Kids' Basics: The Battery Experiment

Teaching kids about batteries

Daniel Koch

Issue 34, May 2020

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An open-ended and adaptable experiment to keep you thinking while at home.


This month, we’re going for a different approach for Kids’ Basics. With isolation our current reality thanks to COVID-19, we’re not going with our regular-type step-by-step build a basic circuit explanation. This is fine most of the time, but it’s an ‘over-and-done-with’ project, with little room to expand your thinking once the project is complete. In education circles, we call this a ‘closed’ task.

This month, we’re presenting an ‘open’ task investigative experiment and a tool with which to conduct it. We’re going to show you how to build your own batteries, and then step you through running an experiment to see which one performs the best. It will be variable and adaptable, with no one right answer and no one fixed project. That means you can make it simple or complex, and go for an hour or a month. While we start with a fair bit of theory, younger makers can skip to the experiment and run it with just a little help from an adult. The result is more like our ‘Classroom’ column, but with more experiments and aimed at kids.

Batteries underpin modern life as we know it. From the basics such as torches to our Information and Communication Technology (ICT) devices, to household alarms, emergency lighting, clocks and remote controls, they’re in nearly every direction we look. They also come in a variety of forms. You may have seen someone loading tiny batteries into their hearing aids, while goods are delivered to your local supermarket in a truck which has a much bigger starter battery than the ones you may see under the bonnet of the family car.


A battery is a collection of individual cells. A cell is an arrangement of chemicals that take part in a chemical reaction, which produces an electric current. The basic components are two electrodes of different material, and an electrolyte that helps make the reaction happen. More on this in a moment. All materials on earth (and indeed, in the universe) are chemicals in one form or another. The word ‘chemical’ in social use often implies something manufactured and often toxic, but in science, the word means any pure or compound substance.


For our purposes, atoms are the smallest independent particle of matter (stuff) in our universe. They are made up of Protons, Neutrons, and Electrons. The atom is the smallest independent particle because protons, neutrons, and electrons generally don’t exist on their own, only within an atom. There are smaller particles and times where protons, neutrons, and electrons exist by themselves, but that isn’t taught until senior high school and doesn’t affect our battery lesson at all.

Most atoms have a ‘nucleus’ of protons and neutrons, surrounded by layers of electrons. The exception is Hydrogen, which has no neutron. Protons have a positive electrical charge, while electrons are negatively charged. Neutrons have no charge, hence the name coming from the word ‘neutral’. A proton and an electron balance each other, but don’t touch. Instead, the electrons fly around in layers a lot like satellites orbit earth. Starting from hydrogen, each atom has more protons and neutrons in its nucleus than the one before, and therefore more electrons. The layers of electrons in the ‘shell’ have a fixed number; the first shell can hold two electrons, the second can hold eight, the third eighteen and the fourth can hold thirty two.

The thing is, the layers are not always full. Each atom contains, generally, the same number of electrons as it has protons. Oxygen, for example, has eight protons, and eight electrons. The first shell has its two electrons, while the second shell is left with six. However, electron shells like to be full. Atoms may give away electrons, or take them, depending on how many they need to reach a state where their shells are complete. Exactly how this works is a chemistry lesson (or several) on its own, but for now, you just need to know that unless an atom’s electron shells are full, it can react chemically with other elements. Also note that the shells do not always fill as described above, and sometimes a shell will start filling before the shell below it is full. The outer shell is the one that reacts, or takes or gives electrons.

Any substance you encounter in the world can be grouped into three main categories: Elements, which are pure substances all of the same atoms; compounds, in which different atoms are stuck to each other in different ways to form molecules; and mixtures, where different elements or compounds are mixed together but not chemically bound together. All pure elements are shown on a table called the Periodic Table of the Elements. It shows the atomic number, which is the number of protons in the nucleus, as well as other information that doesn’t concern us now. Some versions display the number of electrons in each shell, which is useful to us.

