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

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Balancing Redox Equations

Hoe levert een batterij elektriciteit om draagbare apparaten te laten werken? Waarom wordt een gesneden avocado bruin of roest metaal onder water? Deze processen worden aangedreven door specifieke soorten chemische reacties waarbij elektronenoverdracht van het ene atoom naar het andere betrokken is.Het atoom dat elektronen verliest, wordt geoxideerd, terwijl het atoom dat elektronen wint, wordt verminderd. Deze reacties worden oxidatie-reductie-of redoxreacties genoemd en worden gekenmerkt door veranderingen in de oxidatietoestand voor een of meer reactanten. De interactie tussen magnesium en geconcentreerd zoutzuur is een voorbeeld van een redoxreactie.Hier wordt magnesium geoxideerd tot magnesium 2 ionen, en de protonen worden gereduceerd tot waterstofgas. Eenvoudige chemische vergelijkingen kunnen gemakkelijk worden uitgebalanceerd. Omdat de meeste redoxvergelijkingen echter complex zijn, moet het aantal verloren en gewonnen elektronen worden gerechtvaardigd.Bij het balanceren van redoxvergelijkingen is het essentieel om de wet van behoud van massa in acht te nemen. De hoeveelheid van elk element en elke winst of verlies van elektronen moet aan beide kanten van de reactie in evenwicht worden gebracht;dus als een reactant wordt gereduceerd, moet een andere worden geoxideerd. De methode van halfreacties wordt gebruikt om redoxvergelijkingen met succes te balanceren.Denk aan de reactie van permanganaat met oxalaat. Wijs voor waterige zure oplossingen de oxidatietoestanden toe en verdeel de reactie in de samenstellende halfreacties. Breng vervolgens elke halve reactie in evenwicht, waarbij waterstof en zuurstof worden weggelaten.Hier is mangaan al in evenwicht, maar het oxalaat heeft een coëfficiënt van twee nodig;ten tweede, breng zuurstofatomen in evenwicht met de toevoeging van watermoleculen;en ten derde, breng waterstofatomen in evenwicht door waar nodig protonen toe te voegen. Breng de ladingen in evenwicht door elektronen toe te voegen. Omdat mangaan wordt verminderd, worden vijf elektronen toegevoegd aan de reactantzijde.Omgekeerd wordt oxalaat geoxideerd;zo worden aan de productzijde twee elektronen toegevoegd. Vermenigvuldig de halfreacties met een geheel getal, zodat het aantal toegevoegde elektronen gelijk is. Voeg ten slotte de gebalanceerde halfreacties toe en vereenvoudig ze door soorten aan beide kanten op te heffen om een gebalanceerde redoxreactie te verkrijgen.Voor basische waterige oplossingen is de procedure vergelijkbaar, maar omvat een extra stap. Hier wordt een gelijk aantal hydroxylionen aan beide zijden van de reactie toegevoegd om de protonen te neutraliseren voordat de gebalanceerde halfreacties worden opgeteld.

redox-vergelijkingen in evenwicht brengen

Elektrochemie is de wetenschap die zich bezighoudt met de onderlinge omzetting van elektrische en chemische reacties. Dergelijke reacties worden reductie-oxidatie of redoxreacties genoemd. Deze belangrijke reacties worden gedefinieerd door veranderingen in de oxidatietoestanden van een of meer reactantelementen en omvatten een subset van reacties waarbij de overdracht van elektronen tussen reactantsoorten betrokken is. Elektrochemie als vakgebied is geëvolueerd om voldoende inzichten te verschaffen over de fundamentele principes van redoxchemie en meerdere technologieën, variërend van metallurgische processen op industriële schaal tot robuuste, oplaadbare batterijen voor elektrische voertuigen. Omdat reacties waarbij elektronen worden overgedragen essentieel zijn voor de studie van elektrochemie, omvat een kort overzicht van de redoxchemie het volgende. Oxidatiegetallen Per definitie omvat een redoxreactie een verandering in het oxidatiegetal of de oxidatietoestand voor een of meerdere reactanten. Het oxidatiegetal van een element is een beoordeling van hoe de elektronische omgeving van zijn atomen verschilt in vergelijking met atomen van het pure element. Volgens deze definitie heeft een atoom in een element een oxidatiegetal van nul. Voor een atoom is het oxidatiegetal gelijk aan de lading van het atoom in de verbinding als de verbinding ionisch zou zijn. De som van de oxidatiegetallen voor alle atomen in een molecuul is dus gelijk aan de lading van het molecuul. Ionische bestanddelen 

 Eenvoudige ionische verbindingen zijn de eenvoudigste voorbeelden van dit formalisme, aangezien de elementen oxidatiegetallen hebben die gelijk zijn aan hun ionische ladingen. Natriumchloride, NaCl, is samengesteld uit Na+ kationen en Cl−anionen, waarbij de oxidatiegetallen voor natrium en chloor respectievelijk +1 en −1 zijn. Calciumfluoride, CaF2, is samengesteld uit Ca2+ kationen en F− anionen, met oxidatiegetallen voor calcium en fluor, +2 en −1. 

