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

As reações eletroquímicas são vitais para muitos processos na tecnologia e no corpo humano e envolvem a transferência de elétrons de uma espécie química para outra. Essas reações são chamadas de reações de redução-oxidação ou reações redox.
Quando uma espécie perde elétrons, ela sofre oxidação, e dizemos que está oxidada. Quando uma espécie ganha elétrons, ela sofreu redução, e dizemos que está reduzida. Em uma célula galvânica eletroquímica, ocorrem espontaneamente reações redox, que geram energia elétrica. No entanto, em uma célula eletrolítica, a energia elétrica é aplicada, o que faz com que as reações redox ocorram.
Uma célula eletrolítica consiste na solução de reação chamada eletrólito, que contém qualquer espécie química que seja oxidada ou reduzida, bem como íons necessários para permitir o fluxo de elétrons. A célula eletrolítica possui dois eletrodos metálicos que são imersos na solução eletrolítica. Um circuito externo conecta os dois eletrodos, que completa o circuito e nos permite aplicar tensão ou corrente.
A corrente ou tensão aplicada é o que faz com que a reação eletroquímica ocorra. A semi-reação de oxidação ocorre no ânodo e a semi-reação de redução ocorre no cátodo. Para ajudá-lo a diferenciar entre os dois, lembre-se da frase 'gato vermelho', que significa 'redução no cátodo'.
Mas como determinamos qual eletrodo é qual? Podemos usar o potencial de redução padrão dos eletrodos de metal, que é uma medida da tendência do metal de perder elétrons. Quanto maior o potencial de redução, maior a probabilidade de o metal ser reduzido.
Digamos que colocamos um eletrodo de prata e um eletrodo de cobre em uma célula eletrolítica e aplicamos corrente. O eletrodo de prata tem um potencial de redução maior, por isso é reduzido e atua como cátodo. O eletrodo de cobre tem um potencial de redução menor, por isso é oxidado e atua como ânodo.
Um uso das células eletrolíticas é realizar galvanoplastia, que é uma reação em que um metal é oxidado e depois reduzido na superfície de outro metal. Como o ânodo é o metal que faz o revestimento, em nosso exemplo de prata e cobre, o eletrodo de prata é revestido com uma fina camada de cobre.
Neste laboratório, você montará uma célula eletrolítica e executará o processo de galvanoplastia revestindo uma chave de latão com cobre.