The last thing you need to know is about ions. There are situations where an atom has more or less electrons than protons. It, therefore, has an electrical charge and is called an ion. Ions want to satisfy the balance between protons and electrons, so they will react with other elements or compounds to achieve this. They can also be carried by ‘electrolytes’, chemicals which can weakly bond with an ion and allow it to move. They are usually liquids but liquids in a chemistry sense may still be very thick and not runny at all.

That chemistry lesson is very compressed, and probably wouldn’t even stand up to a high school class, however, it does the job for now.

That brings us to the battery. Some chemical elements and compounds will react together, and when they do, electrons move to satisfy the shells of different atoms. Some substances need electrons, others want to give them away. Some reactions can produce an excess of electrons while others produce a deficit, or too few electrons. Some types of chemical bonds are stronger than others. Some chemicals will strip an atom from the rest of the mass as an ion, leaving behind an excess or deficit of electrons. The right combination of these things forms an electrochemical cell. A battery is two or more cells connected together. An excess of electrons builds up on one electrode as the ions carry across the electrolyte, while there is a lack of them on the other electrode. If the electrodes are connected by a conductor, electrical current flows.


Batteries aren’t a new invention, although some of the chemical combinations used to make them are very recent technology. A quick look at their history may help shape your experiments moving forward.

The generally recognised first battery invented was a creation of Alessandro Volta. He built on observations of another Italian scientist, Luigi Galvani, who noticed muscle movements when dissecting animals with metal implements. Volta’s device was a series of metal plates separated by cloth or cardboard soaked in saltwater. Saltwater is an effective electrolyte. This creation was called a ‘pile’, and Volta thought the electrical current came from the metals touching each other. We now know that’s not accurate, but the experiment produced electricity nonetheless.

Volta’s pile had a few issues, the main one being the weight of the metal plates. He had used zinc and copper in his final design, and the pressure squeezed out the electrolyte from the cloth. Scotsman William Cruickshank laid Volta’s pile on its side in a box, which solved this problem. The other main problem with Volta’s battery was that it produced hydrogen gas around the copper electrode, which reduced the surface area of the copper exposed to the electrolyte. Despite this, our experiment will be based on the Voltaic pile, albeit in Cruickshank’s horizontal arrangement, because it is the most easily constructed.

The next major development of relevance for us was from John Frederic Daniell, an English chemistry professor. He got around the hydrogen issue by using two different electrolytes, separated by a porous membrane. He used a copper pot as both container and electrode, with copper sulphate solution in it. Sitting in this was a clay container with sulphuric acid in it, with a zinc electrode in the middle. The clay barrier allows transfer of ions but keeps the solutions generally separate. Others improved this system over time.

Unfortunately, that’s where our history lesson ends, because from here on out, materials used go beyond household availability. Even the ‘heavy duty’ type household batteries, which are often called ‘carbon-zinc’ batteries, actually use manganese dioxide and an electrolyte of ammonium chloride.


With the chemistry behind batteries covered, let’s recap the physical essentials. We need two electrodes of different materials, so that each has a different number of electrons in its outer shell. We also need an electrolyte, a chemical to carry the ions between the electrodes. If we’re making the Daniell cell, we’ll need two electrolytes and a porous barrier between them.

In any electrochemical cell, one electrode becomes positively charged, and the other negatively charged. Because electrons have a negative charge, it’s the electrode with an excess of electrons which becomes the negative electrode, called the anode. The positively charged electrode is called the cathode. The challenge here is that electrons flow from the negative electrode, to the positive electrode, where there are less electrons. This is counter-intuitive but happens because the electrons are negatively charged. The terminology does not reflect where there are more or less, like positive or negative would in maths.


For this to happen, we need materials with different numbers of electrons in their outer shells. For our experiment, you could use the periodic table to choose materials, or you could just try some if you don’t have confidence reading the periodic table. We’ll need an electrolyte, too. We’ve already given enough theory, so we’ll just say for now that most electrolytes used in basic electrochemical cells are acids. We don’t mean highly corrosive car battery acid, either. Acids around the house can be things like vinegar, lemon or orange juice, the starches in potatoes, or even grape juice. And that leads us to our first experiment.