 Covalente verbindingen 

 Covalente verbindingen vormen een grotere uitdaging bij het gebruik van formalisme. Water is een covalente verbinding die bestaat uit twee H-atomen die via polaire covalente O-H-bindingen aan een O-atoom zijn gebonden. De gedeelde elektronen die een O-H-binding maken, worden sterker aangetrokken door het meer elektronegatieve O-atoom. Zuurstof krijgt dus een gedeeltelijke negatieve lading, vergeleken met een O-atoom in elementaire zuurstof. Als gevolg hiervan vertonen H-atomen in een watermolecuul een gedeeltelijk positieve lading vergeleken met waterstofatomen in waterstofgas. De som van de gedeeltelijke negatieve en positieve ladingen voor elk watermolecuul is nul, waardoor het watermolecuul neutraal is. 

 Als de polarisatie van gedeelde elektronen binnen de O-H-bindingen van water compleet zou zijn, zou het resultaat de volledige overdracht van elektronen van H naar O zijn, en zou water een ionische verbinding zijn die bestaat uit O2-anionen en H+-kationen. En dus zijn de oxidatiegetallen voor zuurstof en waterstof in water respectievelijk −2 en +1. Het toepassen van dezelfde logica op tetrachloorkoolstof, CCl4, levert oxidatiegetallen op van +4 voor koolstof en −1 voor chloor. In het nitraation, NO3−, is het oxidatiegetal voor stikstof +5 en dat voor zuurstof −2, opgeteld gelijk aan de −1 lading op het molecuul: 

 <img alt="Eq1" src="/files/ftp_upload/11432/11432_Equation_1.png" style="text-align: center; height: 70px;"> 

 Redox-vergelijkingen in evenwicht brengen 

 De onevenwichtige vergelijking hieronder beschrijft de ontleding van natriumchloride: 

 <img alt="Eq2" src="/files/ftp_upload/11432/11432_Equation_2.png" style="text-align: center; height: 48px;"> 

 Deze reactie komt overeen met het criterium voor een redoxreactie, aangezien het oxidatiegetal voor Na afneemt van +1 naar 0 (door reductie te ondergaan) en dat voor Cl toeneemt van −1 naar 0 (door oxidatie te ondergaan). Het vergelijkingsgeval kan gemakkelijk in evenwicht worden gebracht door de stoichiometrische coëfficiënt van 2 voor NaCl en Na toe te voegen: 

 <img alt="Eq3" src="/files/ftp_upload/11432/11432_Equation_3.png" style="text-align: center; height: 48px;"> 

 Redoxreacties die voorkomen in waterige oplossingen komen vaak voor in de elektrochemie, en bij veel daarvan is water of de ionen ervan betrokken, H+ (aq) en OH− (aq), hetzij als reactanten of producten. 

 In deze gevallen kan het een hele uitdaging zijn om vergelijkingen die redoxreacties vertegenwoordigen alleen door inspectie in evenwicht te brengen, en het gebruik van een systematische aanpak die bekend staat als de halfreactiemethode is nuttig. Deze aanpak omvat de volgende stappen: 
 <ol> 
 	<li>Splits de vergelijking op in de componenten skeletoxidatie en reductiehalfreacties.</li> 
 	<li>Breng elke halfreactie in evenwicht voor alle elementen behalve O en H.</li> 
 	<li>Breng elke halfreactie voor O-atomen in evenwicht door indien nodig watermoleculen toe te voegen in de vergelijking.</li> 
 	<li>Breng elke halfreactie voor H-atomen in evenwicht door protonen toe te voegen zoals vereist voor de vergelijking.</li> 
 	<li>Breng ten slotte de ladingen op de elementen in evenwicht door indien nodig elektronen toe te voegen.</li> 
 	<li>Vermenigvuldig de halfreacties met een willekeurig geheel getal dat nodig is om het aantal elektronen dat verloren gaat in de oxidatie-halfreactie gelijk te maken aan het aantal gewonnen elektronen in de reductie-halfreactie.</li> 
 	<li>Voeg beide halfreacties toe en vereenvoudig het verder door algemene soorten aan beide kanten van de vergelijking weg te laten.</li> 
 	<li>Als de reactie plaatsvindt in een alkalisch medium, voeg dan OH− ionen toe aan de vergelijking verkregen in stap 7 om de protonen te neutraliseren (toegevoegd in gelijke aantallen aan beide kanten van de vergelijking) en vereenvoudig.</li> 
 	<li>Controleer de vergelijking om er zeker van te zijn dat de ladingen op de atomen in evenwicht zijn.</li> 
 </ol> 

 Dit tekst is aangepast van  <a href="https://openstax.org/books/chemistry-2e/pages/17-1-review-of-redox-chemistry">Openstax, Chemistry 2e, Chapter 17: Introduction,</a> en <a href="https://openstax.org/books/chemistry-2e/pages/17-1-review-of-redox-chemistry">Openstax, Chemistry 2e, Section17.2: Review of Redox Chemistry</a>.

Redox-vergelijkingen in evenwicht brengen

Electrochemistry is the science involved in the interconversion of electrical and chemical reactions. Such reactions are called reduction-oxidation, or ...

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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

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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...

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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...

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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...

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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...

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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...

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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...

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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 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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