  • In order to calculate the standard potential, we have to look up the half-reactions of copper and zinc.
  • The standard cell potential for the reaction is then +0.34 V - (-0.76 V) = +1.10 V.
  • Therefore, the standard reduction potential for zinc is more negative than that of copper.
  • In this equation, E is the cell potential, Eo is the standard cell potential (i.e., measured under standard conditions), F is Faraday's constant, R is the universal gas constant, T is the temperature in degrees Kelvin, Q is the reaction quotient (which has the same algebraic from as the equilibrium constant expression, except it applies to any time during the reaction's progress), and n is the number of moles of electrons that are transferred in the balanced chemical equation of the redox process.

• Electrochemical Cell Notation

  • Recall that standard cell potentials can be calculated from potentials E0cell for both oxidation and reduction reactions.
  • A positive cell potential indicates that the reaction proceeds spontaneously in the direction in which the reaction is written.
  • Conversely, a reaction with a negative cell potential proceeds spontaneously in the reverse direction.
  • Recall that oxidation takes place at the anode and reduction takes place at the cathode.
  • If no concentration or pressure is noted, the electrolytes in the cells are assumed to be at standard conditions (1.00 M or 1.00 atm and 298 K).

• Predicting Spontaneous Direction of a Redox Reaction

  • To figure this out, it is important to consider the standard electrode potential, which is a measure of the driving force behind a reaction.
  • The sign of the standard electrode potential indicates in which direction the reaction must proceed in order to achieve equilibrium.
  • What happens to the standard electrode potential when the reaction is written in the reverse direction?
  • However, what will change is the sign of the standard electrode potential.
  • Predict the direction of electron flow in a redox reaction given the reduction potentials of the two half-reactions

• The Nernst Equation

  • In electrochemistry, the Nernst equation can be used to determine the reduction potential of an electrochemical cell.
  • In electrochemistry, the Nernst equation can be used, in conjunction with other information, to determine the reduction potential of a half-cell in an electrochemical cell.
  • Find the cell potential of a galvanic cell based on the following reduction half-reactions where [Ni2+] = 0.030 M and [Pb2+] = 0.300 M.
  • First, find the electromotive force for the standard cell, which assumes concentrations of 1 M.

• Concentration of Cells

  • The standard potential of an electrochemical cell requires standard conditions for all of the reactants.
  • When reactant concentrations differ from standard conditions, the cell potential will deviate from the standard potential.
  • The Nernst equation can be used to calculate the output voltage changes in a pair of half-cells under non-standard conditions.
  • Under standard conditions, the output of this pair of half-cells is well known.
  • Discuss the implications of the Nernst equation on the electrochemical potential of a cell