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Introduction

The Earth, a dynamic and interconnected system, relies fundamentally on the Sun as its primary energy source. All naturally occurring raw energy on our planet, whether directly or indirectly, stems from the radiant power of the Sun. From the growth of plants through photosynthesis to the formation of fossil fuels over millions of years, the Sun’s energy plays a pivotal role in shaping the Earth’s energy landscape. Wind patterns and water movements, too, are consequences of the Sun’s heating of the Earth’s surface. Over the course of evolution, humans have acquired the knowledge to harness and convert these raw natural energies into forms, such as electricity and heat, useful for sustaining daily life. This journey of transformation has resulted in various interchangeable forms of energy, which have become the foundation of our modern existence. In Figure 1.1, we show the six forms of energy that humans have learned to interchangeably transform, underscoring their essential role in supporting and shaping our daily lives.

However, the law of energy conservation requires that energy cannot be created nor destroyed. This means that even though energy changes its form, the total quantity of energy always stays the same. For example, many of the energy converters widely used today involve the transformation of chemical energy through thermal energy into electrical energy. The efficiency of such systems is limited by the fundamental laws of thermodynamics. As of today, the conversion efficiency from naturally occurred fossil fuels to useful electricity is ∼35% on average. Such a low efficiency implies a faster depletion of fossil fuels and more emissions of greenhouse gas (CO2) into the atmosphere. In this aspect, those direct energy-conversion devices, such as fuel cells and electrolyzers that bypass the intermediate step of chemical-to-thermal conversion, are advantageous. Unbound by the thermodynamic law, these direct energy converters can achieve high efficiency, thus less fuel consumption and carbon emissions in producing the same amount of electricity. Electrochemical cells are a representative class of such high-efficiency direct energy converters. They are currently deemed one of the best technologies to address all aspects of energy-related challenges such as emissions/pollutions, efficiency, intermittency, cost, and supply chains.

Figure 1.1 Six forms of energy and their interchangeable transformations.

1.1 Brief History of Electrochemical Cells

The history of electrochemical cells began with Italian physicist Alessandro Volta. In 1800, he demonstrated a so-called “voltaic pile”, see Figure 1.2 [1], in which Cu and Zn discs were separated by cardboard or felt spacers soaked in salt water (the electrolyte). This is believed to be the first prototype of the modern battery producing electrical current from chemical reactions. Following Volta’s pioneering work, British scientist Humphry Davy was the first to link the production of electricity with the occurring chemical reactions (precisely Gibbs free energy change of the reactions). His student Michael Faraday took a step further predicting how much product can be produced by passing a certain amount of electric current though a chemical compound, a process that he called “electrolysis” and later known as Faraday’s law. The above two important works laid the foundation for today’s electrochemical cells to produce power and chemicals.

Figure 1.2 Volta’s voltaic pile [1].

Source: acrogame/Adobe Stock.

Modern electrochemistry has now become a branch of chemistry studying the interplay of electrical and chemical energy. Governed by Faraday’s law, a large portion of this field deals with the study of changes in chemical energy (or Gibbs free energy) caused by the passage of an electrical current or vice versa, the production of electrical current by chemical reactions [2]. These basic laws of electrochemistry have been successfully applied to a wide range of fields from fundamental phenomena (e.g. electrophoresis and corrosion) and to technologies (batteries, sensors, fuel cells, water electrolyzers, smelters, and metal platers), making significant impacts on every part of our life and economy.

1.2 Configuration of Electrochemical Cells

In electrochemical cells, chemical and electrical energy can be reversibly transformed into each other. An electrochemical cell consists of three basic components: electrolyte, cathode, and anode. Electrolyte is an ionic conductor and electron insulator. Cathode is the electrode where reduction reactions (accepting electrons from external circuit) occur, while anode is the opposite electrode to cathode, where oxidation reactions (releasing electrons to external circuit) take place. Therefore, both cathode and anode are typically electron conductors and catalytically active to the respective reduction and oxidation reactions. As such, they are often labeled as electrocatalysts by electrochemists to differentiate from conventional catalysts in chemical catalysis that do not involve electron transfer from/to external circuit. At the device level (e.g. batteries, electrolyzers, and fuel cells), current collector is also considered as an integral component of the cell. The performance (e.g. power, or chemical production rate) of an electrochemical cell is generally determined by the ohmic resistance of electrolyte and current collector, and polarization resistances (activation and concentration) of the two electrodes.

The modern electrochemical cells primarily comprise fuel cells, electrolyzers, and batteries, but are often extended to include pseudocapacitors (supercapacitors) and photoelectrochemical cells that involve electron transfer. Fuel cells are a type of concentration or galvanic cell operated in an open system for the purpose of producing electrical power. For a continuous power generation, the underlying chemical reactions must be spontaneous, which is characterized by a negative Gibbs free energy change and a positive electromotive force (EMF) or Nernst potential (En). Conversely, electrolyzer is a type of electrolytic cell operated in an open system for the purpose of producing chemicals. Since the underlying chemical reactions are non-spontaneous, which is characterized by positive Gibbs free energy change and negative En, electrical current is needed to drive the reaction. If the produced chemicals can be stored externally for later use, electrolyzers are also viewed as an energy storage device. The capacity of fuel cells and electrolyzers is scaled by the surface area (not by mass) because they are open systems.

In contrast to fuel cells and electrolyzers, batteries operate as a closed (or semi-closed) system with alternating power (En > 0) and chemical production (En < 0) modes. The chemical energy is shuttled between the two electrodes with electrons as the charge regulator through the external circuit. Therefore, the capacity of a battery is influenced by the mass of active electrodes and typically scaled by weight and/or volume.

1.3 Half-Reactions in Electrochemical Cells

The overall reaction of an electrochemical cell is represented by a chemical reaction. This chemical reaction is made up of two independent half-electrode reactions that describe the real chemical changes at the two electrodes. Each half-reaction describes electro-active species involved in electron transfer at the corresponding electrode and defines a fixed potential E (E0 under standard condition); refer to Appendix B for these values. The open circuit voltage OCV of a full cell is, therefore, the difference of the electrode potentials of the two half-electrode reactions,

For the half-reaction of interest, the electrode at which it occurs is called the working electrode (WE). To accurately characterize its behaviors, a three-electrode cell configuration is needed, where a reference electrode (RE) and counter electrode (CE) are included, see Figure 1.3.

Example 1.1

Calculate the OCV of the cell: (−) Cu(s)|Cu2+(aq) ||salt bridge|| Ag+(aq)|Ag(s) (+).

Solution

From Appendix B, E0(Cu/Cu2+) = 0.520 V, E0(Ag/Ag+) = 0.7996 V, OCV = 0.7996 − 0.520 = 0.2796 V.

The functionality of RE is to provide a constant potential to which the potential of WE is referred. Therefore, RE is an electrode made up of phases of a constant composition, which enables it, by Gibbs’s phase rule, to exhibit a fixed potential at a given temperature and pressure, even under the passage of small currents. Thus, they are also termed “nonpolarizable electrode.” The primary reference, chosen by convention, is the normal hydrogen electrode (NHE), aka. the standard hydrogen electrode (SHE), Pt|H2 (a = 1)|H+ (a = 1), represented by the half-reaction of

Its thermodynamic potential has been assigned zero at all temperatures. However, SHE is not convenient to use. In practice, the saturated calomel electrode (SCE), Hg/Hg2Cl2/KCl (saturated in water), is commonly used as a RE, which has a standard potential of 0.242 V versus SHE and the half-reaction of