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

Las reacciones electroquímicas son vitales para muchos procesos en la tecnología y el cuerpo humano e implican la transferencia de electrones de una especie química a otra. Estas reacciones se denominan reacciones de reducción-oxidación o reacciones redox.
Cuando una especie pierde electrones, sufre oxidación, y decimos que se oxida. Cuando una especie gana electrones, ha sufrido una reducción, y nosotros decimos que se reduce. En una celda galvánica electroquímica, las reacciones redox ocurren espontáneamente, lo que genera energía eléctrica. Sin embargo, en una celda electrolítica, se aplica energía eléctrica, lo que hace que ocurran las reacciones redox.
Una celda electrolítica consiste en la solución de reacción llamada electrolito, que contiene cualquier especie química que se oxida o reduce, así como los iones necesarios para permitir el flujo de electrones. La celda electrolítica tiene dos electrodos metálicos que se sumergen en la solución electrolítica. Un circuito externo conecta los dos electrodos, lo que completa el circuito y nos permite aplicar voltaje o corriente.
La corriente o voltaje aplicado es lo que hace que se produzca la reacción electroquímica. La semirreacción de oxidación se produce en el ánodo y la semirreacción de reducción se produce en el cátodo. Para ayudarlo a diferenciar entre los dos, recuerde la frase 'gato rojo', que significa 'reducción en el cátodo'.
Pero, ¿cómo determinamos qué electrodo es cuál? Podemos usar el potencial de reducción estándar de los electrodos metálicos, que es una medida de la tendencia del metal a perder electrones. Cuanto mayor sea el potencial de reducción, más probable es que se reduzca el metal.
Digamos que colocamos un electrodo de plata y un electrodo de cobre en una celda electrolítica y aplicamos corriente. El electrodo de plata tiene un mayor potencial de reducción, por lo que se reduce y actúa como cátodo. El electrodo de cobre tiene un menor potencial de reducción, por lo que se oxida y actúa como ánodo.
Uno de los usos de las celdas electrolíticas es realizar la galvanoplastia, que es una reacción en la que un metal se oxida y luego se reduce a la superficie de otro metal. Dado que el ánodo es el metal que hace el recubrimiento, en nuestro ejemplo de plata y cobre, el electrodo de plata está recubierto con una fina capa de cobre.
En este laboratorio, ensamblará una celda electrolítica y realizará el proceso de galvanoplastia recubriendo una llave de latón con cobre.

