A Latimer diagram of a chemical element is a summary of the standard electrode potential data of that element. This type of diagram is named after Wendell Mitchell Latimer, an American chemist.
A Latimer diagram of a chemical element is a summary of the standard electrode potential data of that element. This type of diagram is named after Wendell Mitchell Latimer, an American chemist.
In a Latimer diagram, the most highly oxidized form of the element is on the left side, with successively lower oxidation states to the right side. The species are connected by arrows, and the numerical value of the standard potential (in volts) for the reduction is written at each arrow. For example, for oxygen, the species would be in the order O2 (0), H2O2 (–1), H2O (-2):
The arrow between O2 and H2O2 has a value +0.68 V over it, it indicates that the standard electrode potential for the reaction:
is 0.68 volts.
Latimer diagrams can be used in the construction of Frost diagrams, as a concise summary of the standard electrode potentials relative to the element. Since ΔrGo = -nFEo, the electrode potential is a representation of the Gibbs energy change for the given reduction. The sum of the Gibbs energy changes for subsequent reductions (e.g. from O2 to H2O2, then from H2O2 to H2O) is the same as the Gibbs energy change for the overall reduction (i.e. from O2 to H2O), in accordance with Hess's law. This can be used to find the electrode potential for non-adjacent steps, which gives all the information necessary for the Frost diagram.
A simple examination of a Latimer diagram can also indicate if a species will disproportionate in solution under the conditions for which the electrode potentials are given: if the potential to the right of the species is higher than the potential on the left, it will disproportionate. Therefore, hydrogen peroxide is unstable and will disproportionate (see diagram above).
Electrochemistry is the branch of physical chemistry concerned with the relationship between electrical potential difference and identifiable chemical change. These reactions involve electrons moving via an electronically-conducting phase between electrodes separated by an ionically conducting and electronically insulating electrolyte.
In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. The voltage that is needed for electrolysis to occur is called the decomposition potential. The word "lysis" means to separate or break, so in terms, electrolysis would mean "breakdown via electricity".
In electrochemistry, electrode potential is the electromotive force of a galvanic cell built from a standard reference electrode and another electrode to be characterized. By convention, the reference electrode is the standard hydrogen electrode (SHE). It is defined to have a potential of zero volts. It may also be defined as the potential difference between the charged metallic rods and salt solution.
In chemistry, a half reaction is either the oxidation or reduction reaction component of a redox reaction. A half reaction is obtained by considering the change in oxidation states of individual substances involved in the redox reaction. Often, the concept of half reactions is used to describe what occurs in an electrochemical cell, such as a Galvanic cell battery. Half reactions can be written to describe both the metal undergoing oxidation and the metal undergoing reduction.
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:
Redox is a type of chemical reaction in which the oxidation states of substrate change. Oxidation is the loss of electrons or an increase in the oxidation state, while reduction is the gain of electrons or a decrease in the oxidation state.
In electrochemistry, the Nernst equation is a chemical thermodynamical relationship that permits the calculation of the reduction potential of a reaction from the standard electrode potential, absolute temperature, the number of electrons involved in the redox reaction, and activities of the chemical species undergoing reduction and oxidation respectively. It was named after Walther Nernst, a German physical chemist who formulated the equation.
In chemistry, a reducing agent is a chemical species that "donates" an electron to an electron recipient. Examples of substances that are common reducing agents include the alkali metals, formic acid, oxalic acid, and sulfite compounds.
In electrochemistry, standard electrode potential, or , is a measure of the reducing power of any element or compound. The IUPAC "Gold Book" defines it as; "the value of the standard emf of a cell in which molecular hydrogen under standard pressure is oxidized to solvated protons at the left-hand electrode".
In electrochemistry, the standard hydrogen electrode, is a redox electrode which forms the basis of the thermodynamic scale of oxidation-reduction potentials. Its absolute electrode potential is estimated to be 4.44 ± 0.02 V at 25 °C, but to form a basis for comparison with all other electrochemical reactions, hydrogen's standard electrode potential is declared to be zero volts at any temperature. Potentials of all other electrodes are compared with that of the standard hydrogen electrode at the same temperature.
The values below are standard apparent reduction potentials (E°') for electro-biochemical half-reactions measured at 25 °C, 1 atmosphere and a pH of 7 in aqueous solution.
In chemistry, disproportionation, sometimes called dismutation, is a redox reaction in which one compound of intermediate oxidation state converts to two compounds, one of higher and one of lower oxidation states. The reverse of disproportionation, such as when a compound in an intermediate oxidation state is formed from precursors of lower and higher oxidation states, is called comproportionation, also known as synproportionation.
Comproportionation or synproportionation is a chemical reaction where two reactants containing the same element but with different oxidation numbers, form a compound having an intermediate oxidation number. It is the opposite of disproportionation.
Redox potential is a measure of the tendency of a chemical species to acquire electrons from or lose electrons to an electrode and thereby be reduced or oxidised respectively. Redox potential is expressed in volts (V). Each species has its own intrinsic redox potential; for example, the more positive the reduction potential, the greater the species' affinity for electrons and tendency to be reduced.
In electrochemistry, and more generally in solution chemistry, a Pourbaix diagram, also known as a potential/pH diagram, EH–pH diagram or a pE/pH diagram, is a plot of possible thermodynamically stable phases of an aqueous electrochemical system. Boundaries (50 %/50 %) between the predominant chemical species are represented by lines. As such a Pourbaix diagram can be read much like a standard phase diagram with a different set of axes. Similarly to phase diagrams, they do not allow for reaction rate or kinetic effects. Beside potential and pH, the equilibrium concentrations are also dependent upon, e.g., temperature, pressure, and concentration. Pourbaix diagrams are commonly given at room temperature, atmospheric pressure, and molar concentrations of 10−6 and changing any of these parameters will yield a different diagram.
A Frost diagram or Frost–Ebsworth diagram is a type of graph used by inorganic chemists in electrochemistry to illustrate the relative stability of a number of different oxidation states of a particular substance. The graph illustrates the free energy vs oxidation state of a chemical species. This effect is dependent on pH, so this parameter also must be included. The free energy is determined by the oxidation–reduction half-reactions. The Frost diagram allows easier comprehension of these reduction potentials than the earlier-designed Latimer diagram, because the “lack of additivity of potentials” was confusing. The free energy ΔG° is related to the reduction potential E shown in the graph by the formula: ΔG° = −nFE° or nE° = −ΔG°/F, where n is the number of transferred electrons, and F is the Faraday constant (F = 96,485 J/(V·mol)= 96,485 Coulomb). The Frost diagram is named after Arthur Atwater Frost, who originally invented it as a way to "show both free energy and oxidation potential data conveniently" in a 1951 paper.
Potassium hypomanganate is the inorganic compound with the formula K3MnO4. Also known as potassium manganate(V), this bright blue solid is a rare example of a salt with the hypomanganate or manganate(V) anion, where the manganese atom is in the +5 oxidation state. It is an intermediate in the production of potassium permanganate and the industrially most important Mn(V) compound.
The Haber–Weiss reaction generates •OH (hydroxyl radicals) from H2O2 (hydrogen peroxide) and superoxide (•O2−) catalyzed by iron ions. It was first proposed by Fritz Haber and his student Joseph Joshua Weiss in 1932.
Thermochemical cycles combine solely heat sources (thermo) with chemical reactions to split water into its hydrogen and oxygen components. The term cycle is used because aside of water, hydrogen and oxygen, the chemical compounds used in these processes are continuously recycled.
Water oxidation is one of the half reactions of water splitting:
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