Before we do, we’ll introduce two tools you’ll need for this experiment. All electrochemical cells and batteries show a higher ‘potential difference’ when unloaded. Potential difference is the name for the difference in charge between one electrode and the other, and is measured in ‘volts’, hence the somewhat incorrect but convenient term ‘voltage’.

To make the measurement more realistic, we need a load, and the easiest way to do that is with a resistor. This way, we don’t have to worry that a voltage may be too high or low for a given component like an LED or light globe. We suggest a 100Ω (100 Ohm) resistor for this.

The second tool after a consistent load is a multimeter. This tool is available from major electronic retailers, starting from around $10. It measures a variety of characteristics, but we are interested in its ability to measure voltage and current. It will also tell us the ‘polarity’ of the current, which is the way it is flowing.


These parts are the electronic parts required for all of the experiments following. However, you'll need to check each experiment for the additional parts required to complete them!

ELECTRONIC Parts Required:JaycarAltronicsCore Electronics
1 x 100Ω Resisitor *RR0548R7534CE05092
1 x MultimeterQM1529Q1129TOL-12966
10 x Alligator Clip Jumper Leads *WC6010P0415AP0415
1 x LEDZD0150Z0800CE05103

ELECTRONIC Parts Required:

* Quantity shown, may be sold in packs. You’ll also need a breadboard and prototyping hardware.

Experiment 1:

Lemon Battery

The first experiment involves making a battery from lemons. Yes, lemons! Lemon juice is acidic, filled with acetic (or citric) acid, with a few other acidic compounds thrown in. It happens to make a viable electrolyte. We need two dissimilar (not the same) metals as well. For this, we need copper and zinc.

ADDITIONAL Parts Required:
Experiment 1 Recording Sheet
2 x Lemons
1 x Zinc Metal Sample (See Text)
1 x Copper Metal Sample (See Text)
1 x Multimeter
1 x 100Ω Resistor
2 x Alligator Clip Jumper Leads

ADDITIONAL Parts Required:


Copper and zinc aren’t as hard to find as you might think. Zinc is used to ‘galvanise’ other metals to protect them from chemical reactions from weathering, which otherwise corrode them. This means most galvanised metals are zinc coated and suitable for use. Nails are a great place to start but as you’ll see, surface area matters in this experiment. You may be able to use lots of nails side by side or use galvanised washers instead. Look for ‘hot dipped’ rather than ‘electroplated’ galvanising, as it’s much thicker. You may also consider coach screws or bolts instead of nails.

Copper isn’t quite as easy to find. Copper nails were once common but now rare, and because they’re often used only decoratively, they’re usually thinly electroplated, too thin for our use. The main option is copper water pipe. You can buy this in one metre lengths from chain hardware stores, and have an adult cut small sections off. If this is not something your adults can do, most hardware stores and plumbing suppliers sell copper joiners as well. These and the pipe sections can be used as-is or hammered flat. Other sources could be copper-coated welding rod, earth rod, or tubes from hobby shops.

Step 1:

Arrange your materials: The lemon, electrodes, resistor load, jumper wires and multimeter. Also, print the ‘Experiment 1 Recording Sheet’ or draw one up yourself.

Step 2:

Carefully push the electrodes into the lemon, roughly in the position shown. Measure the distance between them and write it in the ‘separation distance’ field.

Step 3:

Connect the jumper wires to the electrodes. Be careful they don’t touch. Nothing dangerous will happen, but your battery will go flat very quickly.

Step 4:

Plug the leads into your multimeter, and set the dial to measure voltage. On most multimeters available today at retailers, the setting will have the symbol shown. If you have a manual range meter, set it to 2V.

Step 5:

Measure the voltage between the two electrodes. Put the red lead on the zinc electrode, the anode, and the black on the copper electrode, the cathode. Write the voltage down on the recording sheet under ‘Open-circuit Voltage’. Check to see if the polarity symbol appears in the multimeter display.

