Batteries are classified as either primary or secondary. The easiest way of thinking about this is that secondary batteries are rechargeable. By reversing the current through the cell, a secondary battery can be recharged and used again. That is, electrical energy is used to restore a secondary cell back to its original state, whereas a primary battery is used only once—the reactions cannot be easily reversed. Both types are common, as you will recognize from your experience with batteries. Primary batteries usually have higher specific energy [W·h·kg−1] and power [W·kg−1] than secondary batteries, although that is not always the case. Also, keep in mind that the distinction between primary and secondary batteries is often based on design rather than on chemistry alone. For instance, the silver–zinc and alkaline battery chemistries are each used in both primary and secondary (rechargeable) batteries.
CELL VERSUS BATTERY
A cell is the basic electrochemical device that stores energy. Two or more cells connected together form a battery. It is common to refer to a single cell as a battery even though there is only one cell.
A special category of primary batteries is the reserve battery. In these batteries, a key component of the cell, such as the electrolyte, is kept separate from the rest of the components or is otherwise inactive. By doing this, most self-discharge and side reactions (Section 7.8) are eliminated, leading to an extremely long shelf life under even severe environmental conditions. When it is needed, the reserve battery is activated. Only then are all of the components able to function together. The most common reserve batteries are thermally triggered. When the cell is heated, the electrolyte, which is a solid at room temperature, melts and forms an ionically conductive media that enables operation of the battery.
The chemistries for a number of common batteries are listed in Table 7.1. There is a rich diversity of reactions used in batteries, and a full discussion is beyond the scope of this chapter. Nonetheless, we will find it instructive to consider a few battery chemistries in order to develop a physical understanding of what is happening during the operation of a battery cell. This understanding will help us describe the performance of these cells and will provide a foundation for electrode design (see Chapter 8). Two important classes of electrode reactions for batteries are described by Huggins: reconstruction and insertion.
Table 7.1 Example Chemistries for a Number of Common Batteries
Cell Electrochemical reactions Nominal Cell potential Comments
Lead–acid 1.8 V Secondary
During discharge, lead sulfate is formed at both electrodes. The electrolyte is part of the active material
Alkaline 1.5 V Primary
Common battery
Lithium-ion 3.5–4.2 V Secondary
The spinel manganese dioxide is just one of the chemistries of rechargeable lithium-ion batteries
Silver–zinc 1.5 V Manufactured as both a primary and a secondary battery
Ni/Cd 1.2 V Secondary, out of favor because of environmental issue with cadmium
Ni/Fe 1.2 V Secondary, also called the Edison cell
NiMH 1.2 V Secondary
Na-S 1.9 V Secondary high temperature
Li-SO2 3 V Primary
Lithium–sulfur dioxide
Li-SOCl 3.4 V Primary
Lithium thionyl chloride
Li/FeS2 1.5 V Reserve
Thermally activated
LiCl/KCl eutectic for electrolyte
Mg/AgCl 1.4 V Reserve water activated
Reconstruction (Also Referred to as Conversion Reactions)
As the reaction proceeds, new phases are formed and grow, while other phases shrink and disappear. One or more electrons are transferred, and the reactants and products are distinct stoichiometric species. Reconstruction reactions can further be designated as formation or displacement reactions.
An example of a reconstruction/formation reaction is that found in the lithium sulfur dioxide primary cell. The overall reaction is lithium and sulfur dioxide reacting to form solid lithium dithionite.
Of course, in a battery the overall reaction proceeds as two separate electron-transfer reactions. At the negative electrode, lithium metal is oxidized; and at the positive electrode, lithium ions react with sulfur dioxide dissolved in an organic electrolyte.
Thus, a new solid phase, lithium dithionite, forms and grows during the discharge process. Note that a key reactant, SO2, is found in the electrolyte, a feature that is present in a number of batteries.
Displacement is a second kind of reconstruction reaction.
