Introduction

We have previously seen reactions involving the transfer of one or more electrons from one species to another, typically called oxidation-reduction reactions or, more colloquially, redox reactions. Since redox reactions involve the transfer of electrons, they are also accompanied by changes in the oxidation states of the reactant and product. The reactant that donates the electron is dubbed the reductant, while the species that accepts the electron is dubbed the oxidant.

We often think of redox reactions as occurring through two half-reactions. Specifically, we can consider the oxidation half-reaction and the reduction half-reaction. As an example, let’s consider a reaction that you did last semester as part of the “Copper Cycle” experiment:

which can be rewritten as the net ionic equation

This is an example of a single displacement reaction in which the solid zinc gives up two electrons to the copper ion to form solid copper and a zinc ion. Since we know that the zinc is being oxidized and the copper is being reduced, we can write the individual half-reactions as in equations 2 and 3.

It’s important to notice here that the chloride ions have been left out of the half-reactions. This is because the chloride ions are not important to the reaction and are merely included simply for charge balancing purposes. Additionally, if you add the two half-reactions together, you’ll notice that the electrons will “cancel out” and you’ll be left with just the net ionic equation. Remember that no half-reaction can occur by itself — both half-reactions are occurring simultaneously, because the electrons have to travel from the neutral zinc species to the ionic copper species as outlined in Figure 1.

One important aspect to consider regarding redox reactions is the reduction potential of the species undergoing electron transfer. Most often measured in volts, V, reduction potential measures the tendency for a redox reaction to occur. Under standard conditions, 25°C, and 1.0 M concentration of the aqueous ions, the measured voltage of the reduction half-reaction is defined as the standard reduction potential, E°. Thankfully, standard reduction potentials are well-tabulated and readily available from online sources and most chemistry textbooks/review papers.

When you are looking at tabulated standard potentials, it’s important to keep in mind that the values are determined with respect to that ion’s ability to accept electrons (be reduced). To better understand this, let’s go back to our example with zinc and copper. The only difference here is that we will need to rewrite equation 4 to show the zinc ion being reduced by two electrons.

Now we now how the standard reduction potentials compare to one another! But…what do these actually mean? The values are determined as compared to a standard hydrogen electrode (SHE), meaning negative values are stronger reducing agents than an SHE, while positive values and stronger oxidizing agents than an SHE. Therefore, by comparing equations 5 and 6, we can say that copper is the stronger oxidizing agent, meaning Cu²⁺ is more likely to accept electrons, while Zn²⁺ is more likely to donate electrons.

In the “Copper Cycle” experiment from last semester, this reduction was carried out in a single solution, resulting in a very fast reaction that forms solid copper. If we want to utilize the electric potential, the electrons being transferred must be channeled through a conducting wire so that the current can be controlled. This is the basis of a voltaic cell (also referred to as a galvanic cell). A voltaic cell consists of two separated aqueous solutions of the ions involved in the half reactions that are connected by a salt bridge. Additionally, each solution contains a solid, metallic form of the species in the half reactions that are connected by the conducting wire (Figure 2).

The salt bridge can be comprised of numerous objects — oftentimes the bridge will be a gelled salt solution, but can be as simple as a piece of string soaked in a salt solution. The important factor is that the bridge must contain a salt that does not appear in either of the half-reactions of interest. The electrolyte-soaked bridge allows charge to travel from one solution to the other.