Image source: Andy Simmons / Flickr.

How do batteries power our phones, computers and other devices?

How a battery works

Expert reviewers

Dr Anand Bhatt

Research Team Leader, Advanced Energy Storage Technologies


Professor Maria Forsyth FAA

Chair, Electromaterials and Corrosion Sciences

Deakin University

Professor Ray Withers FAA

Research School of Chemistry

The Australian National University

Professor Guoxiu Wang

Director, Centre for Clean Energy Technology

University of Technology Sydney


  • A battery is a device that stores chemical energy and converts it to electrical energy.
  • The chemical reactions in a battery involve the flow of electrons from one material (electrode) to another, through an external circuit. 
  • The flow of electrons provides an electric current that can be used to do work.
  • To balance the flow of electrons, charged ions also flow through an electrolyte solution that is in contact with both electrodes.
  • Different electrodes and electrolytes produce different chemical reactions that affect how the battery works, how much energy it can store and its voltage. 

Imagine a world without batteries. All those portable devices we’re so dependent on would be so limited! We’d only be able to take our laptops and phones as far as the reach of their cables, making that new running app you just downloaded onto your phone fairly useless.

Luckily, we do have batteries. Back in 150 BC in Mesopotamia, the Parthian culture used a device known as the Baghdad battery, made of copper and iron electrodes with vinegar or citric acid. Archaeologists believe these were not actually batteries but were used primarily for religious ceremonies.

The invention of the battery as we know it is credited to the Italian scientist Alessandro Volta, who put together the first battery to prove a point to another Italian scientist, Luigi Galvani. In 1780, Galvani had shown that the legs of frogs hanging on iron or brass hooks would twitch when touched with a probe of some other type of metal. He believed that this was caused by electricity from within the frogs’ tissues, and called it ‘animal electricity’.

Luigi Galvani found that the legs of frogs suspended on brass hooks would twitch when prodded with a probe made of another type of metal. He thought this response was caused by ‘animal electricity’ from within the frog. Image source: Luigi Galvani / Wikimedia Commons.

Volta, while initially impressed with Galvani’s findings, came to believe that the electric current came from the two different types of metal (the hooks on which the frogs were hanging and the different metal of the probe) and was merely being transmitted through, not from, the frogs’ tissues. He experimented with stacks of layers of silver and zinc interspersed with layers of cloth or paper soaked in saltwater, and found that an electric current did in fact flow through a wire applied to both ends of the pile.

Alessandro Volta’s battery: a pile of zinc and silver sheets interspersed with cloth or paper soaked in saltwater. Imagine using that to power your phone. Image source: Luigi Chiesa / Wikimedia Commons.

Volta also found that by using different metals in the pile, the amount of voltage could be increased. He described his findings in a letter to Joseph Banks, then president of the Royal Society of London, in 1800. It was a pretty big deal (Napoleon was fairly impressed!) and his invention earned him sustained recognition in the honour of the ‘volt’ (a measure of electric potential) being named after him.

I myself, joking aside, am amazed how my old and new discoveries of ... pure and simple electricity caused by the contact of metals, could have produced so much excitement.Alessandro Volta

So what exactly was happening with those layers of zinc and silver, and indeed, the twitching frogs’ legs?

The chemistry of a battery

A battery is a device that stores chemical energy, and converts it to electricity. This is known as electrochemistry and the system that underpins a battery is called an electrochemical cell. A battery can be made up of one or several (like in Volta's original pile) electrochemical cells. Each electrochemical cell consists of two electrodes separated by an electrolyte.

So where does an electrochemical cell get its electricity from? To answer this question, we need to know what electricity is. Most simply, electricity is a type of energy produced by the flow of electrons. In an electrochemical cell, electrons are produced by a chemical reaction that happens at one electrode (more about electrodes below!) and then they flow over to the other electrode where they are used up. To understand this properly, we need to have a closer look at the cell's components, and how they are put together. 


