Have you ever looked inside a battery? What you probably found was a cylinder of long, wet, spiral wound strips, kind of like a spool of film. (If you’ve done this, we hope you did so in a laboratory environment under qualified supervision since the chemicals inside batteries may be very dangerous!) How does this funny configuration store energy? In this post, we will discuss how batteries, ubiqituous in modern life, actually work.
Batteries are frequently heralded as a critical component of a low-carbon energy infrastructure, but they don’t actually generate energy – they just store it. Energy storage not only allows for all types of portable electronics, but also increases the efficiency of the grid. A battery can store electricity when it is cheap (whether it be inexpensive base load nuclear power at night or a sudden burst of wind), and then release it again when demand spikes, reducing the total required installed generation capacity. Batteries can also be used in cars: as starter batteries, to recover energy from regenerative braking, or to run the car entirely off electricity. This last application eliminates tail-pipe CO2 emissions and would allow the car to be run off low-carbon sources of electricity.
Different battery chemistries are best suited to each of these applications, but they all work in the same way. A battery is an electrochemical cell that takes advantage of energy used or released by a chemical redox (reduction-oxidation) reaction. Structurally, batteries are composed of two electrodes separated by a porous separator and soaked in electrolyte. The separator allows ions in the electrolyte to travel from one side to the other and transfer charge, but it blocks the flow of electrons. Outside of the electrodes are metal current collectors, which connect the electrodes to the circuit. When a battery is charged or discharged, a complete circuit is formed with electrons flowing outside the cell and through the load and ions transferring charge within the cell.
This flow of electrons causes redox reactions to occur at each electrode. In one direction, these reactions store energy, while the reverse reactions release energy. The chemical potential of the reactions at each electrode determines the voltage of the cell. Batteries can be categorized as either primary or secondary cells. In primary cells, the chemical reactions are not easily reversible, and these batteries are usually discarded after one use. Secondary cells are rechargeable, but the repeated chemical reactions at each electrode cause them to degrade over time.
To make this more concrete, let’s look at a lithium-ion cell. A common chemistry for this type of cell is lithium cobalt oxide (LiCoO2), named for its positive electrode. The negative electrode for this cell is typically made of graphite (carbon). The electrolyte consists of a lithium salt dissolved in a solvent. The reaction at the negative electrode is:
Li+ + C6 + e- <–> LixC6
where x <1. What this formulation means is that, on charging, electrons (e-) move into the graphite electrode (C6) and combine with lithium ions (Li+) from the electrolyte, leaving lithium atoms inserted in the graphite lattice. During discharge, the electron moves the opposite direction back through the circuit and the lithium ions move to the positive electrode. The reaction there is:
LiCoO2 <–> yLi+ + Li(1-y)CoO2 + ye-
where y < 0.5. Together, the full cell reaction is:
LiCoO2 + C6 <–> Li(1-x)CoO2 + LixC6
The potential of this reaction is 3.6 V, more than twice that of your typical AA alkaline battery. This high voltage is one of the reasons that lithium-ion batteries have a higher energy density than many other chemistries which makes them a popular choice for laptops and other portable devices.
The amount of energy that can be stored in the battery is also dependent on the total number of reactants available in the electrodes. In order to make as much of the electrode material available as possible, most electrodes are porous. This porosity greatly increases the surface area of the electrode, allowing for more reaction sites. Batteries also benefit from having the electrodes as close together as possible, which increases volumetric energy density (by reducing wasted space) and improves conductivity. Higher rates of charging and discharging can be achieved with thinner electrodes because the ions don’t have to travel as far.
Together, these design elements lead to the spiral wound construction we discussed at the beginning: the current collectors, electrodes and separator are made into long strips, soaked in electrolyte, and tightly wound to reduce space. Not all batteries have this construction; small button cells or certain large lead-acid batteries have arrangements that favor different applications.
The choice of battery chemistry depends on the application. Battery researchers are constantly looking for better materials to make batteries, whether it be for cheaper inputs, higher voltages, longer cycle life, or high power capabilities. Unfortunately, few materials encompass all of these attributes. Sometimes, different types of energy storage devices must be used together to meet the needs of a given application. One such device is the ultracapacitor, which we will look at in the next installment of the Energy Storage series.