Do batteries bounce?

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Do batteries bounce if they’re dead?

BatteriesLately, there’s been a bunch of videos (such as this one) floating around on the internet, that show a ‘trick’ that supposedly allows you to find out whether batteries are dead or not, by simply letting them bounce on a surface.

The question is whether this trick really works. To answer it, I’ll first need to explain a bit about how batteries work. After that, we’ll take a look at how some researchers used the scientific method to find the answer.

Reactions with electrons

Electrons, the tiny, negatively charged particles that form a ‘cloud’ around every single atom, are essential for any chemical reaction. It’s the electrons that are responsible for bonds between atoms. When atoms bind together to form a molecule, the atoms start sharing their electrons among each other. As long as the electrons are shared, the atoms are being held together by bonds.

In some reactions, electrons aren’t just shared, but actually move to another compound altogether. These reactions require at least two types of compounds. The first gives away electrons, it’s the electron donor or reductant or reducer. The second compound receives the electrons, it’s the electron acceptor, also known as oxidant or oxidizer. This type of reaction, in which electrons move to another species of chemical, are known as reduction-oxidation reactions, or in short, redox reactions.

Rusty surface

A rusty surface (Source).

Redox reactions are actually quite common. A well-known example is rusting. Iron naturally rusts when it is in contact with air. That’s because the oxygen in the air is quite a good oxidant (as the name suggests), while iron is a decent reducer. Let’s look at the reactions:

Fe -> Fe3+ + 3 e

Iron loses three electrons and turns into the positively charged iron(III) ion.
On the other side, oxygen reacts:

O2 + 4 e -> 2O2-

An oxygen molecule takes up four negatively charged electrons and forms two oxide ions.
In a final step, two of the formed iron(III) ions and three of the oxide ions react to form Fe2O3, one of the main components of rust:

2 Fe3+ + 3 O2- -> Fe2O3

In reality, these reactions are a bit more complicated. Water plays a role, too, and hydroxide ions end up in the product. These reactions were just an example of a redox reaction.

Another important property of redox reactions is that there are never any loose electrons left over. Both sides of the redox reaction happen in a specific proportion. Let’s fix them:

4 Fe -> 4 Fe3+ + 12 e
3 O2 + 12e -> 6 O2-

By multiplying the iron reaction by 4 and the oxygen reaction by 3, oxygen uses up all the electrons iron produces. This is the ratio at which these reactions take place. Because of this property, neither of these reactions can take place by itself. Either nothing happens, or both take place at the same time. That’s why we call these reactions half-reactions.

Conveniently, the 4 iron(III) ions and the 6 oxide ions can combine to form exactly 2 Fe2O3 units.

Rusting is usually an unwanted reaction. There are ways to reverse this reaction and make useful iron metal, which is something people do want. But what I want to do is use redox to make electricity.

Using redox to create an electric current

In an electric current, charged particles, usually electrons, move in a certain direction (for example through a wire). So, the trick is to design a redox reaction in which the electrons don’t simply jump from the reducer to the oxidant, but have to go all the way through a wire before they reach the oxidant. In other words, we need a way to separate the two half-reactions.

Electrochemical cell

An electrochemical cell (Adapted from Wikipedia).

This is usually done by having each half-reaction take place at a separate electrode. The electrodes are connected to a wire, which conducts electrons through any electric device, to the other electrode. Electrodes are often pieces of solid metal, submerged in a watery solution. Each half-reaction takes place on the surface of an electrode. It’s important that the species for one half-reaction never end up at the other electrode (in that case they would react locally, without creating a current through the wire). However, electricity only flows through closed circuits, so the last part of our design is something that electrically connects the watery solutions without allowing reactants to ‘escape’. This is usually done by some kind of membrane or a tube with a conducting, salty gel between the solutions. The latter is known as a salt bridge.

The complete set-up is called an electrochemical cell, and is shown to the right. This set-up, which happens to use a copper and a zinc electrode, produces electrical power as soon as the electrical device is switched on.

The word battery originally meant ‘a set of artillery’. At some point, the word also got the meaning ‘a set of electrochemical cells’, because the cells were often used together, in order to get more power than a single cell could deliver. Later on, single electrochemical cells also started being called a battery.

The only real difference between the electrochemical cell shown in the picture above and commercial batteries, is that the latter are usually more compact and use a gel instead of a liquid, so they don’t leak as easily. They’ve simply been made more portable.

The bouncy battery theory

If you have read the previous paragraphs, you should understand the chemistry in a battery well enough to be able to follow along with my search for the answer to the question: Do dead batteries bounce, and if so, why?

It seems all the websites and videos talking about it are talking about the commonly used non-rechargable alkaline battery, so I’ll focus on that. Luckily, there are more people who asked this question, and researchers at Princeton University decided to sort it out once and for all.1

The first thing the Princeton researchers Shoham Bhadra and Daniel Steingart did was test whether dead batteries really bounce. They found that full batteries don’t bounce at all, while those at 50% charge bounce quite high. After that, it tops off.

So, that’s one question answered with a simple experiment. Dead alkaline batteries do bounce, but so do those with half their charge remaining. A bounce isn’t enough to tell you whether a battery is completely dead or not.