To produce a flow of electrons, you need to have somewhere for the electrons to flow from, and somewhere for the electrons to flow to. These are the cell’s electrodes. The electrons flow from one electrode called the anode (or negative electrode) to another electrode called the cathode (the positive electrode). These are generally different types of metals or other chemical compounds.

In Volta’s pile, the anode was the zinc, from which electrons flowed through the wire (when connected) to the silver, which was the battery’s cathode. He stacked lots of these cells together to make the total pile and crank up the voltage. 

Illustration of a voltaic pile, explained above

But where does the anode get all these electrons from in the first place? And why are they so happy to be sent off on their merry way over to the cathode? It all comes down to the chemistry that’s going on inside the cell. 

There are a couple of chemical reactions going on that we need to understand. At the anode, the electrode reacts with the electrolyte in a reaction that produces electrons. These electrons accumulate at the anode. Meanwhile, at the cathode, another chemical reaction occurs simultaneously that enables that electrode to accept electrons. 

The technical chemical term for a reaction that involves the exchange of electrons is a reduction-oxidation reaction, more commonly called a redox reaction. The entire reaction can be split into two half-reactions, and in the case of an electrochemical cell, one half-reaction occurs at the anode, the other at the cathode. Reduction is the gain of electrons, and is what occurs at the cathode; we say that the cathode is reduced during the reaction. Oxidation is the loss of electrons, so we say that the anode is oxidised.

Each of these reactions has a particular standard potential. Think of this characteristic as the reaction’s ability/efficiency to either produce or suck up electrons—its strength in an electron tug-of-war. 

  • Standard potentials for half-reactions

    Below is a list of half reactions that involve the release of electrons from either a pure element or chemical compound. Listed next to the reaction is a number (E0) that compares the strength of the reaction’s electrochemical potential to that of hydrogen’s willingness to part with its electron (if you look down the list, you will see that the hydrogen half-reaction has an E0 of zero). E0 is measured in volts.

    The reason this list is so interesting is that if you pick two reactions from the list, and combine them to make an electrochemical cell, the E0 values tell you which way the overall reaction will proceed: the reaction with the more negative E0 value will donate its electrons to the other reaction and this determines your cell’s anode and cathode. The difference between the two E0 values tells you your cells electrochemical potential, which is basically the voltage of the cell.

    So, if you take lithium and fluoride, and manage to combine them to make a battery cell, you will have the highest voltage theoretically attainable for an electrochemical cell. This list also explains why in Volta's pile, the zinc was the anode, and silver the cathode: the zinc half-reaction has a lower (more negative) E0 value (-0.7618) than the silver half-reaction (0.7996).

    Standard potentials for reduction half-reactions

    (with respect to the standard hydrogen electrode at 25°C) 

