17.5 Batteries & Fuel Cells

Introduction

This section explores how batteries and fuel cells produce electricity through chemistry, and how or if they recharge.

Batteries work on the same principle as the galvanic cells we have studied. The batteries that power our world are usually multiple galvanic cells linked in series, so their electrical potential adds cumulatively.

Fuel cells work like batteries, but instead of having a sealed container of reactants and conducting medium, a constant supply of fuel flows through a fuel cell which extracts electrical energy from it. A battery adjacent to a fuel cell can store its energy. Because fuel cells do not combust their fuel, they produce energy more cleanly and more efficiently than generators.

Of course, batteries range in size from tiny buttons to behemoths for backing up municipal power grids. They also vary in composition. Some are rechargeable, while others are single use.

Single Use Batteries

The potato which kicked off this chapter was in fact a crude battery. The copper and zinc electrode transfer electrons both through the potato’s flesh which is full of solanaceous acid, and across a wire to complete a circuit. (Image credit: Eileen H. Kramer)

Within the potato the zinc anode gives up two electrons and the zinc ions, Zn+2 dissolve into the potato’s juices, while any copper that has ionized around its electrode accepts two electrons. We saw this equation earlier in the chapter:

Zn + Cu⁺² D Zn⁺² +Cu⁰

Of course we don’t power our gadgets with potatoes. The potato’s electrodes share a common medium, a weak electrolyte, which quickly gives out as a buffer. The current producing equation reaches equillibrium and stops producing power.

Alessandro Volta gets credit for creating the first device that resembles a modern, single use, battery (Barnett, 3, 2017). Volta’s, from whose name we get volts, battery was a pile of metal disks separated by cloth in a tall, glass container. The container leaked electrolytes, and soon used up its charge. (Figure 17:xx Volta’s First Battery. Image credit Steve Shupe https://www.flickr.com/photos/11155423@N00/3138858680)

The single use batteries that power our devices, are of course smaller and more long lived. A drycell is a typical single use battery. In such a battery, zinc is both the container and anode, the “-” terminal. A graphite rod is the cathode or “+ terminal.” The dry cell is not really dry inside, but instead contains an electrolyte paste of manganese(IV) oxide (Mn₂O₂), zinc(II) chloride (ZnCl₂), ammonium chloride (NH₄Cl), and some water. It’s graphite electrode is also insert.

The full chemical reaction inside a dry cell is:

Zn(s) + 2MnO₂(s) + 2NH₄Cl(aq) → Mn₂O₃(s) + H₂O(l) + Zn⁺² + 2NH₃(g)

The half reactions are…

Anode/Oxidation Zn(s) Zn⁺² + 2e^–

Cathode/Reduction: 2Mn⁺⁴O₂^-2 + 2NH₄⁺¹Cl^-1 + 2e^– Mn₂⁺³O₃^-2 + 2NH₃⁰ + H₂O⁰ + Cl^-1

Note, both the Mn⁺⁴ and the NH4⁺¹ get reduced, each taking one electron apiece.

Dry cells come in various sizes, but all share the same components, and produce the same voltage, about 1.5 V. Larger dry cells can of course transfer greater amounts of charge because they have more reactant upon which to draw. Connecting dry cells in series also increases available charge.

Figure above shows a typical dry cell. Note that both the anode and case are made of zinc, the cathode is inert, and that the bulk of the battery is filled with electrolyte paste.

The calculators, electronic games, digital watches and portable audio players that are so familiar to us are all powered by small, efficient batteries. A common primary battery is the dry cell, which uses a zinc can as both container and anode (“–” terminal) and a graphite rod as the cathode (“+” terminal). The Zn can is filled with an electrolyte paste containing manganese(IV) oxide, zinc(II) chloride, ammonium chloride, and water. A graphite rod is immersed in the electrolyte paste to complete the cell. The spontaneous cell reaction involves the oxidation of zinc:

which together yield the cell reaction:

The voltage (cell potential) of a dry cell is approximately 1.5 V. Dry cells are available in various sizes (e.g., D, C, AA, AAA). All sizes of dry cells comprise the same components, and so they exhibit the same voltage, but larger cells contain greater amounts of the redox reactants and therefore are capable of transferring correspondingly greater amounts of charge. Like other galvanic cells, dry cells may be connected in series to yield batteries with greater voltage outputs, if needed.

