Core Concepts
In this article, we explore the electrochemistry behind batteries. We examine common examples including primary batteries, zinc-copper batteries, lead-acid batteries, nickel-cadmium batteries, and fuel cells.
Topics Covered in Other Articles
- What is Electrochemistry?
- The Reactivity Series
- The Captivating Element Copper
- The Zesty Element Zinc
- How to Balance Redox Reactions
Redox: The Chemistry behind Batteries
Though they may seem high-tech, batteries work according to fairly basic physics and chemistry. Specifically, you can explain the activity of a battery in molecular terms, as vessels for a chemical reaction that results in an electric current. On the chemical level, this current is a flow of electrons. Chemists decribe reactions that transfer electrons, including those that occur in batteries, as “redox reactions“.
In battery redox reactions, atoms of one metal element oxidize, or lose electrons, while atoms of another metal reduce, or gain electrons, which ultimately produces the desired electron current. They do this because the two metals have different reduction potentials, meaning they will spontaneously transfer electrons when in contact. In particular, the metal with the lower potential oxidizes, and is thus referred to as the anode, while the higher potential metal reduces, which chemists call the cathode. Students usually remember this terminology through the mnemonic “RED CAT AN OX” which means “REDuction at the CAThode while the ANode OXidizes”.
Let’s take a look at this terminology in action and consider the primary battery.
Primary Batteries
Commericially available batteries can be divided into two categories: primary and secondary batteries. Primary batteries are single use and disposable, which includes the Double-A and Triple-A varieties we buy for common household electronics, like flashlights and remote controls.
As we can see in the above diagram, electrons flow from the high-potential anode to the low-potential cathode through a wire that connects the two metals. Most often, the cathode contains manganese while the anode contains zinc. The electrolyte is a paste containing water, starch, ammonium chloride (NH4Cl), and insoluble manganese (IV) oxide (MnO2). The electrolyte serves to provide manganese cations for reduction, as well as ammonia and chloride to precipitate oxidized zinc cations into insoluble Zn(NH3)2Cl2. This insoluble zinc complex cannot be easily reduced back into neutral Zn, indicating that this reaction is ultimately irreversible. Thus, primary batteries cannot recharge. Despite this, primary batteries do not have a high cost to manufacture, making them a sufficient choice for common electronics.
Secondary Batteries
Secondary batteries, by contrast, are rechargeable. For an example of a secondary battery, consider the Cu/Zn galvanic cell:
The Cu/Zn galvanic cell features a copper cathode and a zinc anode. Both electrodes have a surrounding solution with cationic forms of each metal, in the form of metal sulfate. The salt bridge of the cell, usually involving a simple salt like NaCl ot KCl, allows charged electrolytes to flow into each electrode. As one solution loses and the other gains electrons, these electrolytes maintain an electrochemical balance that allows the redox chemistry to occur. As the galvanic cell runs, the solid zinc atoms will oxidize into aqueous Zn2+, which transfers electrons to the copper solution, reducing the Cu2+ cations to solid copper that deposits on the copper electrode. This activity occurs according to the following reaction equations:
Importantly, unlike primary batteries, the redox reactions involved in the activity of galvanic cells can be reversed with the application of electric current. This current temporarily converts the galvanic cell into an electrolytic cell, which oxidizes the copper back into Cu2+ and reduces Zn2+ into neutral zinc. Thus, secondary batteries can recharge. Because of this impressive ability, secondary batteries are used in a variety of valuable technologies, from mobile phones to cars to power tools. Let’s take a look at a few important examples.
Lead-Acid Batteries
One example of a common secondary battery is the lead-acid battery. Invented in 1859, the lead-acid battery is known for being inexpensive to produce while offering high output when initiated. For this reason, lead-acid batteries exist in virtually all commercially available cars for engine ignition.
Lead-acid batteries involve sulfuric acid solution with two submerged metal plates, a neutral lead anode and a lead (IV) oxide cathode (PbO2). As the battery initiates, the acidic medium helps to oxidize the neutral lead into lead (II) sulfate (PbSO4). Simultaneously, the lead oxide at the cathode also converts into PbSO4, but through reduction. The redox activity of lead-acid batteries can be summarized through the following reaction equations:
Importantly, the PbSO4 product is solid, which means that it precipitates onto the electrodes as the reaction progresses. With an added electric current, the PbSO4 simply reconverts into elemental lead and lead (IV) oxide, allowing recharge.
Nickel-Cadmium Batteries
The nickel-cadmium (Ni/Cd or NiCad) battery provides another example of a secondary battery. Lightweight and rechargeable, Ni/Cd batteries have use in a variety of electrical appliances and power tools.
Ni/Cd batteries implement a “jelly roll” structure where metallic layers of cadmium anodes and nickel cathodes are separated by a polymer with aqueous alkaline solution (the separators). These separators allow close positioning of the electrodes, which reduces internal resistance for efficient electron discharge. The nickel cathodes are actually nickel (III) oxo-hydroxide (NiO(OH)), which reduces into nickel (II) hydroxide (Ni(OH)2). Similarly, the cadmium transforms into its respective hydroxide, cadmium (II) hydroxide (Cd(OH)2), but through reduction. The redox activity of Ni/Cd batteries can be summarized through the following reaction equations:
Like the lead-acid battery, the hydroxide products become a solid precipitate on the surface of the electrodes. With application of a recharging current, the redox reactions reverse which restores the capacity of the battery.
Fuel Cell Batteries
Fuel cell batteries are a promising technology that offers a power source without fossil fuel input nor toxic heavy metals. In few words, fuel cells function like a galvanic cell that requires constant inputs of hydrogen and oxygen gas, with an output of gaseous water. Because of this use of external fuel gas, fuel cells don’t count as primary nor secondary batteries. The following diagram summarizes the redox activity of fuel cells.
In short, the central electrolyte catalyzes the paired conversion of hydrogen gas into protons (H+) and oxygen gas into oxide (O2-), producing water. This activity can be summarized in the following redox reactions:
While the technology of fuel cells seems exciting, current fuel cells remain relatively inefficient. This is due to the fact that the reduction of oxygen to oxide remains slow. Current research aims to find a suitable catalyst to speed up this reduction. One promising research lead involves the use of a platinum catalyst. Hopefully, with further study, efficient, clean fuel cells will provide a new source of energy to help in the green energy transition.