Name: electroplating with electrolytic cells
Uploaded: 2020-03-26T07:22:26.000Z
Duration: 2 min 52 s
Channel: JoVE

Le reazioni elettrochimiche sono vitali per molti processi nella tecnologia e nel corpo umano e comportano il trasferimento di elettroni da una specie chimica all'altra. Queste reazioni sono chiamate reazioni di riduzione-ossidazione o reazioni redox.
Quando una specie perde elettroni, subisce ossidazione, e diciamo che è ossidata. Quando una specie guadagna elettroni, ha subito una riduzione, e noi diciamo che è ridotta. In una cella galvanica elettrochimica avvengono spontaneamente reazioni di ossidoriduzione, che generano energia elettrica. Tuttavia, in una cella elettrolitica, viene applicata energia elettrica, che provoca le reazioni redox.
Una cella elettrolitica è costituita dalla soluzione di reazione chiamata elettrolita, che contiene tutte le specie chimiche ossidate o ridotte, nonché gli ioni necessari per consentire il flusso di elettroni. La cella elettrolitica ha due elettrodi metallici che sono immersi nella soluzione elettrolitica. Un circuito esterno collega i due elettrodi, che completa il circuito e ci consente di applicare tensione o corrente.
La corrente o la tensione applicata è ciò che provoca la reazione elettrochimica. La semireazione di ossidazione avviene all'anodo e la semireazione di riduzione avviene al catodo. Per aiutarti a distinguere tra i due, ricorda la frase "gatto rosso", che sta per "riduzione al catodo".
Ma come si fa a determinare quale elettrodo è quale? Possiamo utilizzare il potenziale di riduzione standard degli elettrodi metallici, che è una misura della tendenza del metallo a perdere elettroni. Maggiore è il potenziale di riduzione, maggiore è la probabilità che il metallo venga ridotto.
Diciamo che posizioniamo un elettrodo d'argento e un elettrodo di rame in una cella elettrolitica e applichiamo corrente. L'elettrodo d'argento ha un potenziale di riduzione più elevato, quindi è ridotto e funge da catodo. L'elettrodo di rame ha un potenziale di riduzione inferiore, quindi viene ossidato e funge da anodo.
Un uso delle celle elettrolitiche è quello di eseguire la galvanica, che è una reazione in cui un metallo viene ossidato e poi ridotto sulla superficie di un altro metallo. Poiché l'anodo è il metallo che esegue la placcatura, nel nostro esempio di argento e rame, l'elettrodo d'argento è placcato con un sottile strato di rame.
In questo laboratorio, assemblerai una cella elettrolitica ed eseguirai il processo di galvanica placcando una chiave di ottone con rame.

