"The data values of standard electrode potentials (E°) are given in the table below, in volts relative to the standard hydrogen electrode, and are for the following conditions..." - Quote: "A temperature of 298.15 K (25.00 °C; 77.00 °F)." - Quote: "An effective concentration of 1 mol/L for each aqueous species or a species in mercury amalgam (an alloy of mercury with another metal)." - Quote: "A partial pressure of 101.325 kPa (absolute) (1 atm, 1.01325 bar) for each gaseous reagent." - Quote: "An activity of unity for each pure solid, pure liquid, or for water (solvent)."
The electrochemical series is a list of half-cells arranged in order of their standard electrode potentials.
Redox Reactions: Understanding the concept of oxidation-reduction reactions which involve the transfer of electrons from one species to another.
Galvanic Cells: Understanding the functioning of electrochemical cells that involve the conversion of chemical energy into electrical energy.
Half-Reactions: The concept of half-reactions involves breaking down an overall redox reaction into two separate reactions for oxidation and reduction.
Standard Reduction Potentials: The standard reduction potential is the tendency of an element to gain electrons, measured at standard conditions.
Nernst Equation: The Nernst equation is used to calculate the potential of an electrochemical cell under non-standard conditions.
Electrolysis: The process of converting electrical energy into chemical energy through the application of an external electric field to drive an otherwise non-spontaneous redox reaction.
Electrode Potential: Understanding the potential difference between electrode and solution phases, which are a function of the particular species that is involved in the electrochemical reaction.
Concentration Cells: The concentration cell is a type of electrochemical cell where two half-cells having different solute concentrations are connected.
Corrosion: Understanding the phenomenon of corrosion, where a metal corrodes due to the reaction with another substance, usually an oxidizing agent.
Batteries: The alkaline battery is a common example of an electrochemical cell that is used to store electrical energy.
Standard Electrode Potential Series: This is the most commonly used type of electrochemical series. It is a list of the standard electrode potentials of all the half-cells in a voltaic cell. The electrode with the highest standard potential is the most reactive and will act as the anode, while the electrode with the lowest standard potential is the least reactive and will act as the cathode.
Galvanic Series: This type of series lists various metals and alloys in order of their relative nobility or reactivity in seawater. Metals that are more reactive than hydrogen, such as magnesium, are at the top of the series, while metals that are less reactive, such as gold, are at the bottom.
Activity Series: This type of series lists various metals and non-metals in order of their relative reactivity. Metals that are more reactive than hydrogen are at the top of the series, while non-metals that are less reactive than hydrogen are at the bottom. The activity series is used to predict the products of single displacement reactions.
Polarization Series: This type of series lists various electrode materials in order of their resistance to polarization. Polarization is a phenomenon that occurs when the build-up of reaction products on the surface of an electrode inhibits further reaction. Electrodes that are more susceptible to polarization are at the top of the series, while electrodes that are more resistant to polarization are at the bottom. The polarization series is used to select the most suitable electrode for a particular electrochemical application.
"This pressure is used because most literature data are still given for this value (1 atm) rather than for the current standard of 100 kPa (1 bar) presently considered in the standard state." - Quote: "This pressure is used because most literature data are still given for this value (1 atm) rather than for the current standard of 100 kPa (1 bar) presently considered in the standard state."
"Although many of the half cells are written for multiple-electron transfers, the tabulated potentials are for a single-electron transfer." - Quote: "Although many of the half cells are written for multiple-electron transfers, the tabulated potentials are for a single-electron transfer."
"All of the reactions should be divided by the stoichiometric coefficient for the electron to get the corresponding corrected reaction equation." - Quote: "All of the reactions should be divided by the stoichiometric coefficient for the electron to get the corresponding corrected reaction equation."
"After dividing by the number of electrons, the standard potential E° is related to the standard Gibbs free energy of formation ΔGf° by: where F is the Faraday constant." - Quote: "After dividing by the number of electrons, the standard potential E° is related to the standard Gibbs free energy of formation ΔGf° by: where F is the Faraday constant."
"For example, in the equation Fe2+ + 2 e− ⇌ Fe(s) (–0.44 V),..." - Quote: "For example, in the equation Fe2+ + 2 e− ⇌ Fe(s) (–0.44 V),..."
"The Gibbs energy required to create one neutral atom of Fe(s) from one Fe2+ ion and two electrons is 2 × 0.44 eV = 0.88 eV..." - Quote: "The Gibbs energy required to create one neutral atom of Fe(s) from one Fe2+ ion and two electrons is 2 × 0.44 eV = 0.88 eV..."
"The Nernst equation will then give potentials at concentrations, pressures, and temperatures other than standard." - Quote: "The Nernst equation will then give potentials at concentrations, pressures, and temperatures other than standard."
"Note that the table may lack consistency due to data from different sources." - Quote: "Note that the table may lack consistency due to data from different sources."
"Calculating the potential using Gibbs free energy (E3 = 2E2 – E1) gives the potential for E3 as 0.154 V, not the experimental value of 0.159 V." - Quote: "Calculating the potential using Gibbs free energy (E3 = 2E2 – E1) gives the potential for E3 as 0.154 V, not the experimental value of 0.159 V."