standard hydrogen electrode

Examples of standard hydrogen electrode 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.
  • The values below in parentheses are standard reduction potentials for half-reactions measured at 25 °C, 1 atmosphere, and with a pH of 7 in aqueous solution.
  • These are simply the negative of standard reduction potentials, so it is not a difficult conversion in practice.
  • Recall that a positive reduction potential indicates a thermodynamically favorable reaction relative to the reduction of a proton to hydrogen

• Electrolytic Properties

  • In other systems, the electrode reactions can involve electrode metal as well as electrolyte ions.
  • In the last example, H+ ions (hydrogen ions) also take part in the reaction, and are provided by an acid in the solution or by the solvent itself (water, methanol, etc.).
  • 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:
  • 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.

• Electrolysis of Water

  • If the object is to produce hydrogen and oxygen, the added electrolyte must be energetically more difficult to oxidize or reduce than water itself.
  • For example, electrolysis of a solution of sulfuric acid or of a salt, such as NaNO3, results in the decomposition of water at both electrodes:
  • Hydrogen will appear at the cathode, the negatively charged electrode, where electrons enter the water, and oxygen will appear at the anode, the positively charged electrode.
  • The number of moles of hydrogen generated is twice the number of moles of oxygen, and both are proportional to the total electrical charge conducted by the solution.
  • The number of electrons pushed through the water is twice the number of generated hydrogen molecules, and four times the number of generated oxygen molecules.

• 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.
  • Note that the table also takes the replacement of hydrogen (H2) into account.
  • The tendency of a metal to "replace" hydrogen gas from acidic solution will determine its solubility in that solution.
  • Therefore, the half-cell potential for the Zn/Zn2+ electrode always refers to the reduction reaction:
  • These values can be determined using standard reduction potentials, which can often be looked up.

• 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.
  • 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.

• Electrolysis of Sodium Chloride

  • This results in chemical reactions at the electrodes and the separation of materials.
  • Rather than producing sodium, hydrogen is produced.
  • Metal ions receive electrons at the negative electrode, and the non-metals lose them at the positive electrode.
  • Electrolysis of aqueous NaCl results in hydrogen and chloride gas.
  • At the cathode (C), water is reduced to hydroxide and hydrogen gas.

• Other Rechargeable Batteries

  • The electrolyte may serve as a simple buffer for internal ion flow between the electrodes, as in lithium-ion and nickel-cadmium cells, or it may be an active participant in the electrochemical reaction, as in lead-acid cells.
  • NiMH batteries use positive electrodes of nickel oxyhydroxide (NiOOH), as does the NiCd, but the negative electrodes use a hydrogen-absorbing alloy instead of cadmium.
  • The lithium-ion battery is a family of rechargeable batteries in which lithium ions move from the negative electrode to the positive electrode during discharge, and back when charging.
  • The negative electrode of a conventional lithium-ion cell is made from carbon.
  • The positive electrode is a metal oxide, and the electrolyte is a lithium salt in an organic solvent.

• Dry Cell Battery

  • A fibrous fabric separates the two electrodes, and a brass pin in the center of the cell conducts electricity to the outside circuit.
  • The manganese(IV) oxide in the cell removes the hydrogen produced by the ammonium chloride, according to the following reaction:
  • In some more modern types of so-called "high-power" batteries that have a much lower capacity than standard alkaline batteries, the ammonium chloride is replaced by zinc chloride.

• 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.

• Electrochemical Cell Notation

  • Recall that standard cell potentials can be calculated from potentials E0cell for both oxidation and reduction reactions.
  • One beaker contains 0.15 M Cd(NO3)2 and a Cd metal electrode.
  • The other beaker contains 0.20 M AgNO3 and a Ag metal electrode.
  • If the electrolytes in the cells are not at standard conditions, concentrations and/or pressure, they are included in parentheses with the phase notation.
  • 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).