celle elettrolitiche

elettrochimica
L'elettrochimica è una branca della chimica che descrive e misura la relazione tra l'energia elettrica e un cambiamento chimico. Le reazioni elettrochimiche coinvolgono il movimento di elettroni da una specie all'altra. Se la reazione è spontanea, può provocare la generazione di una corrente. Se la reazione non è spontanea, può essere guidata dall'applicazione di corrente.
In elettrochimica, la reazione chiave è la reazione di ossidoriduzione, chiamata reazione redox. La reazione redox è composta da due semireazioni; ossidazione, in cui una sostanza nota come agente riducente perde elettroni, e riduzione, in cui una sostanza nota come agente ossidante guadagna elettroni. Le reazioni redox avvengono sempre in coppia e modificano gli stati di ossidazione degli atomi nelle molecole coinvolte. Un modo semplice per ricordare questa relazione è attraverso l'OIL-RIG pneumatico: Oxidation Is Losing – Reduction Is Gaining.
<h4>Celle elettrochimiche</h4>
Una cella elettrochimica viene utilizzata per misurare o indurre reazioni elettrochimiche. È costituito da diversi componenti: una camera contenente la soluzione di reazione, due elettrodi conduttivi, una soluzione elettrolitica conduttiva e un circuito esterno. Esistono due tipi di celle elettrochimiche. Una è la cella elettrolitica, che utilizza l'energia elettrica per guidare una reazione non spontanea. In questo tipo di cella, l'energia elettrica viene fornita da una fonte di alimentazione esterna.
L'altro tipo di cella è la cella galvanica, che utilizza una reazione elettrochimica spontanea per generare energia elettrica. I due elettrodi sono chiamati anodo e catodo e sono collegati da un circuito esterno. La camera di reazione è riempita con un elettrolita, che favorisce il passaggio degli ioni tra gli elettrodi. In una cella galvanica, un ponte salino completa il circuito consentendo il flusso di ioni tra le soluzioni elettrolitiche contenenti gli elettrodi. In una cella elettrolitica, non c'è ponte salino poiché gli elettrodi si trovano tipicamente nella stessa soluzione elettrolitica.
La reazione di riduzione avviene a livello del catodo, mentre la reazione di ossidazione avviene a livello dell'anodo. Questo è ricordato usando il mnemonico "Gatto Rosso", che significa che l'uction rossasi verifica allo zoccolo del gatto.
<h4>Celle elettrolitiche</h4>
Le celle elettrolitiche hanno tipicamente una camera di reazione, che contiene la soluzione elettrolitica. L'elettrolita è solitamente una soluzione acquosa contenente ioni o sali disciolti. Gli ioni nell'elettrolita promuovono il movimento di ioni ed elettroni attraverso la soluzione.
Quando viene applicata una tensione esterna, gli ioni nell'elettrolita vengono attratti dall'elettrodo con la carica opposta. È qui che si verificano le due semireazioni. L'anodo perde elettroni durante l'ossidazione, mentre il catodo guadagna elettroni durante la riduzione.
La galvanica è un processo che utilizza una cella elettrolitica. La galvanica è un processo in cui un metallo viene depositato sulla superficie di un elettrodo, in genere un altro metallo. Lo scienziato britannico Michael Faraday, da cui prende il nome la costante di Faraday, ha dimostrato la relazione molare tra lo ione carico galvanico e la corrente elettrica. Osservando la semireazione specifica per l'argento:
Ag+ + 1e- → Ag
Ci vuole una mole di elettroni, alimentata da una corrente esterna, per ridurre una mole di cationi d'argento in argento massiccio. Pertanto, la stechiometria a semireazione può determinare la quantità di materiale galvanico utilizzato per una quantità nota di elettroni. Ricordiamo che l'elettricità non si misura in base alle moli degli elettroni ma piuttosto in coulomb, dal nome dell'ingegnere francese Charles-Augustin de Coulomb. Un coulomb è la quantità di elettricità che passa attraverso un circuito elettrico con 1 ampere di corrente in un intervallo di 1 secondo, che si basa sul principio della carica elementare, o la carica trasportata da un singolo elettrone. La scoperta di Faraday può quindi mettere in relazione la corrente elettrica con il numero di moli placcate dal catione galvanico:
F = eNA
F è la costante di Faraday, e è la carica elementare espressa in coulomb e NA è il numero di Avogadro. Pertanto, la costante di Faraday è espressa come il numero di coulomb per mole e ha un valore di 96485 coulomb per mole. Come possiamo determinare il numero di moli di elettroni trasferiti all'anodo in base alla corrente? Utilizzando la carica elettrica Q e la definizione di un coulomb:
<img src="/CDNSource/lm/labs/53/53_Concepts_3.jpg">
La carica elettrica è uguale alla corrente in ampere (I) moltiplicata per il tempo in secondi in cui la corrente è stata lasciata fluire. Dividendo Q per la costante di Faraday, che ha le unità di coulomb per mole:
<img src="/CDNSource/lm/labs/53/53_Concepts_4.jpg">
Questa equazione ci permette di calcolare il numero di moli di elettroni e quindi di determinare quanto del catione galvanico è stato ridotto.
Nella galvanica, il metallo comprende le piastre anodiche, o copre, il catodo in un sottile strato di metallo. La quantità di metallo placcato dipende dalla quantità di corrente applicata e dal numero di moli del catione galvanico disponibile. Una fonte di alimentazione esterna, come una batteria, induce il flusso di elettroni dall'anodo al catodo e dal terminale positivo al terminale negativo della batteria.
Ad esempio, si consideri una cella elettrolitica con un elettrodo di rame, una chiave di ottone che funge da secondo elettrodo e una soluzione acquosa di solfato di rame come elettrolita. Qui, il rame dell'elettrolita e l'elettrodo di rame vengono placcati sulla chiave di ottone.
Affinché il rame metallico si depositi sulla chiavetta in ottone, l'elettrodo di rame solido deve essere ossidato per formare ioni di rame. Quindi, i cationi di rame provenienti sia dall'elettrodo che dall'elettrolita vengono ridotti dalla soluzione per formare rame solido sulla chiave di ottone.
Cu2+ + 2e- → Cu(s)
Gli elettroni per la reazione vengono ricevuti dal terminale negativo della batteria. Pertanto, la reazione di riduzione avviene alla chiave di ottone, mentre la reazione di ossidazione avviene all'elettrodo di rame. La soluzione concentrata e acidificata di solfato di rame aumenta la solubilità; Pertanto, maggiore è la concentrazione della soluzione, minore è la resistenza e maggiore è la corrente. A sua volta, una corrente più elevata consente una maggiore placcatura degli ioni di rame sulla chiave di ottone.
Alcuni metalli hanno una tendenza maggiore di altri a perdere elettroni. Il potenziale elettrodico standard (E°) di una sostanza è la misura della tendenza di una sostanza a perdere elettroni. Il metallo con il più alto potenziale di riduzione ha la più alta tendenza a perdere elettroni; quindi, viene prima galvanizzato.
<h4>referenze</h4>
<ol>
	<li style="font-size:17px">Kotz, J.C., Treichel Jr, P.M., Townsend, J.R. (2012). Chimica e reattività chimica. Belmont, CA: Brooks/Cole, Cengage Learning.</li>
	<li style="font-size:17px">Silberberg, M.S. (2009). Chimica: la natura molecolare della materia e il cambiamento. Boston, MA: McGraw Hill.</li>
</ol>

