Electrolysis is a great way to introduce the concepts of elements and molecules, chemical changes and stoichiometry for higher level chemistry students. With this simple apparatus and a DC power supply, a dazzling electrolysis experiment can easily be performed.
- Oxidation–reduction reactions
Sodium sulfate solution, Na2SO4, 1 M, 500 mL
DC power source, 9 volts, or 9-volt battery
Graduated cylinder, 100-mL
Matches or lighter
Rubber stoppers, #0, 2*
Rubber-dam discs, 2*
Stainless steel electrode wires, 2*
Stirring rod, glass
Support block with two holes, 5" x 1⅝" x ½"*
Test tubes, 16 x 150 mm, 2
Test tube holding rings, polyethylene, 2*
Test tube rack
Tongs, crucible or forceps
*Materials included in kit.
Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Do not operate a battery or power supply with wet hands or in wet areas. Be sure the area is dry before turning on the power supply or closing the circuit on the battery. Follow additional safety instructions appropriate to your power supply. Do not use an AC power supply. AC will produce oxygen gas and hydrogen gas equally at both electrodes, which can be an explosive mixture. Please review current Safety Data Sheets for additional safety, handling and disposal information.
Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding.Sodium sulfate solution may be disposed of down the drain according to Flinn Suggested Disposal Method #26b. The Brownlee Electrolysis Apparatus should be rinsed with water, dried, and stored for future use.
Electrolyte Solution Preparation
To prepare 500 mL of 1 M sodium sulfate (Na2SO4) solution, add 71 g of sodium sulfate, anhydrous, to 400 mL of distilled or deionized water in the 600-mL beaker. Stir to dissolve and then dilute the solution to a final volume of 500 mL with distilled or deionized water.
Electrolysis Apparatus Setup
- Pour 100 mL of the 1 M sodium sulfate electrolyte solution into a 100-mL graduated cylinder.
- Dispensing from the 100-mL graduated cylinder, completely fill one test tube with the solution. (Save the remaining solution in the graduated cylinder for the other test tube.)
- Place a rubber-dam disc over the opening of the completely filled test tube. Make sure there is a tight seal, and that there are no air bubbles in the test tube.
- Place an index-finger on the disc to hold it in position. Then turn the test tube upside down, over the 600-mL beaker, making sure the solution does not spill out, and there are no air bubbles inside the test tube. Once this is accomplished, the finger can be removed from the disc and air pressure will keep it in place.
- Place the filled test tube, disc-side down, into the 600-mL beaker containing the sodium sulfate electrolyte solution.
- When the opening of the test tube and rubber disc are below the surface of the solution, a stirring rod or tongs can be used to free the disc from the test tube opening. The disc should float to the surface and can be removed from the solution with crucible tongs or forceps.
- Repeat steps 2–6 for the other test tube. Any remaining sodium sulfate electrolyte solution in the graduated cylinder can be poured in with rest of the solution in the 600-mL beaker.
- Place the support block on the 600-mL beaker and position the test tubes so their bottom-ends go through the holes in the block (see Figure 1).
- Slide a test tube holding ring onto the bottom-end of each test tube.
- Raise the bottom-end of each test tube as high as possible without lifting the test tubes completely out of the liquid.
- Place the U-shaped end of electrode into the solution and position the tip inside the test tube until the bottom of the U-shape bend touches the test tube (see Figure 2).
- Slide the wire of the electrode into the notch of the holding ring to secure the electrode in place (see Figures 2 and 3).
- Repeat steps 11–12 with the other electrode.
- Adjust the height of the test tube by raising the electrode and sliding the holding ring down the test tube.
B. Testing the Gases
- Connect the leads from a DC power supply to the wire electrodes. Warning: Make sure the DC power supply is off before attaching leads to the wire electrodes. Do not use an AC power supply.
- Turn on the power supply and adjust the voltage until a large amount of bubbles emanate from the electrodes (9 volts is recommended).
- Observe the electrolysis of water. What gases are being produced? Which test tube contains what gas?
- Turn the power supply off when the first of the two test tubes is approximately three-fourths full of gas (hydrogen gas).
- Unclip the leads from the wire electrodes and remove the electrodes from the apparatus.
