How Batteries Actually Work

In March of 1800, on a workbench at the University of Pavia, Alessandro Volta did something deceptively simple. He stacked small disks of zinc and silver in alternating layers, slipped a piece of cardboard soaked in salty brine between each pair, and ran a wire from the top of the column to the bottom. When he closed the circuit, a current flowed, not a single static spark like the ones that had amused natural philosophers for decades, but a sustained, continuous flow of electricity that did not fade after an instant. On the twentieth of that month he reported the result to the Royal Society of London, and the world quietly acquired its first reliable source of electric current.

That column, which we now call the voltaic pile, ended a bitter argument and started an entire science. Two centuries later the same fundamental chemistry, refined by a long line of names like Daniell, Planté, Grove, Whittingham, Goodenough, and Yoshino, runs the phone in your pocket, the car in your driveway, and the warehouse-sized banks that store power for the electrical grid. The remarkable thing is that almost nobody, holding a AA cell or plugging in a laptop, knows what is actually happening inside. The honest answer overturns the most common thing people believe about batteries, so let us build it up carefully, from Volta's brine-soaked cardboard to the lithium-ion cell.

The Argument Volta Settled With a Stack of Metal

Before the pile, the great electrical question of the late eighteenth century concerned frogs. Luigi Galvani, an anatomist in Bologna, had noticed that a dead frog's leg twitched when touched by two different metals, and he concluded that the muscle itself held a reservoir of what he called animal electricity, a vital fluid stored in living tissue. It was an attractive idea, and it had the weight of careful experiment behind it.

Volta disagreed. He suspected the frog was merely a sensitive detector, and that the electricity arose not in the animal but at the junction where two dissimilar metals met a moist conductor. His pile was the decisive proof. By stacking zinc and silver separated by brine and producing a steady current with no biological tissue anywhere in sight, he demonstrated that the effect lived in the metal-metal-electrolyte contact, not in the muscle. The dispute was over. The current came from chemistry. In recognition, the unit of electrical potential, the volt, carries his name, and the whole field of converting chemical reactions into electric current took its first confident step.

Why a Battery Stores Chemistry, Not Electricity

Here is the misconception worth correcting at the outset, because it shapes how everything else makes sense. A battery does not store electricity. There is no little reservoir of charge sitting inside a AA cell waiting to be poured out. What a battery stores is chemical energy, locked in the arrangement of its reactants, and it converts that energy into electrical work only on demand, only when you complete a circuit and let the reaction proceed.

While we are clearing the ground, a second point of language deserves attention. Strictly speaking, a battery is not a single unit at all. The word refers to two or more electrochemical cells connected in series or parallel, just as a battery of cannon is several guns acting together. A familiar 9-volt transistor battery, for instance, is really six 1.5-volt cells stacked inside one rectangular case. The single AA you drop into a remote control is, properly, a cell. The distinction is not pedantry; it explains why voltages come in the particular round numbers they do, and it points to the unit that actually does the work, which is the cell.

Inside a Cell: Anode, Cathode, and the Salt Bridge

Strip a cell down to its essentials and you find a strikingly tidy structure. A galvanic cell, the kind that produces current from a spontaneous reaction, is built from two half-cells, each a metal electrode sitting in a solution of its own ions, joined by two connections. One is the external wire, the path through which electrons travel and the only place the electricity does anything useful. The other is an internal bridge that keeps the whole thing electrically balanced.

The chemistry divides cleanly between the two electrodes. At one, called the anode, the metal gives up electrons in a reaction we call oxidation. Those liberated electrons cannot cross the solution directly, so they push out through the wire, doing electrical work along the way, and arrive at the other electrode, the cathode, where they are accepted in a reaction called reduction. Generations of students have kept the assignments straight with the mnemonic AN-OX RED-CAT: oxidation at the anode, reduction at the cathode. It holds for any galvanic cell.

There is a subtlety that the external wire alone cannot resolve. As the anode sheds positive ions into its solution and the cathode pulls them out of its own, charge would rapidly build up on each side and stall the reaction within moments. The fix is the salt bridge, an internal path that lets ions migrate between the two half-cells and so closes the ionic circuit. Electrons flow through the wire on the outside; ions flow through the bridge on the inside; and because charge never accumulates at either electrode, the current can continue. Remove the bridge and the current stops, which is a vivid demonstration that a cell needs both a complete electronic circuit and a complete ionic one.

The Daniell Cell and the Textbook Voltage of 1.10

The cleanest illustration of all this came in 1836, when John Frederic Daniell at King's College London assembled what remains the standard teaching example. On one side he placed a zinc electrode in a solution of zinc sulfate; on the other, a copper electrode in copper sulfate; and between them a porous ceramic barrier that did the work of a salt bridge, letting ions pass while keeping the two solutions from mixing freely.

