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

K12 Standards

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North Carolina K-12 Science Standard Course of Study

9-12

Physical Science

Matter and its Interactions

PS.PSc.2.3

Balancing Redox Equations

In che modo una batteria fornisce elettricità per far funzionare i dispositivi portatili? Perché un avocado tagliato diventa marrone, o il metallo arrugginisce sott'acqua? Questi processi sono guidati da specifici tipi di reazioni chimiche, che coinvolgono il trasferimento di elettroni da un atomo all'altro.

L'atomo che perde elettroni viene ossidato, mentre l'atomo che guadagna elettroni viene ridotto. Queste reazioni sono chiamate reazioni di riduzione-ossidazione o redox, e sono caratterizzate da cambiamenti nello stato di ossidazione per uno o più reagenti. L'interazione fra magnesio e acido cloridrico concentrato è un esempio di reazione redox.

Qui, il magnesio viene ossidato a ioni magnesio 2, e i protoni vengono ridotti a idrogeno gassoso. Semplici equazioni chimiche possono essere facilmente bilanciate. Tuttavia, poiché la maggior parte delle equazioni redox sono complesse, il numero di elettroni persi e guadagnati deve essere giustificato.

Quando si bilanciano le equazioni redox, è essenziale osservare la legge di conservazione di massa. La quantità di ogni elemento e qualsiasi guadagno o perdita di elettroni devono essere bilanciati su entrambi i lati della reazione;quindi, se un reagente viene ridotto, un altro deve essere ossidato. Il metodo delle semi-reazioni viene usato per bilanciare con successo le equazioni redox.

Considerate la reazione del permanganato con l'ossalato. Per soluzioni acquose acide, assegnate gli stati di ossidazione e dividete la reazione nelle sue semireazioni componenti. Quindi, bilanciate ogni semireazione omettendo idrogeno e ossigeno.

Qui il manganese è già bilanciato, ma l'ossalato necessita di un coefficiente di due;successivamente, bilanciate gli atomi di ossigeno con l'aggiunta di molecole d'acqua;e infine, bilanciate gli atomi di idrogeno, aggiungendo protoni dove richiesto. Bilanciate le cariche aggiungendo elettroni. Poiché il manganese è ridotto, vengono aggiunti cinque elettroni sul lato reagente.

Al contrario, l'ossalato è ossidato;quindi, due elettroni vengono aggiunti sul lato del prodotto. Moltiplicate le semireazioni per un numero intero, cosicché il numero di elettroni aggiunti sia uguale. Infine, aggiungete e semplificate le semireazioni bilanciate, cancellando le specie su ambo i lati per ottenere una reazione redox bilanciata.

Per le soluzioni acquose di base, la procedura è simile ma include un passaggio aggiuntivo. Qui, un numero uguale di ioni idrossile è aggiunto ad ambo i lati della reazione per neutralizzare i protoni prima di sommare le semireazioni bilanciate.

bilanciamento delle equazioni redox

L'elettrochimica è la scienza coinvolta nell'interconversione delle reazioni elettriche e chimiche. Tali reazioni sono chiamate reazioni di riduzione-ossidazione o reazioni redox. Queste importanti reazioni sono definite da cambiamenti negli stati di ossidazione per uno o più elementi reagenti e includono un sottoinsieme di reazioni che comportano il trasferimento di elettroni tra specie reagenti. L'elettrochimica come campo si è evoluta per fornire approfondimenti sufficienti sui principi fondamentali della chimica redox e su molteplici tecnologie che vanno dai processi metallurgici su scala industriale alle robuste batterie ricaricabili per veicoli elettrici. Poiché le reazioni che comportano il trasferimento di elettroni sono essenziali per lo studio dell'elettrochimica, una breve rassegna della chimica redox include quanto segue.

