Sounds good, Rahul! But I’ve never heard of a lemon or orange battery before. What’s even in it?

Sure! The fruity battery I made at home (pictured above) has the following components.

• Zinc metal

• Copper metal

• Orange (though any fruit would have worked!)

• Multimeter with alligator clips (for measurement purposes)

When set up in the way I’ve pictured, these turn into a single battery cell with measurable voltage. In fact, this could power some small-scale devices (e.g. LEDs, hearing aids, etc.).

Wait, you’re telling me that all those together are an actual battery? My batteries at home look nothing like a metallic fruit though! Why is this relevant to modern-day batteries?

Yes — these actually do constitute a functioning battery! Besides being a common elementary school science fair entry, the simplicity of this setup is very similar to one of the first batteries ever created: the voltaic pile, created by and named for Italian inventor Alessandro Volta in 1800. In fact, Volta demonstrated this exact battery to Napoleon in 1801: as a result, the newly-in-power French leader made Volta a Count following the demonstration (see below photo and note the stacked pile on the table).

Volta presents the voltaic pile to Napoleon in Paris

(source: Wikimedia Commons)

Volta’s original battery consisted of zinc and copper, just like our science fair orange battery. Instead of bringing an orange to Napoleon, however, Volta used pieces of paper soaked in saltwater and brine to separate the metals (which are vertically placed in between alternating discs of each metal in the image above). Volta’s soaked papers and our orange serve the same purpose: they consist of electrolytes in a medium in which both metals exist. This concept of an electrolyte is present in effectively every battery chemistry devised from Volta’s days to the present. They serve to connect the two metals, otherwise known as electrodes, in a battery such that ions (positively charged atomic particles) can float through and interact with both electrodes while avoiding direct contact between the metals themselves (which would cause a short circuit if so). The word electrode is used to denote the fact that electricity can flow through those metals as a way to harness the electricity generated in the battery.

OK, but how does this orange with metal stuck in it generate electricity?

Great question! If you refer back to the image at the beginning of the blog post (of the setup I made at home), setting up this orange battery cell generates an open circuit voltage of 0.91 V (“V” for the unit of Volts, named after Volta). The open circuit voltage of a battery cell refers to how much energy can be generated from a given setup when no electrical load is present. In this case, an electrical load refers to something that draws a given amount of current, such as a lightbulb or any object that requires electricity (the multimeter is itself battery powered and does not require current from the fruit battery). But where does this come from?

Let’s work backwards: the energy you’re able to harness from a battery comes from the voltage it can produce, which we discussed above: you may have noticed that your batteries at home (such as those in the picture above from Wikimedia Commons) have voltages written on them, usually between 1.5-3.0 V for common household applications. In this case, the voltage is equivalent to the difference in energy states between the two electrodes, made of zinc and copper. Specifically, it refers to the difference in ability for each metal to give or accept electrons (negatively charged particles), the flow of which is defined as electricity.

As a universal law of physics and chemistry, things tend to lose energy. This means that a system will always try to get to the lowest energy state possible. There are entire college courses on this specific topic, but for the purposes of our fruit battery, you need to know that because zinc and copper are at different energy states, chemical reactions will take place in order to lower the overall energy of the system, and these reactions involve the flow of electrons i.e. electricity.

You can think about this electric potential difference as equivalent to a ball rolling down a hill, as seen in the image I generated above on Canva. If a ball is at a higher elevation, then it has potential energy compared to if it were at the bottom of the hill. When it spontaneously rolls down the hill due to gravity, that potential energy gets converted to kinetic energy (i.e. motion). In the same way, in a battery, the electric potential difference between two metals (e.g. zinc and copper) is what allows that chemical energy to get translated to useful electricity. Metals lose electrons and become positive ions at the battery’s anode, which is the zinc electrode in this case. Similarly, electrons are gained at the battery’s cathode, which is the copper electrode. You can imagine zinc at the top of the hill and copper at the bottom — the flow of electrons down this hill is electricity that gets picked up by our multimeter.

If electron flow is electricity and it’s driven by the metals, where do they (the electrons) come from and where do they go? What’s actually happening in this fruit battery?

Excellent question! If you run this experiment, you’ll actually see something subtle occurring near your copper electrode: you’ll hear some fizziness and maybe even see some bubbles. If you were to put a lighter near the copper electrode, you’d actually get a small burst of flames. To cut to the chase: there’s actually hydrogen gas being produced near the copper! But what does this mean and why do we care?

Luckily for us, chemists asked this exact question a long time ago, and have not only developed a methodology for measuring the chemical potential of a given metal (against a reference), but have also tabulated these values (as seen in the table below). These tabulated values are referred to as the electromotive series, and are the bread and butter of how different metals interact with each other to generate a voltage. These values come from measuring the voltage of these individual metals when connected to a standard hydrogen electrode (SHE), which is defined as being +0.000 V.

To determine the theoretical voltage generated by using two given metals, you first identify the metal that has the higher reduction potential, which refers to the voltage tabulated above. In our zinc/copper orange battery example, this would be copper (Cu), which has a reduction potential of +0.34 V. Copper’s higher reduction potential than zinc means that it will more readily pull electrons towards it than zinc would. Given that our electrodes (i.e. the only two metals in our setup) are only copper and zinc, zinc (anode) will lose electrons that will be gained by copper (cathode).

You answered where the electrons come from and where they go, but how does that lead to hydrogen gas being produced? And what actual reactions are taking place?