The reactions are exactly the textbook pair. At the zinc anode, each zinc atom gives up two electrons and dissolves into solution as a zinc ion, written as Zn becoming Zn(II) plus two electrons. Those electrons travel through the wire to the copper cathode, where copper ions already in solution accept them and plate out as solid metal, Cu(II) plus two electrons becoming Cu. The net effect is that the zinc electrode slowly dissolves while the copper electrode grows, and the difference in the two metals' eagerness to hold electrons appears across the terminals as a standard potential of 1.10 volts. That number, born of nothing more exotic than zinc, copper, and their sulfate solutions, is the value every chemistry student calculates first, because it makes the abstract idea of electrode potential concrete and measurable.

One Use or Many: Primary, Secondary, and the Fuel Cell

Once you can build a cell, the practical question becomes whether you can run it backward. This sorts the entire field into two great classes. A primary cell runs its chemistry once and is then finished; the reaction is not practically reversible, so when the reactants are spent the cell is discarded. The ordinary alkaline AA, built on zinc and manganese dioxide in a potassium hydroxide electrolyte, is the canonical example, a design that traces to the work of Lewis Urry around 1959.

A secondary cell, by contrast, can be driven in reverse by an external power source, pushing the products back into reactants and recharging the cell for another cycle. Gaston Planté built the first one in 1859, the lead-acid cell, whose lead plates in sulfuric acid still deliver the heavy burst of current that cranks the starter motor in nearly every internal-combustion car on the road. Lead-acid is more than a century and a half old and remains in service precisely because it is cheap, robust, and good at the one job it is asked to do. The dominant rechargeable chemistry today, though, is lithium-ion, which we will come to shortly.

There is also a third architecture that resembles a cell but escapes the primary-secondary dichotomy entirely. In a fuel cell, the reactants are not stored inside the device but fed in continuously from outside, so the cell produces current for as long as fuel keeps arriving. William Robert Grove, a Welsh judge and amateur chemist, demonstrated the principle in 1839 with his gas voltaic battery: hydrogen bubbling over one platinum electrode and oxygen over the other, both immersed in dilute sulfuric acid, combining to produce water and an electric current. The fuel cell is best thought of as the continuous-feed cousin of the battery, sharing the same electrochemistry but never running down so long as it is fed. NASA put the idea to spectacular use, flying alkaline hydrogen-oxygen fuel cells on every crewed Apollo mission from Apollo 7 in 1968, where the cells supplied both electricity and, conveniently, drinking water for the astronauts.

The Lithium-Ion Arc and the Numbers That Govern It All

The battery that defines modern life was not invented in a single flash but assembled across three laboratories over roughly twelve years. M. Stanley Whittingham, working at Exxon in the early 1970s, demonstrated intercalation, the trick of sliding lithium ions reversibly in and out of the layered structure of a host material, titanium disulfide, without tearing it apart. John B. Goodenough, at Oxford in 1980, found a host that raised the working voltage to roughly 4 volts, lithium cobalt oxide, which became the standard cathode material. Akira Yoshino, at Asahi Kasei in 1985, completed the design by pairing that cobalt-oxide cathode with an anode of petroleum coke, later refined to graphite, producing a cell that was safe and practical to manufacture. Sony commercialized the first such cell in 1991, and in 2019 the Nobel Prize in Chemistry recognized Whittingham, Goodenough, and Yoshino together for the achievement.

Underneath every one of these cells, from Volta's pile to a lithium-ion pack, lies a quantitative spine laid down by Michael Faraday at the Royal Institution in 1834. His two laws of electrolysis state, first, that the mass of substance deposited or liberated at an electrode is proportional to the quantity of electric charge passed through, and second, that for the same amount of charge the masses are proportional to the substances' equivalent weights. From these laws comes the constant that converts between the electrical world of coulombs and the chemical world of moles, the faraday, equal to 96,485 coulombs per mole, which is simply the total charge carried by one mole of electrons. Every honest battery calculation, whether you are estimating how long a cell will last or how much metal it will consume, passes through that number.

Key Takeaways

A battery is best understood not as a tank of stored electricity but as stored chemical energy converted to electrical work on demand, and strictly speaking it is two or more electrochemical cells joined together, which is why a 9-volt battery is really six 1.5-volt cells. Each cell runs on a single tidy scheme proven by Volta's 1800 pile and codified in Daniell's 1836 zinc-copper cell at 1.10 volts: oxidation at the anode and reduction at the cathode (AN-OX RED-CAT), with electrons flowing through the external wire while a salt bridge carries ions internally so charge never piles up. Cells sort by reusability into primary single-use chemistries like the alkaline cell, secondary rechargeable ones beginning with Planté's 1859 lead-acid and now dominated by the lithium-ion design that Whittingham, Goodenough, and Yoshino built between the early 1970s and 1985 (Nobel Prize 2019), and the continuous-feed fuel cell that Grove demonstrated in 1839 and NASA flew on Apollo. Binding the whole subject together quantitatively are Faraday's two 1834 laws of electrolysis and the faraday constant, 96,485 coulombs per mole, the charge of one mole of electrons and the bridge between coulombs and moles in every electrochemistry problem.