The generation of clean, efficient and environmentally friendly energy is a major challenge for science and engineering. Carbon dioxide (CO2) levels are 40 % higher than they were at the beginning of the industrial revolution, and most of the increase has taken place since 1970. During the last 40 years, global energy consumption has accelerated, and the rise in CO2 is largely due to the combustion of fossil fuels. There are growing concerns regarding the increasing greenhouse gas emissions from fossil fuels, as well as diminishing fuel reserves, and therefore scientists are looking towards fuel cell technology for future power. Fuel cells have a wide range of potential applications, including portable device and transport applications.
A fuel cell is an electrochemical cell, which converts chemical energy into electricity via a chemical redox reaction. Fuel cells convert the chemical energy of a fuel gas directly into electrical work. This is much more efficient than using combustion of fuels to drive moving parts. The redox reaction produces a potential difference between the two electrodes. The electrode at which the reduction occurs becomes the cathode, which is positively charged, and the oxidation occurs at the anode, which is negatively charged. Electrons flow from anode to cathode.
Fuel cells can continuously operate as long as fuel is constantly supplied into the system. This is an advantage compared to batteries, which need to be recharged or recycled.
Exercise 1: Get students to list the advantages and disadvantages of using fuel cells and batteries. Download the worksheets here.
Hydrogen fuel cells use hydrogen and oxygen as the reactants:
Exercise 2: Ask students to write a balanced equation of the combination of hydrogen and oxygen to produce water. HT students may also be able to write the half equations for the two electrodes. Download the worksheets here.
|The alkaline fuel cell:(-ve electrode) = Anode: H2 + 2 OH– → 2 H2O + 2 e–
(+ve electrode) = Cathode: 1/2 O2 + H2O + 2 e– → 2 OH–
All the documents for this experiment can be downloaded here.
To make a battery and a hydrogen fuel cell
YOU WILL NEED:
- 6 small potatoes
- 6 x 2 pence (or 1 pence) pieces
- 6 x galvanised screws/nails
- 7 insulated wires with crocodile clips
- a voltmeter
- an LED
- 2 graphite electrodes (2 pencils work well, sharpened with a knife of both ends to extend the graphite, or 6 x 0.7 mm leads for propelling pencils, taped together per electrode)
- Blu tack
- 0.4 M KOH (aq)
- 1 x 9 V battery
- plastic knife
- universal indicator paper
Part 1 The Potato Battery
- Push a penny about half way, into a potato.
- Push a nail or screw into the potato, about 3 cm away from the penny, so that only 1 cm of the nail remains above the potato surface.
- Connect a crocodile wire to the penny, and to one side of the voltmeter, and another, to the nail and the other side of the voltmeter, and record the voltage.
- Now disconnect the voltmeter, and attach the crocodile clips to the LED – does it light? You need to connect the LED the right way around; connect the nail to the negative side of the LED, i.e. the shorter wire, and the copper coin to the positive, longer, wire of the LED.
- Make up 5 more potato batteries, and connect them together, measuring the voltages each time. This activity may need to be done as a class, connecting up with other students’ potato batteries.
|Number of potato batteries||Voltage (volt)||Did LED light?|
6. Take a slice of potato and measure its pH.
- The potato battery is made up of two electrodes, what metals are they made of? What part of your battery is behaving as an electrolyte?
- What is the pH of the potato?
- What parts of the cell are being ‘used up’ in the chemical reaction, i.e. what is behaving as the fuel within the battery? What will happen when these eventually run out? Are they easy to keep renewing in the system?
- What happens to the voltage as you connect up more potatoes? Can you light the LED?
Part 2 The Alkaline Hydrogen Fuel Cell
Step 1 – Make your fuel
First you need to make your fuel; your hydrogen and oxygen gas. We will do this by splitting water, using electricity:
We USE electrical energy to force a chemical reaction – this is called ELECTROLYSIS
1. 2H2O → __H2+ __O2 balance the equation
2. In the process of electrolysis, what type of energy is being converted into chemical energy?
_________ Energy → Chemical energy
Take the two pencil electrodes, and tape them together, using a piece of blu tac in between the pencils, to space them, about a pencil width apart.
Pour 100 mL potassium hydroxide solution in to your beaker. WARNING – this solution causes skin irritation and eye irritation. WEAR GOGGLES. If spilt wash hands immediately.
Take your 9 V battery and connect one end of a wire to the +ve terminal of the battery (circular), and the other end to one of the pencils. Mark the pencil with a ‘+’. Use the other wire to connect the –ve terminal (hexagonal) to the other pencil. Mark it with a ‘-‘.
As soon as you connect the battery, the electrolysis should start to happen. Leave it to make your fuel for 2-3 min.
- What can you see happening at the pencil tips in the solution?
- Which electrode (+ve or –ve) produces bubbles faster?
- Which gas (oxygen or hydrogen) do you think you produce more quickly, using your answer to question 1?
- Sketch a diagram of your electrolysis setup below:
The electrolysis process occurring:
(+ve electrode) = Anode 4 OH–(aq) → O2(g) + 2 H2O(l) + 4 e–
(-ve electrode) = Cathode 4 H2O + 4 e– → 2 H2(g) + 4 OH–(aq)
Step 2 – Make your fuel cell
After 2-3 min of making your oxygen and hydrogen fuel, carefully disconnect the battery. TRY NOT to disturb the bubbles as this is your FUEL!
Connect the electrodes up to the voltmeter, and observe the voltage over 5-10 min. Do this carefully.
- Can you measure a voltage? What happens to the voltage over a few minutes? Why do you think this is?
- Sketch the apparatus for the fuel cell that you have made.
- In this process occurring in the fuel cell, what type of energy is being converted into electrical energy?
________ Energy → Electrical Energy
4. What is the overall product when hydrogen reacts with oxygen? Why is this a ‘greener’ process than combustion?
The alkaline fuel cell:
(-ve electrode): 2 H2 + 4 OH– → 4 H2O + 4 e–
(+ve electrode) : O2 + 2 H2O + 4e– → 4 OH–
Step 3 – Can you light an LED?
Remake your fuel (repeat step one), and then check that you have a measurable voltage. Connect your pencils to an LED and see if it lights? How many fuel cells are needed? You will probably have to connect up with other students fuel cells to get enough power!
2H2O → _H2(g) + _O2 (g) balance the equation
In the process of electrolysis, what type of energy is being converted into chemical energy?
Download the following Origami Fuel Cell, and print onto A4 paper, preferably in colour, to make a 3D cube that shows alkaline and solid-oxide fuel cells:
This resource has been adapted from a teaching resource by Professor Peter Slater.
In the research lab
Researchers working with Professor Peter Slater, at the University of Birmingham, are developing a range of new ceramic materials for the electrode and electrolyte materials for solid-oxide fuel cells. Current work involves the development of new ceramics that can reduce the operating temperatures of solid-oxide fuel cells, to extend their operating times, as current zirconium-oxide type structures only conduct oxide ions well in the 800-1000 oC temperature range.
Publications of this work
Abstract above reprinted with permission, copyright (2017) American Chemical Society http://pubs.acs.org/doi/abs/10.1021/jacs.5b13409
This work is licensed under a Creative Commons Attribution 4.0 International License.