Step 6:

Connect the resistor between the jumper leads and check the voltage across the electrodes with the multimeter. Write the voltage in the ‘Loaded Voltage’ field on the recording sheet. Is it different? There is a column on the recording sheet for ‘Observations’ where you can write down things you notice.

Step 7:

Disconnect the resistor from the anode (zinc) side of the jumper lead. Connect the black lead to the jumper wire, and set the multimeter to measure current on the smallest setting. The photo shows how this looks on most multimeters.

Step 8:

Touch the red probe of the multimeter to the free end of the resistor. You should see a reading on the display. Write it down in the ‘Loaded Current’ section of the recording sheet.

Step 9:

Take the red probe off the free end of the resistor and place it on the jumper lead, bypassing the resistor. Write the current value down in the ‘Short-Circuit Current’ field.

Step 10:

Take the electrodes out of the lemon, clean them with a damp rag or paper towel, and insert them again, a little closer together.

Step 11:

Repeat the above set of measurements with the new electrode positions, recording each where relevant. Your sheet should end up with several lines full, depending on how far you move the electrodes each time.

Step 12:

With a fresh lemon and new set of electrodes, repeat the entire above experiment, but this time, roll the lemon on a table with a bit of downward pressure, so that the internal membranes of the lemon are broken.


Look at your results on the recording sheet. Are there trends you notice? How do the observations and measurements change when the electrodes are moved further apart or closer together? Were the numbers any different when you used the lemon with the mooshed membranes? How do the measurements and observations compare to what you thought might happen after reading the theory above?

Experiment 2:

Mixing it Up!

Now that you’ve completed this scaffolded task, which was closer to a closed task, we’re heading into something more open. We suggested zinc and copper because they are known to work well. However, they are not the only metals which can be used to make an electrochemical cell. Now we’re going to investigate other metals. This is where things get a little more open-ended. Your task is to gather as many metal samples from around your house as you can. Ask adult permission of course, then gather anything from steel nails to fishing sinkers. Then, using a lemon, try each combination and record the results. We’re going to use the open circuit voltage and loaded voltage for this test.

ADDITIONAL Parts Required:
Experiment 2 Recording Sheet
Various Metal Samples (See Text)
1 x Lemon or More
1 x Cleaning Rag or Paper Towel
1 x 100Ω Resistor
1 x Multimeter
2 x Alligator Clip Leads

ADDITIONAL Parts Required:

Before you start, list all the variables you can think of on the recording sheet. In a fair test or experiment, we only change one variable and keep all others the same. The variable we are changing is the combination of metals. What are some things that need to stay the same? See the Analysis section if you get stuck.

Step 1:

Print or write your own ‘Experiment 2 Recording Sheet’. Gather your materials and arrange them on a table.

Combining Materials

The number of combinations is not simply double the number of metals. Have a look at the example list here. Do you notice any patterns? The mathematically-minded readers might even like to try to find an algebraic formular to describe this. For everyone else, Write your first material once for every other material. We started with copper. Then, your next material will have one fewer options, because it was already paired with your first option above. Your next material will have one fewer options again, and so on.

Step 2:

Write different combinations of metals in the ‘Combinations’ column. You can try different combinations with one common metal. See the ‘Combining Materials’ box for details.

Step 3:

Make a hypothesis (a prediction based on what you already know and have observed so far) about which metals will work together, and record it at the top of the sheet.

Step 4:

Measure the open-circuit (unconnected besides the meter) voltage of each combination with the multimeter set to Voltage. Take note of the polarity symbol to determine which metal is the anode and which is the cathode, and record this as well as the open circuit voltage. Because the meter is set up for 'conventional current flow', the polarity is opposite what we might think. In a circuit, current flows from positive to negative. Once electrons were discovered and understood, it was realised that electrons flow the other way. That's why the anode electrode is where the electrons flow from and the cathode where they flow to, yet the negative side of a component is the cathode and the positive side the anode.