- Slide the holding rings off the test tubes without removing the test tubes from the solution. Remove the support block from the top of the beaker and allow the test tubes to rest in the solution.
- While the test tubes remain in the electrolyte solution, use crucible tongs to carefully insert a rubber stopper into the opening of one of the test tubes. Make sure the rubber stopper is tightly sealed in the opening of the test tube by carefully, yet firmly, pushing the test tube down so that the rubber stopper presses against the bottom of the beaker.
- Once the rubber stopper is secure, remove the test tube and turn it right-side up. Dry the outside with a paper towel and place the test tube in a test tube rack.
- Repeat steps 7 and 8 for the other test tube. Do not place the test tubes immediately adjacent to each other in the test tube rack.
- Ignite the tip of a wood splint with a match or lighter. Let it burn until the tip has a small, steady flame.
- Obtain the test tube with the most gas, and place the flaming wood splint near the stoppered opening.
- Quickly remove the rubber stopper and insert the flaming wood splint.
- Observe what happens. What gas is this? (A pale blue flame and a loud “pop” or “bark” indicates the gas is hydrogen.)
- Repeat steps 10–13 for the test tube with the least amount of gas. Except this time blow out the flaming wood splint so that only the tip glows red. The oxygen gas test requires a glowing red splint, not a burning splint. (The glowing wood splint should burst into flames revealing the second gas to be oxygen.)
- If necessary, small amounts of tape or putty can be placed on the bottom edges of the block to prevent the block from sliding on the beaker or battery jar.
- The electrodes provided in this kit are inert stainless steel. The curvature of the U-shaped bend may need to be adjusted further in order to allow the arm of the electrode to snap into the notch of the holding ring while the tip of the electrode is inserted into the test tube opening. If necessary, use your hands or pliers to bend the U-shape accordingly.
- The volume of solution recommended to perform this activity allows the test tubes to extend above the support block in order to better observe the collected gas. A smaller volume of solution can be made and used, but the height of the test tubes will need to be adjusted appropriately.
- The concentration of sodium sulfate is not critical. It also does not need to be made from anhydrous sodium sulfate. Hydrated forms will also work. A 1 M concentration of sodium sulfate is only one option; 0.5 M, 2 M, 5 M, and even saturated solutions will also work. The higher the concentration the faster the reaction will proceed, because more current will be allowed to flow through the solution (see Discussion).
- Many other electrolyte solutions can be used for this demonstration, including dilute solutions of sodium hydroxide (NaOH) or sulfuric acid (H2SO4). The concentration is not critical, but 0.5 to 1 M solutions are recommended. Sodium chloride and hydrochloric acid solutions should not be used as electrolyte solutions because they will dissociate to form both oxygen and chlorine gas at the anode (unless the concentration of chloride ions is very small). (See Discussion.)
- Bromthymol blue indicator can be added to the sodium sulfate solution to show where oxidation and reduction occur (change in color as a result of pH change). Add enough indicator solution to give the solution a deep green color. Make sure the solution starts at a neutral pH (green solution). Add 3 M sodium hydroxide or 3 M sulfuric acid dropwise to adjust the pH, if necessary. If the solution is blue (basic) add sulfuric acid dropwise; if it is yellow (acidic) add sodium hydroxide dropwise. A good way to accomplish this is to dip a stirring rod into an acid or base solution, and then use it to stir the sodium sulfate solution until it is a nice green color. During electrolysis the solution will turn yellow at the anode (where hydrogen ions form) and blue at the cathode (where hydroxide ions form).
Correlation to Next Generation Science Standards (NGSS)†
Science & Engineering Practices
Using mathematics and computational thinking
Analyzing and interpreting data
Obtaining, evaluation, and communicating information
Planning and carrying out investigations
Developing and using models
Disciplinary Core Ideas
MS-PS1.A: Structure and Properties of Matter
MS-PS1.B: Chemical Reactions
MS-PS2.B: Types of Interactions
MS-PS3.A: Definitions of Energy
HS-PS1.A: Structure and Properties of Matter
HS-PS1.B: Chemical Reactions
HS-PS2.B: Types of Interactions
HS-PS3.A: Definitions of Energy
Energy and matter
Structure and function
Cause and effect
HS-PS1-1: Use the periodic table as a model to predict the relative properties of elements based on the patterns of electrons in the outermost energy level of atoms.