células eletrolíticas

eletroquímica
A eletroquímica é um ramo da química que descreve e mede a relação entre a energia elétrica e uma mudança química. As reações eletroquímicas envolvem o movimento de elétrons de uma espécie para outra. Se a reação for espontânea, pode resultar em uma corrente gerada. Se a reação não for espontânea, ela pode ser impulsionada pela aplicação de corrente.
Em eletroquímica, a reação chave é a reação de oxidação-redução, chamada de reação redox. A reação redox é composta por duas meias-reações; oxidação, onde uma substância conhecida como agente redutor perde elétrons, e redução, onde uma substância conhecida como agente oxidante ganha elétrons. As reações redox sempre ocorrem em pares e alteram os estados de oxidação dos átomos nas moléculas envolvidas. Uma maneira fácil de lembrar essa relação é através do pneumático PLATAFORMA DE PETRÓLEO: Oxidação Is Losing – Reduction Is Gaining.
<h4>Células eletroquímicas</h4>
Uma célula eletroquímica é usada para medir ou induzir reações eletroquímicas. Consiste em vários componentes: uma câmara contendo a solução de reação, dois eletrodos condutores, uma solução eletrolítica condutora e um circuito externo. Existem dois tipos de células eletroquímicas. Uma é a célula eletrolítica, que usa energia elétrica para conduzir uma reação não espontânea. Nesse tipo de célula, a energia elétrica é fornecida por uma fonte de energia externa.
O outro tipo de célula é a célula galvânica, que usa uma reação eletroquímica espontânea para gerar energia elétrica. Os dois eletrodos são chamados de ânodo e cátodo e são conectados por um circuito externo. A câmara de reação é preenchida com um eletrólito, que promove a passagem de íons entre os eletrodos. Em uma célula galvânica, uma ponte de sal completa o circuito, permitindo o fluxo de íons entre as soluções eletrolíticas que contêm os eletrodos. Em uma célula eletrolítica, não há ponte salina, pois os eletrodos estão normalmente na mesma solução eletrolítica.
A reação de redução ocorre no cátodo, enquanto a reação de oxidação ocorre no ânodo. Isso é lembrado usando o mnemônico "Gato Vermelho", o que significa que a ucção vermelhaocorre no casco do gato.
<h4>Células eletrolíticas</h4>
As células eletrolíticas normalmente têm uma câmara de reação, que contém a solução eletrolítica. O eletrólito é geralmente uma solução aquosa contendo íons ou sais dissolvidos. Os íons no eletrólito promovem o movimento de íons e elétrons através da solução.
Quando uma tensão externa é aplicada, os íons no eletrólito são atraídos para o eletrodo com a carga oposta. É aqui que ocorrem as duas meias-reações. O ânodo perde elétrons durante a oxidação, enquanto o cátodo ganha elétrons durante a redução.
A galvanoplastia é um processo que utiliza uma célula eletrolítica. A galvanoplastia é um processo no qual um metal é depositado na superfície de um eletrodo, normalmente outro metal. O cientista britânico Michael Faraday, que deu nome à constante de Faraday, demonstrou a relação molar entre o íon carregado de galvanoplastia e a corrente elétrica. Olhando para a meia-reação específica para a prata:
Ag+ + 1e- → Ag
É preciso um mol de elétrons, fornecido por uma corrente externa, para reduzir um mol de cátions de prata em prata sólida. Portanto, a estequiometria de meia reação pode determinar a quantidade de material de galvanoplastia utilizado para uma quantidade conhecida de elétrons. Lembre-se de que a eletricidade não é medida com base nos moles dos elétrons, mas sim em coulombs, em homenagem ao engenheiro francês Charles-Augustin de Coulomb. Um coulomb é a quantidade de eletricidade que passa por um circuito elétrico com 1 ampere de corrente em um intervalo de 1 segundo, que se baseia no princípio da carga elementar, ou a carga que um único elétron carrega. A descoberta de Faraday pode, portanto, relacionar a corrente elétrica ao número de moles que o cátion de galvanoplastia revestiu:
F = eNA
F é a constante de Faraday, e é a carga elementar expressa em coulombs e NA é o número de Avogadro. Portanto, a constante de Faraday é expressa como o número de coulombs por mol e tem um valor de 96485 coulombs por mol. Como podemos determinar o número de moles de elétrons transferidos para o ânodo com base na corrente? Usando a carga elétrica Q e a definição de um coulomb:
<img src="/CDNSource/lm/labs/53/53_Concepts_3.jpg">
A carga elétrica é igual à corrente em amperes (I) multiplicada pelo tempo em segundos que a corrente foi deixada fluir. Dividindo Q pela constante de Faraday, que tem as unidades de coulombs por mol:
<img src="/CDNSource/lm/labs/53/53_Concepts_4.jpg">
Essa equação nos permite calcular o número de moles de elétrons e, assim, determinar quanto do cátion galvanoplastia foi reduzido.
Na galvanoplastia, o metal compreende as placas anódicas, ou cobre, o cátodo em uma fina camada de metal. A quantidade de metal chapeado depende da quantidade de corrente aplicada, bem como do número de moles do cátion de galvanoplastia disponível. Uma fonte de energia externa, como uma bateria, induz o fluxo de elétrons do ânodo para o cátodo e do terminal positivo para o terminal negativo da bateria.
Por exemplo, considere uma célula eletrolítica com um eletrodo de cobre, uma chave de latão atuando como o segundo eletrodo e uma solução aquosa de sulfato de cobre como eletrólito. Aqui, o cobre do eletrólito e do eletrodo de cobre é revestido na chave de latão.
Para que o metal de cobre seja depositado na chave de latão, o eletrodo de cobre sólido deve ser oxidado para formar íons de cobre. Em seguida, os cátions de cobre do eletrodo e do eletrólito são reduzidos da solução para formar cobre sólido na chave de latão.
2+ + 2e- →(s)
Os elétrons para a reação são recebidos do terminal negativo da bateria. Assim, a reação de redução ocorre na chave de latão, enquanto a reação de oxidação ocorre no eletrodo de cobre. A solução concentrada e acidificada de sulfato de cobre aumenta a solubilidade; Portanto, quanto maior a concentração da solução, menor a resistência e maior a corrente. Por sua vez, uma corrente mais alta permite que ocorra mais revestimento dos íons de cobre na chave de latão.
Alguns metais têm uma tendência maior do que outros a perder elétrons. O potencial de eletrodo padrão (E°) de uma substância é a medida da tendência de uma substância de perder elétrons. O metal com maior potencial de redução tem a maior tendência a perder elétrons; assim, é galvanizado primeiro.
<h4>Referências</h4>
<ol>
	<li style="font-size:17px">Kotz, J.C., Treichel Jr, P.M., Townsend, J.R. (2012). Química e Reatividade Química. Belmont, CA: Brooks / Cole, Cengage Learning.</li>
	<li style="font-size:17px">Silberberg, M.S. (2009). Química: A Natureza Molecular da Matéria e da Mudança. Boston, MA: Colina McGraw.</li>
</ol>

Células eletrolíticas

Eletroquímica
A eletroquímica é um ramo da química que descreve e mede a relação entre a energia elétrica e uma mudança química. ...

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JoVE Lab Manual

Chemistry

Química

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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Separating Mixtures with Solvent Extraction

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Separating Liquids by Boiling Point

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Steam Distillation and Low-Boiling Separation

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TLC Separation by Polarity

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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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