Celle elettrolitiche

Elettrochimica
L'elettrochimica è una branca della chimica che descrive e misura la relazione tra l'energia elettrica e un cambiamento chimico. ...

Education

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Chemistry

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Electrolytic Cells

electroplating with electrolytic cells

Electrochemistry links electrical energy with chemical change, and electrolytic cells show how electricity can drive a nonspontaneous reaction. In these reactions, electrons move from one species to another. The key chemistry is a redox reaction, or oxidation-reduction reaction, which happens in two half-reactions.

Oxidation is the loss of electrons, and reduction is the gain of electrons. The substance that loses electrons is the reducing agent, and the substance that gains electrons is the oxidizing agent. Redox reactions always occur in pairs and change the oxidation state of the atoms involved. A simple memory aid is OIL-RIG: oxidation is losing, reduction is gaining.

An electrochemical cell is used to measure or cause electrochemical reactions. It includes a reaction chamber, two conductive electrodes, a conductive electrolyte, and an external circuit. The two main types are electrolytic cells and galvanic cells. An electrolytic cell uses outside electrical energy to drive a nonspontaneous reaction, while a galvanic cell uses a spontaneous reaction to produce electrical energy. The electrodes are the anode and cathode, and the reaction at the cathode is reduction while the reaction at the anode is oxidation. This is often remembered with “Red Cat,” which means reduction occurs at the cathode.

Electrolytic cells usually have one reaction chamber filled with an electrolyte solution. The electrolyte is often an aqueous solution with ions or dissolved salts. These ions help charge move through the solution. When voltage is applied, ions move toward the electrode with the opposite charge, and the two half-reactions occur there. The anode loses electrons during oxidation, and the cathode gains electrons during reduction.

A major use of an electrolytic cell is electroplating. In electroplating, one metal is deposited as a thin layer on the surface of another electrode. Michael Faraday showed the link between the charge of the ions being plated and the electrical current. For silver, the half-reaction is Ag+ + 1e- -&gt; Ag, so one mole of electrons reduces one mole of silver cations to solid silver. This means half-reaction stoichiometry can be used to calculate how much plating occurs for a known amount of charge.

Electricity is measured in coulombs, not in moles of electrons. One coulomb is the amount of electricity that passes through a circuit with 1 ampere of current in 1 second. Faraday’s constant connects these ideas. It is equal to the elementary charge times Avogadro’s number, and it has a value of 96485 coulombs per mole. The electrical charge is found by multiplying current, I, by time in seconds, and dividing Q by Faraday’s constant gives the moles of electrons transferred.

This relationship shows how current controls electroplating. The amount of metal plated depends on the current applied and on how much electroplating cation is available. An external power source, such as a battery, pushes electrons from the anode to the cathode and from the positive terminal to the negative terminal on the battery. In a copper plating example, a copper electrode, a brass key, and aqueous copper sulfate are used. Copper from the electrode and from the solution is deposited onto the brass key.

For copper to coat the key, the solid copper electrode must first be oxidized to copper ions. Then copper cations in the solution are reduced to solid copper on the brass key, following Cu2+ + 2e- -&gt; Cu(s). The electrons come from the negative terminal of the battery. So the brass key is the site of reduction, and the copper electrode is the site of oxidation. A concentrated and acidified copper sulfate solution increases solubility, lowers resistance, and raises the current, which allows more copper to plate onto the key.

Some metals lose electrons more easily than others. The standard electrode potential, E°, measures a substance’s tendency to lose electrons. In electroplating, the metal with the highest reduction potential has the highest tendency to lose electrons, so it is plated first.

Electroplating with Electrolytic Cells

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Video completi e incentrati sul curriculum sia per gli istruttori che per gli studenti nei corsi di laboratorio di Chimica introduttiva.

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Learn how electrolytic cells drive electroplating, move electrons through redox reactions, and use Faraday’s constant to calculate plated metal.

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Electroplating with Electrolytic Cells

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