Name: redox reactions and half-reaction balancing
Uploaded: 2020-09-24T15:06:52.000Z
Duration: 2 min 58 s
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Balancing Redox Equations

Comment une batterie fournit-elle de l'électricité pour faire fonctionner des appareils portables? Pourquoi un avocat coupé devient-il marron, ou le métal rouille-t-il sous l'eau? Ces processus sont régis par des types spécifiques de produits de réactions chimiques impliquant un transfert d'électrons d'un atome à un autre.L'atome perdant des électrons est oxydé, tandis que l'atome gagnant des électrons est réduit. Ces réactions sont appelées réactions d'oxydoréduction ou réactions redox et se caractérisent par des changements d'état d'oxydation pour un ou plusieurs réactifs. L'interaction entre le magnésium et l'acide chlorhydrique concentré est un exemple de réaction redox.Ici, le magnésium est oxydé en ions magnésium 2, et les protons sont réduits en hydrogène gazeux. Des équations chimiques simples peuvent être facilement équilibrées. Cependant, comme la plupart des équations redox sont complexes, le nombre d'électrons perdus et gagnés doit être justifié.En équilibrant les équations d'oxydoréduction, il est indispensable d'observer la loi de conservation de masse. La quantité de chaque élément et tout gain ou perte d'électrons doit être équilibré des deux côtés de la réaction;ainsi, si un réactif est réduit, un autre doit être oxydé. La méthode des demi-réactions est utilisée pour équilibrer les équations redox avec succès.Considérez la réaction du permanganate avec l'oxalate. Pour les solutions aqueuses acides, attribuez les états d'oxydation et divisez la réaction en ses demi-réactions de composant. Ensuite, équilibrez chaque demi-réaction en omettant l'hydrogène et l'oxygène.Ici, le manganèse est déjà équilibré, mais l'oxalate a besoin d'un coefficient de deux;deuxièmement, équilibrez les atomes d'oxygène avec l'ajout de molécules d'eau;et troisièmement, équilibrez les atomes d'hydrogène en ajoutant des protons si nécessaire. Équilibrez les charges en ajoutant des électrons. Puisque le manganèse est réduit, cinq électrons sont ajoutés sur le côté réactif.Inversement, l'oxalate est oxydé;ainsi, deux électrons sont ajoutés sur le côté produit. Multipliez les demi-réactions par un entier donc le nombre d'électrons ajoutés est égal. Enfin, ajoutez et simplifiez les demi-réactions équilibrées en annulant les espèces des deux côtés pour obtenir une réaction redox équilibrée.Pour les solutions aqueuses basiques, la procédure est similaire mais comprend une étape supplémentaire. Ici, un nombre égal d'ions hydroxyles est ajouté aux deux côtés de la réaction pour neutraliser les protons avant d'additionner les demi-réactions équilibrées.

