Oxidizing agent and reducing agent react to make reduced molecule and oxidized molecule, respectively, by transferring one or more electrons {redox reaction, inorganic} {oxidation-reduction reaction}.
equilibrium constant
Find redox-reaction equilibrium constant from balanced equation. From balanced equation, separate reduction and oxidation half-reactions. Find half-reaction standard potentials and total potential. Note number of electrons transferred. Equilibrium constant equals exponential of standard potential V times number of electrons transferred n times one faraday F, divided by gas constant k times temperature T: exp(V*n*F/k*T).
Molecules can lose electrons {oxidation}. Losing electrons increases positive charge.
Molecules can gain electrons {reduction}|. Gaining electrons decreases positive charge.
If all electrons shared in chemical bonds go to the more-electronegative atom, each atom has resultant charge {oxidation number}|. Oxidation number is number of electrons added to, or subtracted from, outer shell to make full shell. In molecules, atoms with higher electronegativity tend to gain electrons from atoms with lower electronegativity.
metals
Metals have positive oxidation numbers, because they lose electrons and empty outer shell. Metals are reducers, because they themselves oxidize.
non-metals
Oxygen and fluorine have negative oxidation numbers, because they gain electrons to fill outer shell. Oxygen and fluorine are oxidizers, because they themselves reduce.
hydrogen
Hydrogen can gain or lose one electron, making oxidation number +1 or -1.
others
Atoms can have several oxidation numbers, because they can fill or empty outer shell in different ways, through orbital hybridization.
functional group
Functional-group oxidation number is atom oxidation-number sum.
Metal oxidation states in acids and bases show specific relations {Frost diagram}.
Slow metal ionization {rusting} {corrosion}| can use oxidation, from oxygen in air or acidic water. Metal impurities make circuit from impurity to metal. Corrosion rate depends on exposed area. Aluminum oxide covers aluminum and prevents further corrosion.
Zinc coatings {anodized}| protect metal, because zinc corrodes first.
Because it is electron source, magnesium can prevent corrosion by returning electrons to metal {cathode protection}.
Galvanic cells {battery, redox}| can connect to make electric current by solution chemical reactions. Batteries {dry cell battery} can use paste. Batteries {wet cell battery} can use liquid solution.
voltage
Metals used for electrodes depend on battery solution. Metal combinations make different battery voltages across electrodes. Typical voltage is one volt to three volts.
series
Electrochemical cells can connect in series, so cell voltages add. Automobile batteries use lead plates and sulfuric acid solution.
recharging
Some batteries can recharge by applied electric voltage and current.
additives
Battery additives are useless.
Some batteries {primary cell}| cannot recharge.
Some batteries {secondary cell} can recharge.
battery {storage cell}|.
Current flows from battery negative terminal through circuit and power-using device {load, circuit} to positive terminal and then solution.
Electric voltage can electrolyze or electroplate material {electrolysis}|. Electrolysis uses potential to drive electrochemical reaction. Electric current can split solute molecules into two ions. Mass deposited in electroplating, or split in electrolysis, is proportional to total charge transfer. Mass deposited in electroplating, or split in electrolysis, is proportional to atomic weight to ion charge ratio {chemical equivalent}.
calculation
Find material electrolyzed or electroplated, or charge needed to electrolyze or electroplate material amount, from balanced equation. From balanced equation, separate reduction and oxidation half-reactions. Note number of electrons transferred. Moles of electrons used are coulombs divided by 96,500 Coulombs. Coulombs used is current amperes times seconds. Ratio of electrolyzed or electroplated product coefficient to transferred-electron number is ratio of electrolyzed or electroplated product moles to electrons-used moles.
potential
Nernst potential is minimum voltage needed to reverse spontaneous reaction at given conditions. Concentration gradient at electrode surface can make overpotential. Total needed potential {decomposition potential} includes Nernst potential, overpotential, and electrical-resistance potential.
activation energy
Temperature, current-to-area ratio, electrode surface, and electrode type affect reaction activation energy.
types: constant current
Constant-current electrolysis is for metals with reduction potential greater than hydrogen and for potential greater than hydrogen decomposition potential. Hydrogen ions in high-acidity solution carry constant current, because they are much more numerous than metal ions. Substrate keeps hydrogen gas low, so gas does not cover hydrogen electrode. Constant current times time makes total charge.
types: constant voltage
Constant-voltage electrolysis keeps potential high enough to lower metal-ion concentration to optimum level but low enough to stop hydrogen-gas evolution or other-metal deposition. In this method, current decreases over time.
types: controlled potential
Controlled-potential electrolysis uses third electrode (SCE) as reference to keep oxidation potential at cathode constant, keep current high enough, and prevent unwanted reactions. Current decreases over time exponentially. Cathode potential determines decomposition potential, so, as metal deposits, ion concentration goes down, and decomposition potential goes up.
In electrolysis, material moles formed or reacted is electron moles divided by metal or other ion charge {equivalent, chemistry}|.