An example of a displacement reaction is the silver oxide cell discussed earlier.
Referring back to the half-cell reactions discussed previously, the equilibrium potential for the Zn (−1.251 V) reaction is more negative than that of Ag (0.340 V). Therefore, Zn will displace the Ag in the oxide spontaneously. For the electrochemical cell, of course, these reactions occur as separate electron-transfer reactions as we have already seen.
Although useful in categorizing chemistry of the cell, these designations of formation or displacement do not provide the mechanistic details of the reaction. Often it is necessary to have a more in-depth picture of the electrode reactions in order to explain the behavior of cells. The physical phenomena associated with reconstruction reactions can be complicated and include dissolution and precipitation, solid-state ion transport, and film formation. Because of its importance in a number of batteries, it is worth taking time to explore the dissolution–precipitation mechanism in more detail. In this mechanism, a reaction product is formed that may subsequently react with a species in solution to form a precipitate. For example, a metal (M) used as the negative electrode of a cell (see Table 7.2) may dissolve as a result of electron transfer:
This electrochemical reaction may be accompanied by the subsequent precipitation of a salt from the solution
The amount precipitated depends on the solubility of MXz in the solvent. Other reactions besides metal dissolution can result in the dissolved product and subsequent precipitation. We will examine the behavior under three solubility conditions: (1) an insoluble salt, (2) a highly soluble salt, and (3) an intermediate or sparingly soluble salt. Each case will be explored through specific chemistries.
Table 7.2 Common Negative Electrode Materials for Batteries. Values Are Theoretical
Element MW [g·mol−1] Standard potential [V] Density [g·cm−3] Valence Specific capacity [A·h·g−1] Volumetric capacity [A·h·cm−3]
Li 6.941 −3.01 0.534 1 3.86 2.06
Na 22.99 −2.714 0.971 1 1.16 1.12
Mg 24.305 −2.4 1.738 2 2.20 3.8
Al 26.98 −1.7 2.699 3 2.98 8.1
Fe 55.845 −0.44 7.86 2 0.96 7.5
Zn 65.39 −0.763 7.13 2 0.82 5.8
Cd 112.41 −0.403 8.65 2 0.48 4.1
Pb 207.21 −0.126 11.34 2 0.26 2.9
For an example of an insoluble salt, we examine the positive electrode of a lithium thionyl chloride battery. From Table 7.1, the reaction is
Chloride is produced as a result of reduction. The electrolyte is a mixture of LiAlCl4 (lithium tetrachloroaluminate) and SOCl2 (thionyl chloride), which serves as both a reactant and part of the electrolyte for the cell. Each time a chloride ion is produced at the positive electrode, it will encounter a lithium ion in the electrolyte. For all intents and purposes, LiCl is insoluble in the electrolyte and precipitates out. Consequently, LiCl(s) is likely to be found close to where the thionyl chloride is reduced. Because of the insolubility of the salt, there is essentially no Cl− in solution; therefore, the reverse of the dissolution reaction cannot occur. Thus, because of the poor solubility of the discharge products, the reactions are not easily reversed and the LiSOCl2 is a primary battery.
Next, let’s consider the zinc electrode, which provides an example of dissolution to form a highly soluble salt. The reaction takes place in an alkaline electrolyte and follows a dissolution–precipitation mechanism:
The ultimate product ZnO(s) has a relatively high solubility in aqueous KOH. For now, let’s imagine that the solubility is so high that the precipitation reaction doesn’t occur at all. At first, this doesn’t appear to be a problem. When we want to recharge the cell, zincate ions are readily available in solution to be reduced back to zinc. However, in contrast to the LiCl in the previous example, which is expected to be located at the site where the reduction occurs, the soluble zincate can diffuse or be transported by convection to another part of the cell. Now imagine, as is inevitably the case, that the current distribution on the electrode is nonuniform. The sites where deposition is higher will preferentially grow in size leading to changes in the shape of the electrode and possible dendrites (branching tree-like crystals). The shape changes and dendrites can cause shorting in the cell and lead to premature failure. Consequently, cells using a zinc negative electrode behave in a mediocre fashion as rechargeable batteries.
