Electrochemistry 1: Redox reactions

27th Dec 2020

Electrochemistry is one of the few topics in chemistry, presented at this level, which can be readily applied and explored directly in the laboratory. As you work through this series and indeed from reading other sources, think about how you would test the ideas through practical work.

Before reading this series, you should already be aware of the ideas of half equations and electrolysis, and know a few examples of electrolytes and electrodes. I will mention some of the connections between electrochemistry, spontaneity and chemical equilibrium appropriate to this standard along the way.

The main theme of this series is the study of a cell which generates an electric current as a result of a spontaneous reaction. This new type of cell differs from the cells used to carry out electrolysis. The differences can sometimes confuse students so I will attempt to provide a gradual introduction. Note that this will not be a complete overview of the topic and is provided to supplement your own learning materials. I will not outline Faraday's Laws or electrochemical cell applications, such as fuel cells or rechargeable cells, since these aspects are well explained elsewhere.

The anode and the cathode

Before we continue, let us re-state some of the key terms. The cathode is an electrode where reduction takes place. The anode is an electrode where oxidation takes place. The mnemonic "RedCAT" or "REDuction at the CAThode" is often provided to help students remember this.

The definition of the cathode and anode is based only on the type of electron transfer taking place at the time that an electric current is flowing (the circuit is closed). The definition is not related to the electric charge on the electrode. This means that the cathode is not always negatively charged and, likewise, the anode is not always positively charged. In the case of electrolysis, the cathode is negatively charged and the anode is positively charged however, as we will soon see, this is not the case for other types of electrochemical cells.

Electrolytic cells

An electrolytic cell is a type of electrochemical cell where an external power supply drives electrolysis, an example of a non-spontaneous process. Cast your mind back to the electrolysis of water, for example, where an electric current is applied to water. Hydrogen is produced at the cathode and oxygen is produced at the anode (Figure 15.1).

Electrolysis of water Figure 15.1 An electrolytic cell

Overall, protons are reduced to hydrogen and hydroxide is oxidised to oxygen. The chlor-alkali process, the extraction of aluminium from bauxite or the electrolytic purification of copper are also valid examples where electrolytic cells are applied.

In regards to spontaneity, note that the conversion of water to hydrogen and oxygen is non-spontaneous at atmospheric pressure and room temperature. During electrolysis, an external power supply is driving an electric current which causes a non-spontaneous reaction to occur. In contrast, we know that the reaction of hydrogen with oxygen is spontaneous.

We now move on to a different class of electrochemical cells where a spontaneous reaction produces an electric current, that is, the 'mirror-image' of an electrolytic cell.

Voltaic (or galvanic) cells

A voltaic cell, also known as a galvanic cell, is an electrochemical cell in which a spontaneous reaction generates an electric current. Examples of voltaic cells include lithium-ion batteries, fuel cells, or rechargeable batteries in their discharge state.

The "RedCAT" definition of the anode and cathode still applies here; however, the charges are opposite to that found with electrolytic cells. In other words, the cathode is positively charged and the anode is negatively charged. To see why, first we need to understand what is going on at the atomic level. This will be outlined shortly.

Let us first summarise the main points in the table below before getting to the details. If you have difficulty recalling the main points, I suggest that you memorise the observations for one type of cell (probably the more familiar electrolytic cell) and then when asked to recall the other, just reverse or toggle the statements.

Cell type Reaction taking place Cathode charge Anode charge
Voltaic (or galvanic) Spontaneous Positive Negative
X ⇌ Y + energy
Electrolytic Non-spontaneous Negative Positive
X + energy ⇌ Y

Redox processes at the electrode

There are numerous theories which attempt to explain what is going on near the surface of each electrode in a voltaic cell, most of which are beyond the scope of this series. We will apply a simplified theory which explains what occurs when an electrode is immersed in an electrolyte.