Équilibrer une réaction d'oxydoréduction

L'électrochimie est la science impliquée dans l'interconversion des réactions électriques et chimiques. De telles réactions sont appelées réactions d'oxydoréduction, ou réactions redox. Ces réactions importantes sont définies par des changements de l'état d'oxydation d'un ou de plusieurs éléments réactifs et comprennent un sous-ensemble de réactions impliquant le transfert d'électrons entre les espèces de réactifs. L'électrochimie en tant que domaine a évolué pour donner suffisamment d'informations sur les principes fondamentaux de la chimie redox et de multiples technologies allant des procédés métallurgiques à l'échelle industrielle aux batteries robustes et rechargeables pour les véhicules électriques. Comme les réactions impliquant le transfert d'électrons sont essentielles à l'étude de l'électrochimie, un bref récapitulatif de la chimie redox comprend ce qui suit. 
 Nombre d'oxydation 
 Par définition, une réaction d'oxydoréduction implique un changement du nombre d'oxydation ou du degré d'oxydation d'un ou plusieurs réactifs. Le nombre d'oxydation d'un élément est une évaluation de la façon dont l'environnement électronique de ses atomes diffère par rapport aux atomes de l'élément pur. D’après cette définition, un atome dans un élément porte un nombre d'oxydation de zéro. Pour un atome, le nombre d'oxydation est égal à la charge de l'atome dans le composé si le composé est ionique. Ainsi, la somme des nombres d'oxydation de tous les atomes d'une molécule est égale à la charge de la molécule. 
 Composés ioniques 
 Les composés ioniques simples sont les exemples les plus simples de ce formalisme puisque les éléments ont des nombres d'oxydation égaux à leurs charges ioniques. Le chlorure de sodium, NaCl, est composé de cations Na+ et d'anions Cl−, les nombres d'oxydation pour le sodium et le chlore étant respectivement +1 et −1. Le fluorure de calcium, CaF2, est composé de cations Ca2+ et d'anions F−, avec des nombres d'oxydation pour le calcium et le fluor de +2 et −1. 
 Composés covalents 
 Les composés covalents sont plus compliqués en ce qui concerne l'utilisation du formalisme. L'eau est un composé covalent constitué de deux atomes H liés à un atome O via des liaisons O−H covalentes polarisées. Les électrons communs qui forment une liaison O−H sont plus fortement attirés par l'atome O plus électronégatif. Ainsi, l'oxygène acquiert une charge partielle négative, par rapport à un atome O de l'oxygène élémentaire. Par conséquent, les atomes H d'une molécule d'eau présentent une charge partielle positive par rapport aux atomes d'hydrogène du gaz hydrogène. La somme des charges partielles négatives et positives pour chaque molécule d'eau est nulle, ce qui rend la molécule d'eau neutre. 
 Si la polarisation des électrons communs à l'intérieur des liaisons O−H de l'eau était complète, le résultat serait le transfert total des électrons de H à O, et l'eau serait un composé ionique composé d'anions O2− et de cations H+. Ainsi, les nombres d'oxydations de l'oxygène et de l'hydrogène dans l'eau sont respectivement −2 et +1. En appliquant cette même logique au tétrachlorure de carbone, CCl4, cela donne des nombres d'oxydation de +4 pour le carbone et de −1 pour le chlore. Dans l'ion nitrate NO3−, le nombre d'oxydation est de +5 pour l'azote et celui de l'oxygène est de −2, ce qui s'additionne pour être égal à la charge de −1 sur la molécule : 
 <img alt="Eq1" src="/files/ftp_upload/11432/11432_Equation_1.png" style="text-align: center; height: 70px;" />
 Equilibrer des équations redox 
 L'équation déséquilibrée ci-dessous décrit la décomposition du chlorure de sodium : 
 <img alt="Eq2" src="/files/ftp_upload/11432/11432_Equation_2.png" style="text-align: center; height: 48px;" />
 Cette réaction correspond au critère d'une réaction d'oxydoréduction, car le nombre d'oxydation du Na diminue de +1 à 0 (en subissant une réduction) et celui du Cl augmente de −1 à 0 (en subissant une oxydation). Le cas de cette équation est facilement équilibré en ajoutant le coefficient stœchiométrique de 2 pour le NaCl et Na : 
 <img alt="Eq3" src="/files/ftp_upload/11432/11432_Equation_3.png" style="text-align: center; height: 48px;" />
 Les réactions d'oxydoréduction qui se produisent en solutions aqueuses se rencontrent couramment en électrochimie, et beaucoup impliquent l'eau ou ses ions, H+ (aq) et OH− (aq), soit comme réactifs soit comme produits. 
 Dans ces cas, les équations représentant les réactions redox peuvent être très difficiles à équilibrer par simple inspection, et l'utilisation d'une approche systématique connue sous le nom de la méthode des demi-équations est utile. Cette approche implique les étapes suivantes : 
 <ol>
 <li>Divisez l'équation en demi-réactions d'oxydation et de réduction composant son squelette. </li>
 <li>Équilibrer chaque demi-équation pour tous les éléments autres que O et H. </li>
 <li>Équilibrer chaque demi-équation pour les atomes O en ajoutant des molécules d'eau selon les besoins dans l'équation. </li>
 <li>Équilibrer chaque demi-équation pour les atomes H en ajoutant des protons comme requis pour l'équation. </li>
 <li>Enfin, équilibrez les charges des éléments en ajoutant des électrons selon les besoins. </li>
 <li>Multiplier les demi-équations par tout entier qui est nécessaire pour égaliser le nombre d'électrons perdus dans la demi-réaction d'oxydation par le nombre d'électrons acquis dans la demi-réaction de réduction. </li>
 <li>Ajoutez les deux demi-équations et simplifiez davantage en annulant les espèces communes des deux côtés de l'équation. </li>
 <li>Si la réaction se produit dans un milieu alcalin, ajouter des ions OH− à l'équation obtenue à l'étape 7 pour neutraliser les protons (ajoutés en nombres égaux des deux côtés de l'équation) et simplifiez. </li>
 <li>Vérifiez l'équation pour vous assurer que les charges sur les atomes sont équilibrées. </li>
 </ol>
 Ce texte est adapté de <a href="https://openstax.org/books/chemistry-2e/pages/17-1-review-of-redox-chemistry"> Openstax, Chemistry 2e, Chapter 17 : Introduction,</a> et <a href="https://openstax.org/books/chemistry-2e/pages/17-1-review-of-redox-chemistry"> Openstax, Chemistry 2e, Section 17.2 : Récapitulatif sur la chimie rédox</a>.