Numeri di ossidazione

Per definizione, una reazione redox comporta un cambiamento nel numero di ossidazione o nello stato di ossidazione per uno o più reagenti. Il numero di ossidazione di un elemento è una valutazione di come l'ambiente elettronico dei suoi atomi differisce rispetto agli atomi dell'elemento puro. Secondo questa definizione, un atomo in un elemento porta un numero di ossidazione pari a zero. Per un atomo, il numero di ossidazione è uguale alla carica dell'atomo nel composto se il composto è ionico. Pertanto, la somma dei numeri di ossidazione di tutti gli atomi di una molecola è uguale alla carica della molecola.

Composti ionici

I composti ionici semplici sono gli esempi più semplici di questo formalismo poiché gli elementi hanno numeri di ossidazione pari alle loro cariche ioniche. Il cloruro di sodio, NaCl, è composto da cationi Na^+ e anioni Cl^−, con numeri di ossidazione per sodio e cloro rispettivamente +1 e −1. Il fluoruro di calcio, CaF_2, è composto da cationi Ca^2+ e anioni F^−, con numeri di ossidazione per calcio e fluoro, +2 e −1.

Composti covalenti

I composti covalenti sono più impegnativi nell'uso del formalismo. L'acqua è un composto covalente costituito da due atomi di H legati a un atomo di O tramite legami covalenti polari OH. Gli elettroni condivisi che formano un legame OH sono attratti più fortemente dall'atomo O più elettronegativo. Quindi l'ossigeno acquisisce una carica negativa parziale, rispetto a un atomo di O nell'ossigeno elementare. Di conseguenza, gli atomi di H in una molecola d’acqua mostrano una carica positiva parziale rispetto agli atomi di idrogeno nell’idrogeno gassoso. La somma delle cariche parziali negative e positive per ciascuna molecola d'acqua è zero, rendendo la molecola d'acqua neutra.

Se la polarizzazione degli elettroni condivisi all’interno dei legami O−H dell’acqua fosse completa, il risultato sarebbe il completo trasferimento di elettroni da H a O, e l’acqua sarebbe un composto ionico composto da anioni O_2^− e cationi H^+. Pertanto, i numeri di ossidazione dell'ossigeno e dell'idrogeno nell'acqua sono rispettivamente −2 e +1. Applicando questa stessa logica al tetracloruro di carbonio, CCl_4, si ottengono numeri di ossidazione di +4 per il carbonio e -1 per il cloro. Nello ione nitrato, NO_3^−, il numero di ossidazione dell'azoto è +5 e quello dell'ossigeno è −2, la cui somma equivale alla carica −1 della molecola:

<img alt="Eq1" src="/files/ftp_upload/11432/11432_Equation_1.png" style="text-align: center; height: 70px;" />

Bilanciamento delle equazioni Redox

L'equazione sbilanciata mostrata di seguito descrive la decomposizione del cloruro di sodio:

<img alt="Eq2" src="/files/ftp_upload/11432/11432_Equation_2.png" style="text-align: center; height: 48px;" />

Questa reazione corrisponde al criterio per una reazione redox, poiché il numero di ossidazione del Na diminuisce da +1 a 0 (subendo riduzione) e quello del Cl aumenta da −1 a 0 (subendo ossidazione). Il caso dell'equazione è facilmente bilanciato aggiungendo il coefficiente stechiometrico di 2 per NaCl e Na:

<img alt="Eq3" src="/files/ftp_upload/11432/11432_Equation_3.png" style="text-align: center; height: 48px;" />

Le reazioni redox che si verificano in soluzioni acquose si incontrano comunemente in elettrochimica e molte coinvolgono l'acqua o i suoi ioni, H^+ (aq) e OH^− (aq), sia come reagenti che come prodotti.