In oxidation-reduction reactions, electrons transfer. Electron transfer requires voltage {potential, reaction}. Oxidation-reduction half-reactions have potentials. Total oxidation-reduction reaction potential is sum of half-reaction potentials.
level
Electronegative atoms have high reduction potentials.
calculation
Cell standard potential depends on half-reaction potentials and balanced equation. From balanced equation, separate reduction and oxidation half-reactions. Tables have half-reaction reduction potentials at 25 C. Subtract reduction potential for oxidation half-reaction from reduction potential for reduction half-reaction to find chemical-reaction potential.
spontaneity
If oxidation-reduction chemical-reaction potential is greater than zero, chemical reaction is spontaneous. If reduction potential is more than oxidation potential, ionization potential is higher, hydration energy is lower, and sublimation energy is more.
equilibrium
When oxidation-reduction reaction is complete, potential is zero.
Half-cell reduction-reaction voltage depends on oxidized and reduced concentrations, temperature, and number of electrons transferred {Nernst equation}. V = V0 - R * T / (n * ln([Co] - [Cr])), where V is reaction potential, V0 equals standard unit cell potential, R is gas constant, T is temperature, n is number of electrons transferred, Co is oxidized-ion concentration, and Cr is reduced-ion concentration.
Activation energy comes from electric field and from temperature.
valence charge
Electrode voltage is V = V0 + R * T * ln(K / (z * F)), where V is voltage, V0 is standard potential, R is gas constant, T is temperature, K is equilibrium constant, z is absolute value of transferred charge {valence charge}, and F is 1 Faraday. Therefore, exp(-V / (R*T)) = K / (z * F) and K = z * F * exp(-V / (R*T)). Standard potential is at concentration 1 M, pressure 1 atmosphere, and temperature 25 C. Solids have concentration = 1. Voltage is always positive. If V > 0, reaction is spontaneous. At equilibrium, voltage = 0 and current = 0.
Molecules {oxidizing agent}| can gain electrons from another molecule. Oxygen, halogens, permanganates, and chromates are oxidizing agents.
Molecules {reducing agent}| can lose electrons to another molecule. Small, light metals are reducing agents.
Redox reactions can be in solutions {cell, redox reaction}|. Conducting plates can be in two connected half-cells. Current from one electrode goes through wire to other electrode and then through solution.
potential
Metal and metal ion have potential difference, because electrons and ions separate. If potential > 0, metal favors reduction. If potential < 0, metal favors metal oxidation.
Nernst equation
If electrode voltage is > 0, reaction is spontaneous. At equilibrium, voltage = 0 and current = 0.
tables
Tables can show reduction-half-reaction potentials. Hydrogen electrode has standard potential 0 V.
diffusion
In fast reactions, diffusion controls reaction rate. Moving electrode or stirring solution minimizes diffusion effects.
salt bridge
Agar and potassium-chloride salt bridge can connect half-cells.
membrane
Membranes that block and allow ion flows can have potential difference, because membrane sides have different ion concentrations. Membrane-permeable ion diffuses through until electrostatic repulsion from higher-concentration side stops ion flow.
Redox reaction has two parts {half-reaction}| {half-cell} that interact, reducing-agent oxidation and oxidizing-agent reduction. Cell redox reactions oxidize reducing agent and reduce oxidizing agent. For example, half-reaction for hydrogen electrode is 2 H+ + 2 e- [+ and - are superscripts] <-> 1 H2 [2 is subscript].
Potential difference produced by open circuit {electromotive force}| (emf) is voltage that can make electricity.
Devices that make electric current have resistance force {back electromotive force}| to current. Maximum battery power has battery back electromotive force equal to circuit resistance.
Voltage applied to cells {electrolytic cell}| can force current through cell and cause reverse redox reaction. In electrolytic cells, cathode is electron source, and anode is electron sink. Applying voltage and current can split molecules by electrolysis. For example, water can form hydrogen and oxygen. Aluminum salts can make aluminum.
Spontaneous redox reaction in cells {galvanic cell}| can make current. In galvanic cells, anode oxidizes and is negative, while cathode reduces and is positive. Metal or metal-oxide electrode can be in solution {wet cell} or paste {dry cell}. Different metal or metal oxide can be in same solution or paste, or another solution or paste connected to first by conductor, to make two coupled cells and a battery. Metal reacts with solution to make ions. Metal anode loses electrons and becomes positive. Metal cathode gains electrons and becomes negative. Electrodes have potential difference.
battery types
Batteries can have nickel and cadmium in acid solution. Edison cells have nickel oxide and iron electrodes in alkaline solution. Batteries can have lead and lead oxide in acid solution.
Galvanic cells {fuel cell}| can have continuous fuel supply. Fuel cells make electric current by oxidizing hydrides or other substances. Fuel cells are efficient, cool, and clean.
Electrolysis can put metal on conducting surfaces {electroplating}|. In silver plating, gold plating, and zinc plating, metal derives from salt solution. Electric current adds electrons to change ion to metal at electrode.
In electrolytic cells, oxidation is at positively charged electrode {anode, cell}|. In galvanic cells, reduction is at positively charged electrode.
In electrolytic cells, reduction is at negatively charged electrode {cathode, cell}|. In galvanic cells, oxidation is at negatively charged electrode.
5-Chemistry-Inorganic-Oxidation-Reduction
Outline of Knowledge Database Home Page
Description of Outline of Knowledge Database
Date Modified: 2022.0225