    E° (V)
    Li+(aq) + e is equivalent to Li(s) –3.040
    Be2+(aq) + 2e is equivalent to Be(s) –1.99
    Al3+(aq) + 3e is equivalent to Al(s) –1.676
    Zn2+(aq) + 2e is equivalent to Zn(s) –0.7618
    Ag2S(s) + 2e is equivalent to 2Ag(s) + S2−(aq) –0.71
    Fe2+(aq) + 2e is equivalent to Fe(s) –0.44
    Cr3+(aq) + e is equivalent to Cr2+(aq) –0.424
    Cd2+(aq) + 2e is equivalent to Cd(s) –0.4030
    PbSO4(s) + 2e is equivalent to Pb(s) + SO42−(aq) –0.356
    Ni2+(aq) + 2e is equivalent to Ni(s) –0.257
    2SO42−(aq) + 4H+(aq) + 2e is equivalent to S2O62−(aq) + 2H2O(l) –0.25
    Sn2+(aq) + 2e is equivalent to Sn(s) −0.14
    2H+(aq) + 2e is equivalent to H2(g) 0
    Sn4+(aq) + 2e is equivalent to Sn2+(aq) 0.154
    Cu2+(aq) + e is equivalent to Cu+(aq) 0.159
    AgCl(s) + e is equivalent to Ag(s) + Cl(aq) 0.2223
    Cu2+(aq) + 2e is equivalent to Cu(s) 0.3419
    O2(g) + 2H2O(l) + 4e is equivalent to 4OH(aq) 0.401
    H2SO3(aq) + 4H+(aq) + 4e is equivalent to S(s) + 3H2O(l) 0.45
    I2(s) + 2e is equivalent to 2I(aq) 0.5355
    MnO42−(aq) + 2H2O(l) + 2e is equivalent to MnO2(s) + 4OH(aq) 0.6
    O2(g) + 2H+(aq) + 2e is equivalent to H2O2(aq) 0.695
    H2SeO3(aq) + 4H+ + 4e is equivalent to Se(s) + 3H2O(l) 0.74
    Fe3+(aq) + e is equivalent to Fe2+(aq) 0.771
    Ag+(aq) + e is equivalent to Ag(s) 0.7996
    NO3(aq) + 3H+(aq) + 2e is equivalent to HNO2(aq) + H2O(l) 0.94
    Br2(aq) + 2e is equivalent to 2Br(aq) 1.087
    MnO2(s) + 4H+(aq) + 2e is equivalent to Mn2+(aq) + 2H2O(l) 1.23
    O2(g) + 4H+(aq) + 4e is equivalent to 2H2O(l) 1.229
    Cr2O72−(aq) + 14H+(aq) + 6e is equivalent to 2Cr3+(aq) + 7H2O(l) 1.36
    Cl2(g) + 2e is equivalent to 2Cl(aq) 1.396
    Ce4+(aq)+eis equivalent toCe3+(aq) 1.44
    PbO2(s) + HSO4(aq) + 3H+(aq) + 2e is equivalent to PbSO4(s) + 2H2O(l) 1.69
    H2O2(aq) + 2H+(aq) + 2e is equivalent to 2H2O(l) 1.763
    F2(g) + 2eis equivalent to 2F(aq) 2.87

    Source: UC Davis ChemWiki

Any two conducting materials that have reactions with different standard potentials can form an electrochemical cell, because the stronger one will be able to take electrons from the weaker one. But the ideal choice for an anode would be a material that produces a reaction with a significantly lower (more negative) standard potential than the material you choose for your cathode. What we end up with is electrons being attracted to the cathode from the anode (and the anode not trying to fight very much), and when provided with an easy pathway to get there—a conducting wire—we can harness their energy to provide electrical power to our torch, phone, or whatever. 

Illustration overview of an electrochemical cell

The difference in standard potential between the electrodes kind of equates to the force with which electrons will travel between the two electrodes. This is known as the cell’s overall electrochemical potential, and it determines the cell’s voltage. The greater the difference, the greater the electrochemical potential, and the higher the voltage. 

To increase a battery’s voltage, we’ve got two options. We could choose different materials for our electrodes, ones that will give the cell a greater electrochemical potential. Or, we can stack several cells together. When the cells are combined in a particular way (in series), it has an additive effect on the battery’s voltage. Essentially, the force at which the electrons move through the battery can be seen as the total force as it moves from the anode of the first cell all the way through however many cells the battery contains to the cathode of the final cell. 

When cells are combined in another way (in parallel) it increases the battery’s possible current, which can be thought of as the total number of electrons flowing through the cells, but not its voltage.


But the electrodes are just part of the battery. Remember Volta’s bits of paper soaked in salty water? The salty water was the electrolyte, another crucial part of the picture. An electrolyte can be a liquid, gel or a solid substance, but it must be able to allow the movement of charged ions. 

Electrons have a negative charge, and as we’re sending the flow of negative electrons around through our circuit, we need a way to balance that charge movement. The electrolyte provides a medium through which charge-balancing positive ions can flow.

As the chemical reaction at the anode produces electrons, to maintain a neutral charge balance on the electrode, a matching amount of positively charged ions are also produced. These don’t go down the external wire (that’s for electrons only!) but are released into the electrolyte.