L’électrochimie est la science impliquée dans l’interconversion des réactions électriques et chimiques. De telles réactions sont appelées réduction-oxydation, ou ...

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redox reactions and half-reaction balancing

Redox reactions involve electron transfer and changes in oxidation number. In electrochemistry, these reactions connect chemical changes with electrical energy. They also power technologies such as industrial metal processing and rechargeable batteries used in electric vehicles.

Oxidation number is a way to track how an atom's electronic environment changes compared with the pure element, where the oxidation number is zero. For a simple ionic compound, the oxidation number matches the ion charge. Sodium chloride has Na at +1 and Cl at -1, and calcium fluoride has Ca at +2 and F at -1.

Covalent compounds need a closer look because electrons are shared. In water, the O-H bonds are polar, so oxygen pulls shared electrons more strongly than hydrogen does. That gives oxygen an oxidation number of -2 and hydrogen +1. The same idea gives carbon +4 and chlorine -1 in carbon tetrachloride, and nitrogen +5 with oxygen -2 in the nitrate ion, which adds up to the ion's -1 charge.

A simple redox example is the decomposition of sodium chloride. Sodium changes from +1 to 0, so it is reduced. Chlorine changes from -1 to 0, so it is oxidized. The equation is balanced by placing a coefficient of 2 in front of NaCl and Na.

Redox reactions in aqueous solution are often harder to balance because water, H+ ions, or OH- ions can appear as reactants or products. A systematic half-reaction method helps. First, split the equation into oxidation and reduction half-reactions. Then balance all elements except O and H, add water for oxygen, add protons for hydrogen, and add electrons to balance charge. After that, multiply the half-reactions so the electrons lost and gained are equal, add them together, and cancel species that appear on both sides.

If the reaction occurs in an alkaline medium, add OH- ions to neutralize the protons and simplify the equation. Finish by checking that both atoms and charge are balanced. This method gives a clear path for balancing redox equations in electrochemistry.

Redox Reactions and Half-Reaction Balancing

redox-reactions-and-half-reaction-balancing

balancing redox equations

Redox reactions involve electron transfer and changes in oxidation number. In electrochemistry, these reactions connect chemical changes with electrical energy. They also power technologies such as industrial metal processing and rechargeable batteries used in electric vehicles. Oxidation number is a way to track how an atom's electronic environment changes compared with the pure element, where the oxidation number is zero. For a simple ionic compound, the oxidation number matches the ion...

electromotive force

Cell Voltage and Electron Flow

Electric current is the flow of electric charge through a conducting medium or a solution. In wires, that charge is carried by electrons. In some solutions, ions can also move and carry charge. When electrons move through a wire, they produce an electric current. This idea connects to redox reactions, where electrons are lost and gained. In the spontaneous reaction between zinc and copper ions, zinc is oxidized to zinc ions. Copper ions are reduced to solid copper. Even so, this direct...