In its acid version, the dry cell battery contains a zinc inner case that acts as the anode and a carbon rod in contact with a moist paste of solid MnO₂, solid NH₄Cl, and carbon that acts as the cathode in the following figure. The half-reactions are complex but can be approximated as follow:

The above Figure shows a schematic diagram of a typical dry cell.

The standard automotive battery in today’s vehicles is a 12-volt battery. Each battery has six cells, each with 2.1 volts at full charge. A car battery is considered fully charged at 12.6 volts or higher. When the battery’s voltage drops, even a small amount, it makes a big difference in its performance.

The above figure is automobile lead storage battery where one of the six cells is a 12-V lead storage battery. The anode consists of a lead grid filled with spongy lead, and the cathode is a lead grid filled with lead dioxide. The cell also contains 38% (by mass) sulfuric acid.

An automobile with a dead battery can be “jump-started” by connecting its battery to the battery in a running automobile. But remember, this process may be dangerous, because the resulting flow of current causes electrolysis of water in the dead battery producing hydrogen and oxygen gases. Disconnecting the jumper cables after the disabled car stars causes an arc that can ignite the gaseous mixture. If this happens, the battery may explode, ejecting corrosive sulfuric acid. This problem can be avoided by connecting the ground jumper cable to a part of the engine remote from the battery. Any arc produced when this cable is disconnected will then be harmless.

Reference: Zumdahl DeCoste, 10^th Edition.

Reference: https://www.youtube.com/watch?v=Immuw51iuXA

Reference: https://www.youtube.com/watch?v=r6rJWqQaWxQ

Alkaline batteries 

Alkaline batteries were developed in the year 1950s  to improve on the performance of the dry cell battery, they were designed around the same redox couples. As their name suggests, these types of batteries use alkaline electrolytes, often potassium hydroxide.

The reactions are:

An alkaline battery can deliver about three to five times the energy of a zinc-carbon dry cell of similar size. Alkaline batteries are prone to leaking potassium hydroxide, so these should be removed from devices for long-term storage. While some alkaline batteries are rechargeable, most are not. Attempts to recharge an alkaline battery that is not rechargeable often leads to rupture of the battery and leakage of the potassium hydroxide electrolyte.

The alkaline dry cell battery Lasts for a longer period of time mainly because the Zinc anode corrodes less rapidly under basic conditions.

The above Figure shows Alkaline batteries that were designed as improved replacements for zinc-carbon (dry cell) batteries.

Reference: https://www.youtube.com/watch?v=BW4yIBS59HY

Other types of useful batteries include the “Silver Cell” which has a Zinc anode and a cathode that uses Ag₂O, as the oxidizing agent in a basic environment. Mercury Cells often used in calculators and have a Zinc anode and cathode involving HgO as the oxidizing agent in a basic medium.

The above Figure shows a Mercury battery of the type used in calculator.

https://www.youtube.com/watch?v=A-phEvlGjPs

Reference: https://www.youtube.com/watch?v=A-phEvlGjPs

Rechargeable (Secondary) Batteries

An especially important type of battery is the: Nickel–cadmium battery, in which the electrode reactions are:

As in the lead storage battery, the products adhere to the electrodes. Therefore, a nickle-cadmium battery can be recharged an indefinite number of times.

When properly treated, a NiCd battery can be recharged about 1000 times. Cadmium is a toxic heavy metal so NiCd batteries should never be ruptured or incinerated, and they should be disposed of in accordance with relevant toxic waste guidelines, never put in the regular trash.