In questi casi, le equazioni che rappresentano le reazioni redox possono essere molto difficili da bilanciare semplicemente mediante ispezione, ed è utile l’uso di un approccio sistematico noto come metodo della semireazione. Questo approccio prevede i seguenti passaggi:

<ol>
	<li>Suddividere l'equazione nelle semireazioni di ossidazione dello scheletro e di riduzione dei suoi componenti.</li>
	<li>Bilancia ciascuna semireazione per tutti gli elementi diversi da O e H.</li>
	<li>Bilancia ciascuna semireazione per gli atomi di O aggiungendo molecole d'acqua secondo necessità nell'equazione.</li>
	<li>Bilancia ciascuna semireazione per gli atomi di H aggiungendo protoni come richiesto per l'equazione.</li>
	<li>Infine, bilancia le cariche degli elementi aggiungendo elettroni secondo necessità.</li>
	<li>Moltiplicare le semireazioni per qualsiasi numero intero necessario per uguagliare il numero di elettroni persi nella semireazione di ossidazione con il numero di elettroni acquistati nella semireazione di riduzione.</li>
	<li>Aggiungi entrambe le semireazioni e semplifica ulteriormente il tutto cancellando le specie comuni su entrambi i lati dell'equazione.</li>
	<li>Se la reazione avviene in un mezzo alcalino, aggiungere ioni OH^− all'equazione ottenuta nel passaggio 7 per neutralizzare i protoni (aggiunti in numero uguale su entrambi i lati dell'equazione) e semplificare.</li>
	<li>Controlla l'equazione per assicurarti che le cariche sugli atomi siano bilanciate.</li>
</ol>

Questo testo è adattato da  <a href="https://openstax.org/books/chemistry-2e/pages/17-1-review-of-redox-chemistry">Openstax, Chemistry 2e, Chapter 17: Introduction,</a> and <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>.

Bilanciamento delle equazioni Redox

L'elettrochimica è la scienza coinvolta nell'interconversione delle reazioni elettriche e chimiche. Tali reazioni sono chiamate riduzione-ossidazione, o ...

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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...

cell potential and free energy

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,...

corrosion

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...

key terms and concepts

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

Redox Reactions and Half-Reaction Balancing | JoVE K12

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JoVE Core: Chimica illustra concetti chiave di chimica attraverso video animazioni concise e di immediata comprensione. Inoltre, i video della sezione "scienziati in azione" mostrano procedure metodologiche realizzate da ricercatori provenienti dai migliori laboratori internazionali.

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Chemistry and Physics

IPC.8

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Matter and Its Interactions

CHEM1.PS1.2

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PS.3

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Balancing Chemical Equations and Reaction Outcomes

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Physical Properties of Matter

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Molecular Comparison of Gases, Liquids, and Solids

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Comparing Intermolecular Forces: Melting Point, Boiling Point, and Miscibility

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Carry out investigations to compare physical and chemical properties of matter.

PS.PSc.1.3

The Atomic Theory of Matter

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Subatomic Particles

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The Bohr Model

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Emission Spectra

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The de Broglie Wavelength

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The Uncertainty Principle

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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.

PS.PSc.1.4

Ions and Ionic Charges

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Use models to compare representations of atoms, ions, and isotopes.

PS.PSc.1.5

The Periodic Table

As early chemists discovered more elements, they realized that various elements could be grouped by their similar chemical behaviors. One such grouping ...

Periodic Classification of the Elements

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Ionic Radii

Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom; it ...

Ionization Energy

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Electron Affinity

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Alkali Metals

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Halogens

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Noble Gases


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Use the Periodic Table as a model to predict the relative properties (metallic/nonmetallic character, ionic charge, and reactivity) and arrangement of elements based on the pattern of valence electrons in the outermost energy levels of atoms.

PS.PSc.1.6

Types of Chemical Bonds

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Ionic Bonding and Electron Transfer

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Covalent Bonding and Lewis Structures

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Bonding in Metals

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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

Nomenclature of Alkanes

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Structure and Nomenclature of Ethers

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IUPAC Nomenclature of Ketones

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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.

PS.PSc.2.2

Understand types, properties, and structure of matter.

Redox Reactions and Half-Reaction Balancing

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