Metals are generally easy to oxidise compared to many other species. A metallic electrode placed into water oxidises to some extent, leading to a layer of metal cations at the surface of the electrode, along with an equal build-up of electrons at the electrode. The more readily a metal oxidises, the greater the production of metal cations and free electrons. For example, copper would generate fewer ions near the electrode surface compared to zinc. We can visualise both systems side by side (Figure 15.2).

Oxidation of copper and zinc electrodes in water Figure 15.2 The build-up of free electrons and cations from the partial oxidation of metallic electrodes in water

It is probably worth emphasising here that the cations generated near the surface of the metal are, on average, positioned near or at the surface of the metal. They are probably not free to migrate as hydrated cations in water. It would not be possible to naturally generate a solution of hydrated cations (or indeed anions) because this would lead to considerable repulsion between species with charges of the same sign and require energy to maintain. I like to think of the electrode as a piece of metal with an oxidised layer, as cations stuck to the surface. The atoms/ions at the surface are constantly changing between oxidised and reduced forms.

Note that the oxidation of the metal leads to a state of higher entropy, in particular, the number of configurations of a partially oxidised metal electrode in water is higher than that of a pure metal electrode and water. As such, the partial oxidation of the metal is favourable up to a point. The metallic electrode will generally reach a state of dynamic equilibrium where the extent of oxidation is fixed. Free electrons would re-combine with the surface-bound metal cations while, at the same rate, the metal would oxidise. The equilibrium composition will differ for different metals (Figure 15.2).

The surface of the metal electrode does not necessarily oxidise completely, even at the surface, because this would lead to a much larger build-up of negative charge on the electrode and positive charge near the surface. This would result in a lot a repulsion present on the electrode and near the electrode surface. Any further oxidation of the metal would generate ions that have to overcome this repulsion. In other words, energy would be required to oxidise the electrode further (Figure 15.3).

Excessive oxidation at the electrode surface Figure 15.3 Continued oxidation of the metallic electrode would require energy to overcome the repulsion

The study of the nature and composition of the electrode surface is quite advanced. Related to this is the study of heterogeneous catalysis, a process in which a reaction is sped up by a catalyst that is in a different state or phase to the reactants and products. The description of the surface of iron in the Haber Process, nickel in the hydrogenation of alkenes or the electrodes of an electrochemical cell (be it voltaic or electrolytic) is complex and can easily form the basis of higher level research projects. The 'layered' description of cations and metal atoms at the electrode surface, while limited, will suffice for our needs.

Potential energy and polarity

We can identify the anode and cathode, and compare the extent of oxidation of two electrodes by comparing the Coulombic potential energy (see Figure 6.8). At this stage, many authors will correctly proceed to explain the need to measure the potential energy difference between the electrode and the electrolyte near the electrode surface. Then authors will state that it is physically impossible to achieve. These explanations are perfectly fine. For chemistry students at this level, I usually take a slightly different approach and prefer to focus on the different electron densities of each electrode and measure the potential energy using a test-electron. This is explained next.

One way of comparing the Coulombic potential energy of each electrode is achieved by seeing how a test-electron responds to the resident electrons on the electrode. A test-electron is a single, hypothetical electron used for test purposes and does not originate from the voltaic cell. The electrode will contain far more electrons (of the order of Avogadro's number) so the test-electron is not going to influence the combined negative charge generated by the electrode or affect any ongoing redox processes. After we place a test-electron on the surface of the metal, ask yourself, which test-electron would experience the most repulsion and have a potential energy of greater magnitude (Figure 15.4). Recall that the repulsive potential energy between electrons or between protons is always positive.