Nickel-cadmium, or NiCd, batteries consist of a nickel-plated cathode, cadmium-plated anode, and a potassium hydroxide electrode. The positive and negative plates, which are prevented from shorting by the separator, are rolled together and put into the case. This is a “jelly-roll” design and allows the NiCd cell to deliver much more current than a similar-sized alkaline battery.

The above Figure shows the NiCd batteries use a “jelly-roll” design that significantly increases the amount of current the battery can deliver as compared to a similar-sized alkaline battery.

Reference: https://www.youtube.com/watch?v=Q0VSVy-_IIM

Lithium ion batteries are considered to be among the most popular rechargeable batteries and are used in many portable electronic devices. The reactions are:

The variable stoichiometry of the cell reaction leads to variation in cell voltages, but for typical conditions, x is usually no more than 0.5 and the cell voltage is approximately 3.7 V. Lithium batteries are popular because they can provide a large amount current, are lighter than comparable batteries of other types, produce a nearly constant voltage as they discharge, and only slowly lose their charge when stored.

The Figure above show in a lithium ion battery, charge flows between the electrodes as the lithium ions move between the anode and cathode.

Lithium-ion-batteries involve the migration of Li⁺ ions from the cathode to the anode, where they intercalate (enter the interior) as the battery is charged.

At the same time, Charge-balancing electrons travel to the anode through the external circuit in the charger. On discharge, the opposite process occurs.

The cathode of the first successful Lithium-ion batteries originally contained LiCoO₂ and a Lithium-intercalated carbon (LiC ) anode. More recently manufacturers have included transition metals such as nickle and manganese  in the cathode in addition to cobalt. The mixed metal cathodes have greater charge capacity and power output and shorter recharge times.

Reference: https://www.youtube.com/watch?v=zUlbHMDCosI

Fuel cells  is a Galvanic cell for which reactants are continuously supplied. To illustrate the principles of fuel cell, let’s consider the exothermic redox reaction of methane with oxygen.

Usually the energy from this reaction is released as heat to warm homes and to run machines. However, in a fuel cell designed to use this reaction, the energy is used to produce electric current. The electrons flow from the reducing agent (CH₄) to the oxidizing agent (O₂) through a conductor.

In a galvanic cell that uses traditional combustive fuels, most often hydrogen or methane, that are continuously fed into the cell along with an oxidant. (An alternative, but not very popular, name for a fuel cell is a flow battery.) Within the cell, fuel and oxidant undergo the same redox chemistry as when they are combusted, but via a catalyzed electrochemical that is significantly more efficient. For example, a typical hydrogen fuel cell uses graphite electrodes embedded with platinum-based catalysts to accelerate the two half-cell reactions:

The Figure above shows hydrogen fuel cell, oxygen from the air reacts with hydrogen, producing water and electricity.

In a hydrogen fuel cell, the reactions are:

These types of fuel cells generally produce voltages of approximately 1.2 V. Compared to an internal combustion engine, the energy efficiency of a fuel cell using the same redox reaction is typically more than double (~20%–25% for an engine versus ~50%–75% for a fuel cell). Hydrogen fuel cells are commonly used on extended space missions, and prototypes for personal vehicles have been developed, though the technology remains relatively immature.

Reference: https://www.youtube.com/watch?v=5_lDGna9MBM

Reference: https://youtu.be/9zgx-PlDEKA?t=13

SUMMERY AND KEY CONCEPTS

Galvanic cells designed specifically to function as electrical power supplies are called batteries. A variety of both single-use batteries (primary cells) and rechargeable batteries (secondary cells) are commercially available to serve a variety of applications, with important specifications including voltage, size, and lifetime. Fuel cells, sometimes called flow batteries, are devices that harness the energy of spontaneous redox reactions normally associated with combustion processes. Like batteries, fuel cells enable the reaction’s electron transfer via an external circuit, but they require continuous input of the redox reactants (fuel and oxidant) from an external reservoir. Fuel cells are typically much more efficient in converting the energy released by the reaction to useful work in comparison to internal combustion engines.