The final example is the negative electrode of the lead–acid battery, where lead reacts to form lead sulfate, a nonconducting ionic solid that is sparingly soluble.
As in the previous examples, the first step on discharge is an electron-transfer reaction and dissolution. In this instance, the lead sulfate is sparingly soluble in the electrolyte. The solubility is small enough so that during cycling the lead sulfate product is repeatedly formed and dissolved at roughly the same location. Yet, the solubility is large enough so that the reverse reaction can occur. On charging, not only are the kinetics for the electron-transfer reaction important, but the rate can be limited by dissolution of the lead sulfate and mass transfer of Pb2+ to the lead surface. This situation is explored in Problem 7.26.
Insertion
A second important class of electrode reactions found in batteries is the insertion reaction. Here there is a relatively stable host material, and the guest material is found in unoccupied sites of the host material. When the host material has a layered structure, the insertion is called intercalation. In the case of intercalation, the host material retains its essential crystalline structure and the process can be highly reversible. This process is illustrated by the intercalation of lithium into titanium disulfide (Figure 7.3). TiS2 has a layered structure, where individual sheets of TiS2 are held together by van der Waals forces. In contrast to the reconstruction reactions, the lithium occupies the space between the layers of titanium disulfide, which is called the gallery space, rather than creating a new phase. The intercalation of lithium into the host titanium disulfide, a positive electrode material for a lithium-ion cell, is written as
where x varies from zero to one. When x is zero, the TiS2 is delithiated and the electrode is considered charged; when x is one, the host is lithiated, representing a discharged state. The product of this insertion process results in a kind of solid solution of the guest species in the host, where lithium occupies sites in the gallery. In these instances, a characteristic sloping discharge curve results (Figure 7.5).
Figure 7.3 Insertion of lithium (solid black circles) into the gallery space of titanium disulfide.
There are many types of intercalation and insertion reactions. The interested reader should consult the Further Reading section at the end of the chapter for additional information. Insertion sometimes occurs in stages, where there are different sites within a layer or even preferential filling of a stack of layers. Here, the guest material is not randomly distributed between the layers; instead, occupation of specific low-energy interstitial sites is favored. In these cases, the equilibrium potential can have multiple plateaus and steps. We write the reaction for the lithiation of the material as if lithium were reduced to a valence of zero; however, it may be the oxidation state of the metal (e.g., Mn) in the host material that changes.
The development of lithium rechargeable batteries has brought the importance of insertion reactions to the forefront. For lithium-ion cells, positive electrode materials are generally transition metal oxides or phosphates. The predominant negative electrode material is carbon. Finally, it should also be noted that metals other than lithium can undergo insertion reactions.
In this section, we have examined reconstruction and insertion reactions, both of which are important for practical battery systems. You have also been introduced to several different types of battery chemistries. This short introduction is by no means complete, but it is hoped that it has helped you to gain new insights into batteries and battery chemistries. We encourage you to consult other sources, such as those listed at the end of the chapter, for additional information.
Some standard metrics provide a basis for comparison of different electrode and battery systems. For example, Table 7.2 contains information for elements commonly used as negative electrode materials. Key metrics are found in the last two columns: the specific capacity [A·h·g−1] and the volumetric capacity [A·h·cm−3]. The potential of the electrode is also critical, and low values for the negative electrode potential contribute to higher cell voltages. Given its small mass and low potential, it is clear why lithium is the customary choice for high-energy batteries. When the volumetric capacity is considered, other materials compare more favorably to lithium, for example, magnesium and aluminum. Note that both of these metals have a valence greater than one. The standard potentials in this table refer to the oxidation of the elements, and do not necessarily reflect the actual reactions in a practical battery.
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