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Redox In the Conservation of Copper experiment you were given a brief introduction into redox reactions. Recall, redox reactions are also known as oxidation-reduction reactions. In particular, oxidation and reduction are the terms used to describe the transfer of electrons in a chemical reaction. Continuing, atoms or ions that accept electrons in a reaction are said to have undergone reduction, while species associated with the loss of electrons undergo oxidation. These two definitions can be remembered by two different mnemonics:
When an oxidation and a reduction reaction are paired together, electrons can flow from the oxidized species (losing electrons) to the reduced species (gaining electrons). The reduced species is often referred to as the oxidizing agent, or oxidant, as it is taking electrons away from another atom or ion. Conversely, the species being oxidized is referred to as the reducing agent or reductant, because it is donating electrons. In further detail, the oxidation or reduction of an atom or ion can be followed by observing the species’ oxidation number on the reactant side (left) and seeing if it changed on the product side (right). The rules for assigning oxidation numbers are as follows, but be sure you apply them in the order they are given.
Electricity and Voltaic Cells In general, the transfer of electrons from one place to another is commonly referred to as electricity. In fact, the electricity that you use to power your iPod® is simply the flow of electrons down a metal wire. As you will learn, another way to produce electricity is through the spontaneous redox reactions that occur in voltaic cells. In detail, the oxidation and reduction processes are separated, while the transfer of electrons occurs through an external wire. These ‘separated’ parts of the cell are called half-cells, one for oxidation and one for reduction. As previously stated, the construction of a voltaic cell takes advantage of the spontaneous nature of a redox reaction. For example, let’s look at the following reaction where zinc metal is oxidized and the copper ion is reduced.
If we were to construct the characteristic voltaic cell, we would need 1.0-M solutions of copper nitrate (Cu(NO3)2) and zinc nitrate (Zn(NO3)2). Additionally, we would have to suspend a piece of copper metal in the container holding the Cu(NO3)2 solution. Similarly, a piece of zinc would have to be held in the Zn(NO3)2 as shown in the animation above. In particular, both metals in this setup serve as electrodes, which act as the conductor used to initiate contact with the ions in solution. Overall, the piece of zinc metal is the negative electrode, referred to as the anode, while the piece of copper metal is the positive electrode and termed the cathode. As you can see from the figure above, both solutions are then connected to a third solution which acts as a salt-bridge. Finally, when the two electrodes are connected by means of an external metal wire, usually a voltmeter, the circuit is complete and a current of electrons can be observed. When a voltaic cell is constructed correctly, oxidation always occurs at the anode, while reduction always occurs at the cathode. A little hint to help you differentiate these two components is that oxidation and anode both start with vowels, while cathode and reduction both start with consonants. Thus, the flow of electrons should always have a positive potential value since they are flowing from negative (anode) to positive (cathode). In the case that you do obtain a negative potential value, which indicates that your electrons are flowing in the opposite direction, all you have to do is to switch the electrode connections on the voltmeter. In order to further understand what we have been talking about thus far, let’s break down our zinc-copper cell a little further and look at the reactions taking place at the individual electrodes. Recalling that oxidation occurs at the anode and reduction occurs at the cathode, we can break down our reaction into two half-reactions:
We can then join these two half reactions together in order to generate the overall reaction given previously in this discussion. Overall, the potential of this cell is 1.1-V as measured by the voltmeter. Standard Cells A cell in which all the materials are in their standard states at 25°C is termed a standard cell. In further detail, the standard states of elements and compounds are the most stable form found at 25°C. For gaseous substances the standard state is referred to at 1.00 atm of pressure and 25°C, while for solutions the standard state is 1.00 M. Voltaic Cell Representation Rather than writing out long chemical equations to represent redox reactions occurring in voltaic cells, a short-hand form has been devised using the following rules.
Calculations of a Cell’s Potential A term left unmentioned in the previous sections was the overall cell potential (EoCell), which in brief is simply the sum of the standard reduction potential (EoRed) and the standard oxidation potential (EoOxid). Mathematically, the overall cell potential is expressed as:
Using the example shown in our diagram above, let’s make a reasonable hypothesis as to what we believe our overall potential will be for this particular cell. In the figure, we see that Zn is being oxidized at the anode, while Cu2+ is being reduced at the cathode, as shown from the half reactions provided below.
From Appendix 3, we obtain that the EoRed for Cu2+ is +0.34 V, while the EoOxid for Zn is +0.76 V, thus allowing us to arrive at an estimated overall cell potential.
However, when you construct this cell in the experiment the reading from your voltmeter may not be exactly +1.10 V. But remember, you are not working at standard conditions. Still, this is not the most important thing that needs to be considered—the most important condition that needs to be met is that the ratio of concentrations of the corresponding salts is 1:1. If this ratio is met, any differences in the readings could be rationalized to stem from other sources such as resistance in the salt bridge or lack of calibration of the voltmeter.
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Redox Reactions in Voltaic Cells: Construction of a Potential Series