At the same time, the cathode must also balance the negative charge of the electrons it receives, so the reaction that occurs here must pull in positively charged ions from the electrolyte (alternatively, it may also release negative charged ions from the electrode into the electrolyte). 

So, while the external wire provides the pathway for the flow of negatively charged electrons, the electrolyte provides the pathway for the transfer of positively charged ions to balance the negative flow. This flow of positively charged ions is just as important as the electrons that provide the electric current in the external circuit we use to power our devices. The charge balancing role they perform is necessary to keep the entire reaction running. 

Now, if all the ions released into the electrolyte were allowed to move completely freely through the electrolyte, they would end up coating the surfaces of the electrodes and clog the whole system up. So the cell generally has some sort of barrier to prevent this from happening.

Start animation

When the battery is being used, we have a situation where there is a continuous flow of electrons (through the external circuit) and positively charged ions (through the electrolyte). If this continuous flow is halted—if the circuit is open, like when your torch is turned off—the flow of electrons is halted. The charges will pile/build up and the chemical reactions driving the battery will stop. 

As the battery is used, and the reactions at both electrodes chug along, new chemical products are made. These reaction products can create a kind of resistance that can prevent the reaction from continuing with the same efficiency. When this resistance becomes too great, the reaction slows down. The electron tug-of-war between the cathode and anode also loses its strength and the electrons stop flowing. The battery slowly goes flat. 

Recharging a battery

Some common batteries are single use only (known as primary or disposable batteries). The trip the electrons take from the anode over to the cathode is one-way. Either their electrodes become depleted as they release their positive or negative ions into the electrolyte, or the build-up of reaction products on the electrodes prevents the reaction from continuing, and it’s done and dusted. The battery ends up in the bin (or hopefully the recycling, but that’s a whole other Nova topic). 

But. The nifty thing about that flow of ions and electrons as it takes place in some types of batteries that have appropriate electrode materials, is that it can also go backwards, taking our battery back to its starting point and giving it a whole new lease on life. Just as batteries transformed the way we’ve been able to use various electrical devices, rechargeable batteries have further transformed those devices’ utility and lifespans.

When we connect an almost flat battery to an external electricity source, and send energy back in to the battery, it reverses the chemical reaction that occurred during discharge. This sends the positive ions released from the anode into the electrolyte back to the anode, and the electrons that the cathode took in also back to the anode. The return of both the positive ions and electrons back into the anode primes the system so its ready to run again: your battery is recharged.

Start animation

The process isn’t perfect, however. The replacement of the negative and positive ions from the electrolyte back on to the relevant electrode as the battery is recharged isn’t as neat or as nicely structured as the electrode was in the first place. Each charge cycle degrades the electrodes just a little bit more, meaning the battery loses performance over time, which is why even rechargeable batteries don’t keep on working forever. 

Over the course of several charge and discharge cycles, the shape of the battery's crystals becomes less ordered. This is exacerbated when a battery is discharged/recharged at a high rate—for example, if you drive your electric car in big bursts of speed rather than steadily. High-rate cycling leads to the crystal structure becoming more disordered, with a less efficient battery as a result.


Memory effect and self discharge

The almost-but-not-quite completely reversible discharge and recharge reactions also contribute to something called the ‘memory effect’. When you recharge some types of rechargeable batteries without sufficiently discharging them first, they ‘remember’ where they were up to in earlier discharge cycles and don’t recharge properly. 

In some cells, it is caused by the way the metal and the electrolyte react to form a salt (and the way that salt then dissolves again and metal is replaced on the electrodes when you recharge it). We want our cells to have nice, uniform, small crystals of salt coating a perfect metal surface, but that’s not what we get in the real world! The way some crystals form is very complex, and the way some metals deposit during recharge is also surprisingly complex, which is why some battery types have a bigger memory effect than others. The imperfections mainly depend on the charge state of the battery to start with, the temperature, charge voltage and charging current. Over time, the imperfections in one charge cycle can cause the same in the next charge cycle, and so on, and our battery picks up some bad memories. The memory effect is strong for some types of cells, such as nickel-based batteries. Other types, like lithium-ion, don’t suffer from this problem.