Step 5:

Connect two jumper leads, with the 100Ω resistor in the middle, to the electrodes and measure the loaded voltage. Record this as well.

Step 6:

Take out one or both electrodes, clean them with a rag or paper towel, then try a different combination. Check the lemon at each stage and use another if the holes are becoming too dirty for good contact with the metals.


Starting with the variables, hopefully you had the distance between the electrodes, and the surface area of the electrodes as things to keep the same between tests. These could strongly influence your results. There are other things too, like whether the acidity of the lemon was the same if you had to change. Revisit your variables list and see if you would add or change anything now.

Looking at the data you have gathered, which set of electrodes gave the highest voltage? Were there any surprises in the differences between open-circuit and loaded voltages? Were there any combinations which had a much bigger or smaller drop between open-circuit and loaded voltages?

Experiment 3:


After experimenting with different metals, we’re going to now have a go at using different electrolytes. As noted, a fair test needs to be kept consistent, with all variables kept the same except for one that is carefully changed. There are actually three types of variables. The Independent Variable is the one that we are changing. In this case, that is the electrolyte. The Controlled Variables are the ones that are going to be kept the same. We’ll ask you to think about those in a moment. The Dependent Variables are the ones which change because of changes we make. In other words, the variables we observe, measure, and record. In this test, the dependent variables will be the voltages measured, both open-circuit and loaded.

The sources of electrolytes can be varied. Many fruits and some vegetables work, as will some household liquids. Saltwater is one suggestion, as is vinegar. We’re not suggesting both of these will work, we’re just showing you which directions you might head in. With both food and household chemicals, ask your adult first. Some chemicals are dangerous and even if they’re considered safe, if there is a label on the package, read the details or have your adult read them aloud.

If you’re using liquids, they will need to go into a container. Think about controlled variables and how this will compare to the food items you might be using. How are you going to keep all variables the same?

ADDITIONAL Parts Required:
Experiment 3 Recording Sheet
Various Electrolyte Samples (See Text)
Plastic or Glass Container for Liquid Samples
1 x 100Ω Resistor
1 x Multimeter
2 x Alligator Clip Leads

ADDITIONAL Parts Required:

The point of this experiment is to explore and observe. Some of the products described or shown will not work. we don't want to give all the answers away, we want you to find your own answers by investigation.

With that in mind, have a good look around your house for anything that might work, but: READ LABELS FOR SAFETY INFORMATION! and ask an adult about anything you don't already use, like fruit juice. Be curious, but be safe.

Step 1:

Print or draw your own recording sheet, and gather your materials. You will need to choose two electrodes, based on your previous experiments.

Step 2:

Fill in the sections of the recording sheet for your hypothesis, controlled variables, and the list of electrolytes you’ll try. The electrode materials are two of the controlled variables, but we have a place for those separately.

Step 3:

Insert your electrodes into an electrolyte to be tested. This might be something solid like a fruit or vegetable, or a liquid in a container.

Step 4:

Measure the open-circuit voltage with the multimeter, and record it on the sheet. Remember to set your multimeter to the voltage setting.

Step 5:

Measure the loaded voltage using two jumper leads and the 100Ω resistor.

Step 6:

Remove the electrodes and clean them with a rag or paper towel. If you had liquid in a container, clean the container too.


Analysing experimental results should be becoming familiar to you now. Start by looking for trends or patterns in the recorded data, and use it to see if your hypothesis was correct. Which electrolyte turned out to be the best performing? How else could we test this?

Further Experiments:

Now we’re going to step up the challenge. So far, we’ve presented one structured experiment, and two scaffolded but open tasks. Now, we’re turning you loose. As mentioned earlier, different chemical combinations perform differently to each other. You have tested different metal combinations in lemons, and different electrolytes with the same electrodes. What if we present you with a new idea? It is possible that even though one set of metals performed the best in lemons, those metals may not be the best for other electrolytes. There is a big combination now of different metals to test with different electrolytes. Testing this is your challenge.