Using a test-electron to measure potential energy Figure 15.4 Finding the potential energy of a test-electron placed at the surface of an electrode

You can see that a test-electron would experience more repulsion when placed on the zinc electrode. This result leads to a number of key properties:

  1. The potential energy of the test-electron on the surface of zinc is greater in magnitude than the potential energy of the test-electron on the surface of copper.
  2. With further investigation, one can deduce the potential difference of two electrodes, that is, the difference of potential energy experienced by each test-electron. The determination and calculation of potential difference will be covered in the next article.
  3. If both electrodes were connected and assuming that electrons are free to migrate from one electrode to the other (this assumption is evidently wrong), then the electrons are more likely released from zinc, leading to the reduction copper cations present at the copper electrode surface. The oxidation of zinc and reduction of copper is a familiar reaction and one which is spontaneous under ambient conditions.
  4. Electrons flow from a region of higher potential energy (with more repulsion) to a region of lower potential energy.
  5. If both electrodes are identical, then the potential difference is zero and there would be no net flow of electric current (no preferred direction of electron flow).

Note that the process of electron or charge migration is simplified here. For our purposes, the electron released from a zinc atom will push neighbouring delocalised electrons on the zinc electrode and subsequently copper electrode so that eventually one of the copper electrons binds to a copper cation (Figure 15.5). This is also how one could view the motion of delocalised electrons in metals and graphite. A better explanation of the migration of charge would require a knowledge of electric fields, which is beyond the scope of this series.

Cascading repulsion of electrons and reduction of cations Figure 15.5 Zinc releases an electron which ultimately pushes other delocalised electrons along the electrodes, one of which binds to the cation and results in reduction (note: this only assumes that electrons can migrate, a wrong assumption)

Point 3 above helps us identify the cathode (copper) and the anode (zinc). In summary, the zinc metal is more likely to oxidise and lead to a build-up of negative charge on the electrode. Therefore, the anode is more negatively charged than the cathode. Finally, the copper cations are more easily reduced (compared to zinc cations) and the withdrawal of electrons from the copper electrode results in a lower electron density and fewer negative charges present. The cathode is therefore viewed as the positive electrode.

Knowing the relative potential energy of each test-electron, in short, reveals the polarity of the voltaic cell. The polarity describes which electrode is positive and which is negative. You will probably have been required to consider 'bond polarities' at some point. This sort of question is the same, only this time you would be asked which end of the bond is δ+ and which is δ-.

If you assembled the above voltaic cell in the laboratory then you would probably not observe the oxidation of zinc or reduction of copper cations. To enable electrons to flow freely we need to connect both solutions with a salt bridge, which is the next topic.

Completing the voltaic cell: salt bridges

I have used the above approach to describe something which resembles a voltaic cell, though this is not what you should construct because the circuit is still open. What we have ignored at this stage is the state of the electrolyte.

In practice, each electrode is not immersed in water but in an aqueous solution of its ions. For example, a strip of copper foil would normally be placed in a solution of copper(II) sulfate or some other water-soluble form of copper(II), at a concentration roughly between 0.001 M to 0.100 M.

When an electrode is placed in a solution of its ions (Figure 15.6a), any oxidation or reduction that takes place near the electrode surface would result in an imbalance of electrical charge. For instance, at the cathode there would be a decrease in cation concentration or, put another way, an increase in counter-anion concentration (Figure 15.6b). Similarly, at the anode there would be an increase in cation concentration or a decrease in counter-anion concentration (Figure 15.6b).

Imbalance of electric charge at the cathode and anode Figure 15.6 (a) Electrodes in a salt solution (b) Increase in anion concentration near the cathode/increase in cation concentration near the anode

One can address the imbalance of charge by connecting both electrolytes with a salt-bridge, a junction which maintains the neutrality of the electrolytes and enables electrons to migrate from the anode to the cathode (Figure 15.7).

The main component of the salt-bridge is a saturated solution of an inert salt. The salt should be chosen based on the electrodes and electrolyte involved. Ideally, the salt bridge salt should not take part in any of the ongoing redox reactions. The solution is saturated so as to maximise the conductivity of the junction and also minimise the mixing of the cathode and anode electrolytes. Recall that dissolved ions migrate from regions of higher concentration of salt to regions of lower concentration. Since the electrolyte concentrations are typically well below saturation concentrations, it means that the anode and cathode electrolyte ions are unlikely to migrate from the electrolyte to the salt bridge. If anything, the salt bridge ions would migrate from the salt bridge to each electrolyte.