Ah yes. Remember that the juice inside our orange (as well as the brine in Volta’s soaked papers) have electrolytes. In both cases (orange juice and brine), the electrolyte is actually slightly acidic. If a liquid is acidic, then it has a high concentration of atomic hydrogen ions (referred to as H+) swimming around. When the circuit is connected, the zinc gives up those electrons and positively charged zinc ions get dissolved into the orange juice. Most of the electrons, on the other hand, travel through the exterior circuit and through the multimeter to the copper electrode. Lone electrons cannot travel through the electrolyte from one electrode to another. While the vast majority of electrons from the zinc travel through the exterior circuit to the copper electrode, some remain situated around the zinc electrode itself.

At the copper electrode, as a result, the abundant H+ ions swimming in the electrolyte and the newly arrived electrons from the external circuit combine together to form hydrogen gas molecules (referred to as H2). Because these gas molecules are less dense than the electrolyte, they float to the top and escape into the surroundings. Some of the electrons from the zinc electrode, however, remain on the zinc itself (though far less than what gets to the copper), so a much smaller amount of hydrogen gas is produced at the zinc electrode as well, though not nearly as appreciable as what’s produced at the copper electrode.

Previous experiments involving this exact setup (orange included) have shown that the mass dissolved of the zinc electrode (measured by solid mass loss) is equivalent to the mass evolved of hydrogen gas (source: Kelter et al 1996), further corroborating the finding that the chemical reactions involving zinc are the initial phenomenon leading to hydrogen gas production.

Full chemical reaction of the zinc/copper orange battery

(source: Goodisman, Jerry, "Observations on Lemon Cells", 2001)

Only reactions with zinc have been mentioned. Why is the copper necessary if it’s not really participating in any chemical reaction?

Excellent observation! The purpose of the copper electrode is to serve as a conduit of electric current to/from the external circuit to the electrolyte, which itself interacts with the zinc electrode. Initial observations for this setup suggested that the copper is reactive given that most of hydrogen gas was specifically produced at the copper electrode, but given that the electrons for this gas production come from the zinc, the copper participation is negligible.

In fact, it is possible to use any metal with a higher reduction potential than zinc, including noble metals such as platinum (in this case: “noble” refers to not being particularly reactive). Even though Volta’s setup used alternating zinc and copper discs, his prior experiments also used alternating zinc and silver, the latter of which you’ll see on the electromotive series table with a higher reduction potential than either copper or zinc.

The copper electrode is still critical to this setup, however, because it’s what provides that driving force for the flow of electrons. The energy difference between zinc and copper is what drives the reaction — the copper serves as a “pull” for the electrons from zinc, even though the electrons end up reacting with hydrogen ions after arriving at the copper electrode. More technically, the zinc’s electrochemical potential is higher than copper’s, which is a fancy way of saying that zinc’s electrons are at the top of our metaphorical hill and ready to roll down to the copper. Once they reach the copper, copper’s electrochemical potential is higher than that if the electrons combined with the H+ in the electrolyte, and therefore H2 gas is produced. Remember: things tend to lose energy. Using more noble elements (i.e. those with a higher reduction potential) would lead to a higher voltage due to a higher difference in energy states between zinc and the metal in question.

If silver and platinum produce a higher voltage than copper, why don’t we just use those?

Money. The answer is almost always money.

According to the U.S. Geological Survey, copper is $3.81/pound while silver trades at almost $350/pound and platinum at almost $15,000/pound as of January 2024. I wish there was a fancy scientific reason, but sometimes life gets in the way of that, so the answer here is money.

The Duracell/Energizer batteries in my house still look nothing like a metallic orange. How do I make sense of Volta’s antiquated setup and modern tech?

Well first of all dear reader, it’s kind of harsh to call Volta “antiquated” given the importance of this contribution to science, but it’s a good question nonetheless. Since Volta’s time, several scientists and engineers made incremental improvements to battery technology. Volta already experimented with connecting several of these alternating disks in series to produce a higher voltage (see image above): specifically, each battery cell (denoted in the image as an “element”) contributes approximately ~1.0 V (i.e. a singular orange setup), and when connected in series the voltages are additive. If 6 elements are connected, then we get about 6 V out of the stack-up. Yes — this means that you could power a 12 V car battery with 12 oranges and some zinc and copper sticks!

Another key detail: this orange is an example of a primary battery, which means that it cannot be recharged. Once this battery produces some electricity, it cannot be recharged by plugging it in elsewhere. This makes sense because the key chemical reaction (resulting in hydrogen gas production) is not reversible with this setup as written. If you could indeed reverse the key chemical reaction in this setup by supplying external electricity to the system, then this battery would be secondary, which just means that it’s rechargeable. The batteries in your calculator, smoke detector, and other simple household devices are primary. Those in your more complex devices are secondary: think anything rechargeable, such as your smartphone, shaving razor, or car battery.

Since Volta’s time, however, many aspects of the battery have gone through extensive study and scrutiny. This includes the electrolyte itself, the composition of electrodes, physical setup to optimize performance, connections to other cells, and a host of other variables. Each class of battery setups is called a “battery chemistry” — you could very well call our setup an “orange juice battery chemistry” and that would be accurate! The pack of Duracell batteries you picked up from Costco have different chemistries than our trusty orange, but they still have an anode, cathode, and electrolyte powering it all. Other examples of battery chemistries that you may have heard of are lead-acid batteries, lithium-ion, even metal-air chemistries, and so many more.

Thanks for the info, Rahul!

No problem! The goal of this initial technical blog post was to introduce the world of battery engineering at the most basic level, both for my own understanding as well as that of any curious audience members. I will continue to keep writing up posts on various technical topics at a similar level and will engage more complex content as is required.

Whether you enjoyed or hated this post (or simply have some thoughts), I’m all ears! You can reach me at rahulr[at]alum[dot]mit[dot]edu. Thank you!