Some questions to get you started are: How are you going to test this? What is your exact hypothesis or aim going to be? How are you going to record your observations? Which variables will you keep controlled and which will you vary? What materials and supplies do you need, if you haven’t got everything already?

Challenge 1:

Voltage Across

Now that you’re familiar with home-made electrochemical cells, it’s time to back-track. Earlier, we tested which electrochemical cell in each experiment performed best based on open-circuit and loaded voltages measured at the start. However, here’s a hypothesis for you to test: ‘Different chemical combinations in an electrochemical cell will produce different voltages for different times’. This is a hypothesis only. We aren’t stating this as fact yet, though it does give you a test to run. Maybe the highest voltage cell you made is not the best because it goes flat fast, and a lower voltage combination of electrodes and electrolyte may produce the lower voltage for much longer. If that is the case, which is the ‘better’ cell, and why?

The answer really depends on how you justify it, and what criteria you value over others. If you’re stumped for ideas, see the Hints section at the end of the article. This test will keep you going for a while, we expect. You might run several combinations side by side and test one after another at the same interval, or complete one test then start another. It will really depend on how many jumper leads and resistors you have!

Challenge 2:

Current Through

You will probably have noticed by now that the voltages and currents out of these cells are very small. One thing that we haven’t done since the first experiment is to test the current of the cells, but none of them are very high. How do you make them more useful? Connect them in series to increase the voltage, and parallel to increase the current! Choose one of the cell designs described previously. It may be the original lemon cell, one of your other metal combinations or a different electrolyte, and make sure you can make quite a few of them. Next, you’ll need to measure the voltage and current of each. Measure this as the loaded voltage and current. Then, you’ll need to choose a load to drive. As we discovered in DIYODE Issue 004, you certainly won’t be charging your mobile phone, but you may easily light a small light bulb or an LED. From the data from the supply, write down the current and voltage needed to drive your chosen load. An LED normally needs a resistor to survive its full rated life hours, because they do not current limit themselves, even with the voltage matched. However, for such low-power sources such as these, you can leave the resistor out.

Connect your load to one cell. Measure the voltage across it and current going through it. This is unlikely to light the LED and almost certainly won’t light the globe if you’re using one. This will tell you how the load has affected the performance of the cell. Record your measurements. At this point, how many cells do you think you will need to fully supply our load? Remember that LEDs are polarised, they have a negative and positive side. With such a small power source, you can try both ways if you’re unsure.Now, make a second cell, and connect it in series with the first. Your load now connects across the outer electrodes. You’ve made your first battery! Measure voltage and current again, and record them. Keep adding cells, measuring, and recording, until you reach the right voltage for your load. How many cells have you had to connect in series?

Make more cells so that you have a second set of cells connected in series the same as the first. Connect it in parallel with the first. Measure the current and voltage again. Is it enough to drive the load yet? An LED will start showing light well before its rated current, and so will a light globe. They just won’t reach full brightness until the specifications are met.

If the battery is still not producing the full current of the load, add another series-connected set in parallel to the first two, and so on until the load is fully supplied. How does this compare with your predictions at the start?


EXPERIMENT 1: In step 6, the loaded voltage should be lower than the open-circuit voltage.

EXPERIMENT 2: Because the multimeter works in conventional current flow, the electrons moving from anode to cathode are opposite to the convention. So, if the black lead is on the anode and the red lead on the cathode, the display will show correct polarity. If the red lead is on the anode and the black on the cathode, the ‘-’ sign will show on the display next to the voltage.

EXPERIMENT 3: Don’t forget to try the internet favourite, the potato! Fruit juices and many other drinks work well too. Does milk work? Try powders from the pantry like tartaric acid, sugar, salt or bicarbonate soda, mixed with water. Coca-Cola is high in phosphoric acid, and you doctor would probably argue that this is the best use for it.