The salt bridge Figure 15.7 Using a salt bridge to maintain electrical neutrality in each electrolyte. The sequence of steps is intended to demonstrate the action of the salt bridge and not intended to accurately represent the mechanism.

In a school environment, the salt bridge can be made up of a strip of filter paper soaked in the saturated solution, or, with another beaker containing the saturated solution and connecting the saturated solution to both electrolytes using wide strips of filter paper soaked in the saturated solution. It may also be necessary to choose an appropriate salt for the salt bridge which would not form a precipitate with either electrolyte and potentially block the migration of salt bridge ions.

Non-metallic reactants and products

We have focused on metallic electrodes and their hydrated ions e.g. copper(0) and copper(II). However, there are also examples of reactants and products which are non-metallic or where the zero-oxidation state of the metal is not part of the reaction. Some reagents are gaseous and may not dissolve in water very well. Examples of reagent pairs of interest include:

  • hydrogen gas and aqueous hydrochloric acid
  • potassium sulfide K2S solution and sulfur
  • aqueous mixtures of iron(II) chloride and iron(III) chloride (see note below**)

How would one transfer electrons and facilitate redox processes for the above paired species?

A solution to this problem is platinum black. As the name suggests, platinum black is black in appearance and has clearly observable catalytic properties. Platinum black is a powdered, high-surface area form of elemental platinum and differs in appearance to the moulded form of platinum, which is silver and shiny. Platinum is also quite expensive! The material can be formed into an electrode that is largely inert to the reactants and products involved, and one which resists corrosion and conducts electricity.

Platinum black provides an interface or, in other words, a meeting point where the oxidised and reduced forms of a reactant can transfer electrons. The reacting species adsorb or stick to the surface of the metal, in a similar way to how the nitrogen, hydrogen and ammonia stick to the surface of iron in the Haber process. A simplified description of electron transfer is given in Figure 15.8. In many ways, this is another example of heterogeneous catalysis, mentioned previously.

Action of platinum black Figure 15.8 The transfer of electrons mediated by platinum black electrode (electrolyte not shown). Electron build-up occurs on the anode and electron-withdrawal occurs on the cathode. The anode is assigned the negative pole and the cathode is assigned the positive pole.

**The latter case listed above highlights the use of platinum black as the electrode, not elemental iron. This consideration tends to confuse some students. When planning the design of voltaic cells which involve the redox reaction of mixed oxidation state metal ions like iron(II) and iron(III), you should not use iron(0) as the electrode because iron(0) is not part of the redox reaction. If the redox process between iron(0) to iron(II) is involved, then by all means use iron(0) as the electrode.

More general uses and future goals

You might be wondering why chemists go to such lengths to set up voltaic cells like those outlined here. One of the main reasons is that the assembly provides chemists with a way of measuring potential difference. In the next article we will study what potential difference represents and what factors affect potential difference.

Another, more practical reason why voltaic cells are set up like this is that it provides end-users (not just chemists) with a way of harnessing the energy generated by the spontaneous reaction. Remember that voltaic cells are just like other cells and batteries which are discharging.

Going back a little, placing a zinc electrode in a solution of copper(II) sulfate would result in the same chemical changes only this time there would be no way of harnessing the "electrical energy". The electrons would migrate from a zinc atom (the 'anode') directly to a copper cation (the 'cathode'). By setting up independent electrodes and electrolytes, chemists can temporarily isolate the electron migration and its energy. For example, one could power an LED by connecting its poles to the electrodes of the voltaic cell. You can experiment with other electronic components and construct cells in series or in parallel.

In this first part of the series, we have looked at the underlying microscopic processes which help to explain how voltaic cells work. We have also compared voltaic cells with electrolytic cells and noted some key differences regarding the polarity of the electrodes and whether energy is required or released. In the next article, we shall focus entirely on voltaic cells and the application of potential difference for chemists.