Electrochemical Reactions
Electrochemical reactions
In a previous tutorial, we mentioned an experiment that shows what happens when zinc granules are added to a solution of copper(II) sulfate.
When a sheet of zinc is placed in an aqueous solution of copper(II) sulfate, solid copper forms on the zinc sheet (Screenshots taken from a video by Aaron Huggard on Youtube)
Optional Video: Copper Sulfate + Zinc
In the experiment, the \(\text{Cu}^{2+}\) ions from the \(\color{blue}{\textbf{blue copper(II) sulfate}}\) solution were reduced (gained electrons) to copper metal, which was then deposited as a layer on the solid zinc. The zinc atoms were oxidised (lost electrons) to form \(\text{Zn}^{2+}\) ions in the solution. \(\text{Zn}^{2+}(\text{aq})\) is colourless, therefore the blue solution lost colour. As discussed in another lesson, the half-reactions are as follows:
\(\text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-}\) \(\to\) \(\text{Cu}(\text{s})\) (\(\color{red}{\text{reduction half-reaction}}\))
\(\text{Zn}(\text{s})\) \(\to\) \(\text{Zn}^{2+}(\text{aq}) + 2\text{e}^{-}\) (\(\color{blue}{\text{oxidation half-reaction}}\))
The overall redox reaction is:
\(\text{Cu}^{2+}(\text{aq}) + \text{Zn}(\text{s})\) \(\to\) \(\text{Cu}(\text{s}) + \text{Zn}^{2+}(\text{aq})\)
Solid zinc loses two electrons to form zinc ions (\(\text{Zn}^{2+}\)) in an aqueous solution of copper(II) sulfate. The copper ions (\(\text{Cu}^{2+}\)) gain two electrons and deposit as solid copper. (Photos by benjah-bmm27 and Jurii on wikipedia)
Remember that there was an increase in the temperature of the reaction when you carried out this experiment (it was exothermic). An exothermic reaction releases energy. This raises a few questions:
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Is it possible that this heat energy could be converted into electrical energy?
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Can we use a chemical reaction with an exchange of electrons, to produce electricity?
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If we supplied an electrical current could we cause some type of chemical reaction to take place?
The answers to these questions are the focus of this tutorial:
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The energy of a chemical reaction can be converted to electrical potential energy, which forms an electric current.
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The transfer of electrons in a chemical reaction can cause electrical current to flow.
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If you supply an electric current it can cause a chemical reaction to take place, by supplying the electrons (and potential energy) necessary for the reactions taking place within the cell.
These types of reactions are called electrochemical reactions. An electrochemical reaction is a reaction where:
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a chemical reaction creates an electrical potential difference, and therefore an electric current in the external conducting wires
or
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an electric current provides electrical potential energy and electrons, and therefore a chemical reaction takes place
Definition: Electrochemical reaction
An electrochemical reaction involves a transfer of electrons. There is a conversion of chemical potential energy to electrical potential energy, or electrical potential energy to chemical potential energy.
Electrochemistry is the branch of chemistry that studies these electrochemical reactions. An electrochemical cell is a device in which electrochemical reactions take place.
Definition: Electrochemical cell
A device where electrochemical reactions take place.
There are two types of electrochemical cells we will be looking more closely in this tutorial: galvanic and electrolytic cells. Before we go into detail on galvanic and electrolytic cells you'll need to know a few definitions:
Definition: Electrode
There are two types of electrodes in an electrochemical cell, the \(\color{blue}{\textbf{anode}}\) and the \(\color{red}{\textbf{cathode}}\).
\(\color{blue}{\text{Oxidation}}\) always occurs at the \(\color{blue}{\textbf{anode}}\) while \(\color{red}{\text{reduction}}\) always occurs at the \(\color{red}{\textbf{cathode}}\). So when trying to determine which electrode you are looking at first determine whether oxidation or reduction is occurring there. An easy way to remember this is:
Tip:
Oxidation is the loss of electrons.
Reduction is the gain of electrons.
|
\(\color{blue}{\textbf{O}}\)xidation \(\color{blue}{\textbf{i}}\)s \(\color{blue}{\textbf{l}}\)oss of electrons |
\(\color{blue}{\textbf{OIL}}\) |
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\(\color{red}{\textbf{R}}\)eduction \(\color{red}{\textbf{i}}\)s \(\color{red}{\textbf{g}}\)ain of electrons |
\(\color{red}{\textbf{RIG}}\) |
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\(\color{blue}{\textbf{O}}\)xidation is \(\color{blue}{\textbf{l}}\)oss of electrons at the \(\color{blue}{\textbf{a}}\)node |
\(\color{blue}{\textbf{An Ox}}\) |
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\(\color{red}{\textbf{Red}}\)uction is gain of electrons at the \(\color{red}{\textbf{cat}}\)hode |
\(\color{red}{\textbf{Red Cat}}\) |
Table: A summary of phrases useful to help you remember the oxidation and reduction rules.
The electrode is placed in an electrolyte solution within the cell. If the cell is made up of two compartments, those compartments will be connected by a salt bridge.
Definition: Electrolyte
An electrolyte is a solution that contains free ions, and which therefore behaves as a conductor of charges (electrical conductor) in solution.
Definition: Salt bridge
A salt bridge is a material which contains electrolytic solution and acts as a connection between two half-cells (completes the circuit). It maintains electrical neutrality in and between the electrolytes in the half-cell compartments.
This lesson is part of:
Electrochemical Reactions