voltaic/galvanic cells

Galvanic Cells and Battery Reactions

Galvanic cells use spontaneous redox reactions to produce electric current. A common example is a disposable AA battery, which powers devices by turning chemical energy into electrical energy. Another example is a copper wire placed in silver nitrate solution, where the solution slowly turns bright blue and a grey solid forms on the wire. In the copper and silver reaction, copper is oxidized and silver ions are reduced. Copper forms cuprous ions, which give the solution its blue color. At the...

standard electrode potentials

Using Electrode Potentials to Predict Reactions

Electrode potentials help predict whether a redox reaction will happen on its own. A redox reaction involves electron transfer. In electrochemistry, the measured value is called cell potential or cell voltage. It shows the energy linked to charge transfer and is measured in volts, where 1 volt equals 1 joule per coulomb. A single half-cell cannot be measured by itself because electron transfer needs both a donor and a receiver. For that reason, scientists measure the difference between two...

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Redox Cell Potential, ΔG°, and K

Redox cell potential connects chemical energy, free energy change, and the equilibrium constant in an electrochemical cell. In these reactions, chemical energy is converted into electrical energy. That link helps predict whether a reaction is spontaneous and how far it may proceed. Thermodynamics is the study of how heat and other forms of energy are related. In an electrochemical cell, cell potential can also reflect oxidizing strength or reducing strength. A similar idea appears in acid-base...

the nernst equation

Cell Potential Under Changing Conditions

Cell potential changes when a redox reaction does not run at standard conditions. Standard cell potentials connect with standard free energy change, ΔG°, and the equilibrium constant K, but those relationships only hold at 1 M concentration, 298 K, and 1 atmosphere. A zinc(II) and tin(II) redox cell with E°cell = +0.291 V and ΔG° = −56.2 kJ is spontaneous under those standard conditions. As a cell discharges, the reactant concentration changes and the cell potential steadily drops. Many redox...

concentration cells

Cell Potential from Concentration Differences

Concentration cells are voltaic cells that use two almost identical half-cells. The half-cells use the same electrode and the same half-reaction. They differ only in the concentration of one redox species, so the cell potential depends on that concentration difference. In a concentration cell, the Nernst equation can be used to find the cell potential. For the zinc example shown, the calculated cell potential is positive. A positive cell potential means the overall reaction is spontaneous. <br...

batteries and fuel cells

Primary, Rechargeable Batteries and Fuel Cells

Batteries and fuel cells are galvanic cells that supply electrical power for many applications. They range from tiny button batteries in wristwatches to large systems that can provide backup energy to municipal power grids. Some are primary cells, which are single-use and cannot be recharged. Others are secondary cells, which use reversible reactions and can be charged again by an external power source. A common primary battery is the dry cell. It uses zinc as both the container and the anode,...

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Rusting and Metal Protection

Corrosion is the natural breakdown of metals through electrochemical reactions. It can appear as rust on iron, tarnish on silver, or a blue-green patina on copper. In each case, the metal is oxidized. Some corrosion products can protect the surface, while others keep damaging the metal. Rusting is the harmful type of corrosion seen in iron. It happens when iron is exposed to oxygen and water. A galvanic cell forms on the iron surface, and this drives oxidation at the anodic region and...

electrolysis

Electrolysis in Batteries and Metal Production

Electrolysis uses electrical energy to drive a nonspontaneous redox reaction. In a galvanic cell, a spontaneous redox reaction does electrical work on the surroundings by sending electrons through an external circuit. In electrolysis, an external circuit does work on the redox system by applying a voltage large enough to force the reaction in the opposite direction. This process is used in several real-world settings. Recharging a battery is one example, because an outside power source drives...

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Learn how oxidation numbers, redox changes, and the half-reaction method help balance redox equations in aqueous and alkaline solutions.

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Use models to interpret the data presented in Bohr diagrams and electron dot diagrams for neutral atoms of elements 1 through 18.

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Halogens

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Noble Gases


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Covalent Bonding and Lewis Structures

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Construct an explanation to classify the type of chemical bond that occurs (covalent, ionic, or metallic) in a given substance.

PS.PSc.2.1

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Structure and Nomenclature of Ethers

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Use models to apply International Union of Pure and Applied Chemistry (IUPAC) conventions to name and write formulas for simple compounds.

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Understand types, properties, and structure of matter.

Redox Reactions and Half-Reaction Balancing

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