celdas electrolíticas

electroquímica
La electroquímica es una rama de la química que describe y mide la relación entre la energía eléctrica y un cambio químico. Las reacciones electroquímicas implican el movimiento de electrones de una especie a otra. Si la reacción es espontánea, puede dar lugar a una corriente generada. Si la reacción no es espontánea, puede ser impulsada por la aplicación de corriente.
En electroquímica, la reacción clave es la reacción de oxidación-reducción, llamada reacción redox. La reacción redox se compone de dos medias reacciones; la oxidación, en la que una sustancia conocida como agente reductor pierde electrones, y la reducción, en la que una sustancia conocida como agente oxidante gana electrones. Las reacciones redox siempre ocurren en pares y cambian los estados de oxidación de los átomos en las moléculas involucradas. Una manera fácil de recordar esta relación es a través de la OIL-RIG neumática: Oxidación Is Losing – Reduction Is Gaining.
<h4>Celdas electroquímicas</h4>
Una celda electroquímica se utiliza para medir o inducir reacciones electroquímicas. Consta de varios componentes: una cámara que contiene la solución de reacción, dos electrodos conductores, una solución de electrolito conductor y un circuito externo. Hay dos tipos de celdas electroquímicas. Una es la celda electrolítica, que utiliza energía eléctrica para impulsar una reacción no espontánea. En este tipo de celda, la energía eléctrica se suministra desde una fuente de alimentación externa.
El otro tipo de célula es la célula galvánica, que utiliza una reacción electroquímica espontánea para generar energía eléctrica. Los dos electrodos se llaman ánodo y cátodo, y están conectados por un circuito externo. La cámara de reacción está llena de un electrolito, que promueve el paso de iones entre electrodos. En una celda galvánica, un puente salino completa el circuito permitiendo el flujo de iones entre las soluciones electrolíticas que contienen los electrodos. En una celda electrolítica, no hay puente salino, ya que los electrodos suelen estar en la misma solución electrolítica.
La reacción de reducción ocurre en el cátodo, mientras que la reacción de oxidación ocurre en el ánodo. Esto se recuerda usando el mnemónico "Gato Rojo", lo que significa que la ucción rojase produce en la manguera del gato.
<h4>Celdas electrolíticas</h4>
Las celdas electrolíticas suelen tener una cámara de reacción, que contiene la solución electrolítica. El electrolito suele ser una solución acuosa que contiene iones o sales disueltas. Los iones en el electrolito promueven el movimiento de iones y electrones a través de la solución.
Cuando se aplica un voltaje externo, los iones en el electrolito son atraídos al electrodo con la carga opuesta. Aquí es donde se producen las dos medias reacciones. El ánodo pierde electrones durante la oxidación, mientras que el cátodo gana electrones durante la reducción.
La galvanoplastia es un proceso que utiliza una celda electrolítica. La galvanoplastia es un proceso en el que un metal se deposita en la superficie de un electrodo, normalmente otro metal. El científico británico Michael Faraday, por quien la constante de Faraday toma su nombre, demostró la relación molar entre el ion cargado de galvanoplastia y la corriente eléctrica. En cuanto a la semirreacción específica de la plata:
Ag+ + 1e- → Ag
Se necesita un mol de electrones, suministrado por una corriente externa, para reducir un mol de cationes de plata a plata sólida. Por lo tanto, la estequiometría de media reacción puede determinar la cantidad de material de galvanoplastia utilizado para una cantidad conocida de electrones. Recordemos que la electricidad no se mide en función de los moles de electrones, sino en culombios, llamados así por el ingeniero francés Charles-Augustin de Coulomb. Un culombio es la cantidad de electricidad que pasa a través de un circuito eléctrico con 1 amperio de corriente en un intervalo de 1 segundo, que se basa en el principio de carga elemental, o la carga que lleva un solo electrón. Por lo tanto, el descubrimiento de Faraday puede relacionar la corriente eléctrica con el número de moles que placa el catión de galvanoplastia:
F = eNA
F es la constante de Faraday, e es la carga elemental expresada en culombios y NA es el número de Avogadro. Por lo tanto, la constante de Faraday se expresa como el número de culombios por mol y tiene un valor de 96485 culombios por mol. ¿Cómo podemos determinar el número de moles de electrones transferidos al ánodo en función de la corriente? Usando la carga eléctrica Q y la definición de un culombio:
<img src="/CDNSource/lm/labs/53/53_Concepts_3.jpg">
La carga eléctrica es igual a la corriente en amperios (I) multiplicada por el tiempo en segundos que se permitió que fluyera la corriente. Dividiendo Q por la constante de Faraday, que tiene las unidades de culombios por mol:
<img src="/CDNSource/lm/labs/53/53_Concepts_4.jpg">
Esta ecuación nos permite calcular el número de moles de electrones y así determinar cuánto del catión galvanoplastia se redujo.
En la galvanoplastia, el metal comprende las placas del ánodo, o cubre, el cátodo en una capa delgada de metal. La cantidad de metal chapado depende de la cantidad de corriente aplicada, así como del número de moles del catión galvanoplastia disponibles. Una fuente de alimentación externa, como una batería, induce el flujo de electrones del ánodo al cátodo y del terminal positivo al terminal negativo de la batería.
Por ejemplo, considere una celda electrolítica con un electrodo de cobre, una llave de latón que actúa como segundo electrodo y una solución acuosa de sulfato de cobre como electrolito. Aquí, el cobre del electrolito y el electrodo de cobre se recubre en la llave de latón.
Para que el metal de cobre se deposite en la llave de latón, el electrodo de cobre sólido debe oxidarse para formar iones de cobre. Luego, los cationes de cobre tanto del electrodo como del electrolito se reducen de la solución para formar cobre sólido en la llave de latón.
Cu2+ + 2e- → Cu(s)
Los electrones para la reacción se reciben del terminal negativo de la batería. Por lo tanto, la reacción de reducción ocurre en la llave de latón, mientras que la reacción de oxidación ocurre en el electrodo de cobre. La solución de sulfato de cobre concentrada y acidificada aumenta la solubilidad; Por lo tanto, cuanto mayor sea la concentración de la solución, menor será la resistencia y mayor será la corriente. A su vez, una corriente más alta permite que se produzca una mayor chapa de los iones de cobre en la llave de latón.
Algunos metales tienen una mayor tendencia que otros a perder electrones. El potencial de electrodo estándar (E°) de una sustancia es la medida de la tendencia de una sustancia a perder electrones. El metal con el mayor potencial de reducción tiene la mayor tendencia a perder electrones; por lo tanto, primero se galvaniza.
<h4>Referencias</h4>
<ol>
	<li style="font-size:17px">Kotz, J.C., Treichel Jr, P.M., Townsend, J.R. (2012). Química y reactividad química. Belmont, CA: Brooks/Cole, Cengage Learning.</li>
	<li style="font-size:17px">Silberberg, M.S. (2009). Química: La naturaleza molecular de la materia y el cambio. Boston, MA: McGraw Hill.</li>
</ol>

Celdas electrolíticas

Electroquímica
La electroquímica es una rama de la química que describe y mide la relación entre la energía eléctrica y un cambio químico. ...

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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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Electroplating with Electrolytic Cells

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