Another aspect of rechargeable batteries is that the chemistry that makes them rechargeable also means they have a higher tendency towards self-discharge. This is when internal reactions occur within the battery cell even when the electrodes are not connected via the external circuit. This results in the cell losing some of its chemical energy over time. A high self-discharge rate seriously limits the life of the battery—and makes them die during storage. 

The lithium-ion batteries in our mobile phones have a pretty good self-discharge rate of around 2–3 per cent per month, and our lead-acid car batteries are also pretty reasonable—they tend to lose 4–6 per cent per month. Nickel-based batteries lose around 10–15 per cent of their charge per month, which is not very good if you plan to store a torch for a whole season when you don’t need it! A non-rechargeable alkaline battery only loses around 2–3 per cent of its charge per year.  

Voltage, current, power, capacity … what’s the difference?

All these words basically describe the strength of a battery, right? Well, sort of. But they’re all subtly different. 

Voltage = force at which the reaction driving the battery pushes electrons through the cell. This is also known as electrical potential, and depends on the difference in potential between the reactions that occur at each of the electrodes, that is, how strongly the cathode will pull the electrons (through the circuit) from the anode. The higher the voltage, the more work the same number of electrons can do. 

Current = the number of electrons that happen to be passing through any one point of a circuit at a given time. The higher the current, the more work it can do at the same voltage. Within the cell, you can also think of current as the number of ions moving through the electrolyte, times the charge of those ions.

Power = voltage x current. The higher the power, the quicker the rate at which a battery can do work—this relationship shows how voltage and current are both important for working out what a battery is suitable for.

Capacity = the power of the battery as a function of time, which is used to describe the length of time a battery will be able to power a device for. A high-capacity battery will be able to keep going for a longer period before going flat/running out of current. Some batteries have a sad little quirk—if you try and draw too much from them too quickly, the chemical reactions involved can’t keep up and the capacity is less! So, we always have to be careful when we talk about battery capacity and remember what the battery is going to be used for.

Another popular term is ‘energy density’. This is the amount of energy a device can hold per unit volume, in other words, how much bang you get for your buck in terms of power vs. size. With a battery, generally the higher the energy density the better, as it means the battery can be smaller and more compact, which is always a plus when you need it to power something you want to keep in your pocket. It’s even a plus for electric cars—the battery has to fit in the car somehow!

For some applications, such as storing electricity at a renewable power plant like a wind or solar farm, a high energy density isn’t so much of a problem, as they will most likely have ample space to store the batteries. The main goal for this use would be to simply store as much electricity as possible, as safely and cheaply as possible. 

Video: How do batteries work? (TED-Ed / YouTube). View details and transcript.

Why so many types?

A range of materials (it used to be just metals) can be used as the electrodes in a battery. Over the years, many, many different combinations have been tried out, but there are only a few that have really gone the distance. But why use different combinations of metals anyway? If you’ve got a pair of metals that work well together as electrodes, why bother messing around with others?

Different materials have different electrochemical properties, and so they produce different results when you put them together in a battery cell. For example, some combinations will produce a high voltage, very quickly, but then drop off rapidly, unable to sustain that voltage for long. This is good if you need to produce, say, a sudden flash of light like a camera flash. 

Other combinations will only produce a trickle of current, but they’ll keep that trickle going for ages. We don’t need a huge amount of current to power a smoke detector, for example, but we do want our smoke detectors to keep going for a long time. 

Another reason to use different combinations of metals is that often two or more battery cells need to be stacked to obtain the required voltage, and it turns out that some electrode combinations stack together much more happily than other combinations. For example, the lithium iron phosphate batteries (a type of lithium-ion battery) used in electric cars stack together to make high voltage systems (100 or even more volts), but you’d never do that with those NiCad Walkman batteries that get hot!

Our different needs over time have led to the development of a huge array of battery types. To read more about them, and what the future holds for battery power, check out our other Nova topics.