Now, the real question is, what causes this effect? Let’s look at the chemistry in an alkaline battery.

The redox chemistry of an alkaline battery

In the alkaline battery, the reducer is zinc. It donates electrons to the negative pole (anode), through the following half-reaction, in which zinc oxide and water is formed:

Zn + 2 OH -> ZnO + H2O + 2 e

The oxidant is manganese dioxide. When it takes electrons from the positive pole (cathode), it reacts as follows, using up water in the process (formula corrected for electron ratio):

2 MnO2 + 2 H2O + 2 e –> 2 MnOOH + 2 OH

The presence of the OH ion in these reactions makes this an alkaline battery. The reactions take place in a solution containing OH, and such a solution is by definition alkaline. This alkaline solution is a thick gel with which the rest of the battery is filled.

Alright, you now know about the chemical reactions that happen in batteries. Can you answer the question? What causes the bouncyness in a dead battery?

 

Time for some answers

That was a bit of a trick question. At this point, I had no idea myself!

The Princeton researchers did measurements on each reaction in the battery in order to find the culprit. This is part of how the scientific method works: first you consider possible explanations (hypotheses), and then you devise experiments to find out if any of the explanations are plausible.

Let’s start with their first hypothesis:

Hypothesis 1: Alkaline batteries leak hydrogen gas during operation. This is why they get bouncy.

Another reaction that could possibly take place is the creation of hydrogen gas from the water. This reaction should not normally happen, but it can occur if the battery gets too hot. That’s why you shouldn’t throw batteries into a fire. Not only do they release possibly toxic chemicals, they could also explode. Do not attempt this. The researchers suggested that this might also happen during normal operation, and a small amount of hydrogen gas could leak from the battery. This effect could change the density of the battery, and so the bouncyness changes as well.

There’s a simple test for this hypothesis: compare the mass of a full battery with that of a dead one. The researchers found out that the mass stays exactly the same. Hypothesis 1 is proven wrong.

Next, we need to formulate hypotheses for each of the chemical changes that do happen in the battery. I’ll just list them in the same order as they were tested by the Princeton researchers.

Hypothesis 2: When MnO2 reacts to MnOOH, the battery gets bouncy.

Hypothesis 3: The first reaction creates one water molecule, the second reaction uses up two. When the battery starts running out of water, it gets bouncy.

Hypothesis 4: When Zn (zinc) reacts to ZnO (zinc oxide), the battery gets bouncy.

We don’t need a hypothesis for the OH, because for every OH that’s used up, a new one is formed by the other reaction. OH can migrate freely through this type of battery. The amount of OH stays the same, so it can’t be responsible.

For Hypothesis 2, the researchers measured the state of the positive electrode at several times while the battery was being used. They found that while there’s a clear chemical change, it happens completely linearly. However, the bouncyness doesn’t increase linearly: it starts slow, until a battery has about 70-60% charge remaining. At that point, it quickly becomes a lot bouncier. Beyond 50% charge remaining, the bouncyness stops increasing. So, the measurement at the positive electrode does not match the results of the bouncyness test. Hypothesis 2 is proven wrong.

Hypothesis 3 is more difficult to test. How can you measure the effect of water alone, when the amount of water depends on all the reactions taking place? The researchers decided to carefully open up a battery, dry it, and put it back together. They found that this did change the composition of the gel: the gel turned solid and crumbly. However, the dried battery didn’t bounce much. So, hypothesis 3 is proven wrong as well.

It’s a good thing that the researchers kept Hypothesis 4 for last, because the measurements on the negative electrode gave some interesting results. The zinc electrode starts out as a suspension of tiny grains of zinc in the battery gel. As it starts reacting, the zinc forms a thin layer of zinc oxide around each grain. This zinc oxide doesn’t do much, it just sits there.

As the reaction continues, and each grain of zinc is covered with the oxide, new bits of zinc oxide start moving into the gel between the grains. Eventually, the zinc oxide forms tiny ‘bridges’ between the zinc grains. Afterwards, the rest of the zinc reacts to zinc oxide, which fills up any remaining gaps. The remaining mass is solid, and differently from the dried gel in the third experiment, not crumbly at all.

As it happens, the bridges form right when the battery starts getting bouncy. It seems tiny zinc oxide ‘bridges’ are quite elastic, they cause the battery to get bouncy. If you look back at the reaction formulas, you couldn’t have guessed that the zinc oxide actually forms in the shape of ‘bridges’ between grains. In chemistry, you often need to do real experiments to be certain of something. Theory only goes so far. Anyway, hypothesis 4 is the correct one.

Of course, in retrospect, we could’ve suspected that zinc oxide had something to do with it. Zinc oxide is used commercially to make a number of products, such as golf balls, more bouncy.

In any case, because of the good work of Bhadra, Steingart, and their colleagues at Princeton, we are now sure: alkaline batteries do get bouncy when they’re dead and this is caused by the formation of zinc oxide bridges. However, batteries already get bouncy when they’re still half full, so don’t throw them away as soon as they start bouncing.

1 Sources: http://www.princeton.edu/main/news/archive/S42/72/95S25/index.xml and http://pubs.rsc.org/en/content/articlelanding/2015/ta/c5ta01576f#!divAbstract

 

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