CHALLENGE 1: Which is the better cell really depends on what you expect it to do. Some batteries need to provide a lot of current fast, for a short amount of time, like your family car battery. Others, like smoke detector batteries, need to produce a little current for a long time. If the higher voltage goes flat before it’s done its job, it isn’t the best. On the other hand, if the battery that lasts the longest never gives the current or voltage required, even in series and parallel as you’ll explore later, then it’s no good either. Size may be a factor too. Test each battery’s output against time. Leave the 100Ω load connected and measure the voltage. Will you test every minute, every ten minutes, every hour, or a different time?

Challenge 3:

In the theory section, we described the Daniell cell, a type of electrochemical battery. This is similar to a Galvanic Cell pictured here. The Galvanic cell is two half-cells of an electrolyte and electrode each. The half-cells are connected by a salt bridge, which allows ions to be transferred. The salt bridge in labs is often a glass tube filled with salt solutions, but in kids’ chemistry sets, is often just filter paper soaked in saltwater. You might come up with a different way to make yours.

Can you make a Daniell or Galvanic cell? Will you use a salt bridge or some porous container? You might be adventurous enough to make a thin-walled container out of air-drying clay. You might try an aluminium drink can as an outer container-electrode, but be careful to have an adult cut the top off and cover the sharp edges with tape. Other metal containers work too, or you could use a plastic or glass outer container with a metal electrode (or electrodes connected together) in it.

Where would you source electrolytes for that? The electrolytes used are generally salts of the metals involved. If you use a copper electrode and a zinc electrode, you would most probably be using copper sulphate and zinc sulphate as they dissolve in water as the two electrolytes. Where does someone get chemicals like that? They’re a bit more common than you might think. Zinc sulphate is used as a powder for people with zinc deficiency (too little zinc in their bloodstream), while copper sulphate crystals are used as a pool chemical. Both are used as a fertiliser and nutrient correction additive for plants, and both are available in the garden section of most hardware stores. If you choose other metals and look for other metal salts (sulphates and nitrates), look for similar sources.

Whatever you use, discuss everything with your adult, read all safety precautions on the labels to make sure it’s suitable, and keep the original containers with your experiment. Also record what you do at every step of the way so that at all times, every chemical used in every part of the experiment is documented. Write your planned procedure out first before you start and keep the record with your experiment. We’re not even going to provide a Bill of Materials for parts list, because there are far too many possibilities.


Many Kids’ Basics readers will be using a multimeter for the first time, while others will find a few quirks while running these experiments. The first concerns polarity, which means which side is negative and which is positive. Normally, you connect the red probe to the positive side of anything, and the black probe to the negative. The polarity of the measured voltage is indicated on the screen, usually by the presence of a ‘-’ sign if the probes are reverse-polarised. This happens if the red probe is connected to the negative of the source voltage, and the black to the positive.

However, this is based on ‘conventional current flow’ from positive to negative. As we discuss in the article, electrons actually flow from the negative to the positive, because the terms ‘negative’ and ‘positive’ were decided before electrons were discovered and understood. Before that, researchers working with electricity had theories about what they were seeing, but nobody knew. When electrons were discovered, it was found that electrons are particles that have a negative charge. The electrons move from the anode to the cathode. However, the anode of a component is its positive side where we think of current as flowing in, and the cathode is the negative side where current flows out. Have a close look at the diagram to follow along. The arrows are colour coded to help.


Finally, measuring current works differently on different meters. Traditionally, all multimeters had a ‘COM’ socket, for ‘common’, where the black lead goes. There were two red sockets, one for current measuring, usually marked ‘A’, for ‘Amperes’; and another for everything else, like voltage and resistance. To measure current, you plugged the red lead into the ‘A’ socket and turned the dial to the current setting.

The meters are still made the same way, but many modern meters actually measure currents below a certain value through the ‘everything else’ socket. This is commonly anything below 500mA, but that limit varies in some meters. You still turn the dial to the current setting, usually labelled ‘DC A’ or ‘A DC’. Have a close look at the labelling at the sockets to find out.