HS-PS1-2: Construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.
HS-PS1-4: Develop a model to illustrate that the release or absorption of energy from a chemical reaction system depends upon the changes in total bond energy.
HS-PS1-5: Apply scientific principles and evidence to provide an explanation about the effects of changing the temperature or concentration of the reacting particles on the rate at which a reaction occurs.
Electrolysis is a nonspontaneous redox reaction that is made to occur by passing an electric current through a solution under a sufficiently high voltage. For the electrolysis of water to occur, a high enough voltage (or electric potential) needs to be created that will add and remove the electrons from the water molecules, and a current needs to be generated so that electrons have the ability to flow to and from the electrodes. If there is a high enough voltage across the water solution, but the electric current is too low, then electrolysis will not occur. For this reason, electrolytes (ionic salts) are added to the water. The electrical resistance of pure water is too great to allow electrons to flow through it. Adding electrolytes, which dissociate to form positive and negative ions, reduces the resistance and creates an excellent environment for electrons (current) to flow. When a DC power supply is connected to the solution via electrodes, the electrons will flow from the power source to one electrode, which then becomes negatively charged. Since water is a very polar molecule, the negative charge will attract some water molecules. If the voltage across the system is high enough, the water molecules will be reduced (gain electrons) at the negative electrode to form hydrogen gas and hydroxide ions (Equation 1). In an electrolysis cell, the electrode where reduction occurs is called the cathode.
Meanwhile, the electrons in the other electrode are drawn toward the power supply and are removed, leaving this electrode with a positive charge. This positive charge also attracts some polar water molecules. If the voltage is great enough at this electrode, then electrons are removed from the water molecules and transferred to the positive electrode, producing oxygen gas and hydrogen ions (Equation 2). When water molecules lose electrons, they are said to be oxidized. The electrode where oxidation occurs is call the anode.
Equations 1 and 2 are known as half-reactions. These two equations can be added together to obtain the net equation for the electrolysis of water. (Equation 1 needs to be multiplied by 2 in order to balance the electrons produced by Equation 2 with the electrons consumed by Equation 1. The net reaction is a neutral solution. The four hydroxide ions and four hydrogen ions initially formed in the two half-reactions combine to give four water molecules which then cancel out with four water molecules from the reactant side.) The overall chemical equation for the electrolysis of water is shown in Equation 3.
From Equation 3 it can be seen that two moles of water molecules form two moles of hydrogen molecules and one mole of oxygen molecules. This stoichiometric relationship explains why the volume of hydrogen gas produced is twice the volume of the oxygen gas.
When choosing an electrolyte, it is important to choose one that will not interfere with the results of the electrolysis of water. Sodium chloride (NaCl) and hydrochloric acid (HCl) should not be used as electrolytes because chloride ions are more easily oxidized than water molecules during the electrolysis of water. This means that poisonous chlorine gas will initially form at the anode, instead of oxygen gas. Once the chloride ions are consumed, oxygen gas will be formed, but the collected gas will now be a mixture of the two gases. Sodium sulfate is a good electrolyte to use for electrolysis because water is both oxidized and reduced much more readily than sulfate ions and sodium ions, respectively, and the salt solution is neutral.
Inert electrodes need to be used so that they, too, do not interfere with electrolysis. Copper wire electrodes are not used because copper is more easily oxidized than water. This means that a copper electrode will undergo oxidation (to copper ions) rather than the desired water. Oxygen gas will not be produced and the copper electrode may become corroded with a blue-green insoluble basic copper(II) sulfate (if the sulfate ion is present). Other possible insoluble salts that may form on the copper electrodes include copper(II) hydroxide and copper(II) oxide, depending on the electrolyte.
Bilash, B.; Gross, G. R.; Koob, J. K.; A Demo a Day™—A Year of Chemical Demonstrations; Flinn Scientific: Batavia, IL, 1995; pp 250.
Brown, T. L.; LeMay, H. E.; Bursten, B. E.; Chemistry: The Central Science, 6th edition; Prentice Hall: Englewood Cliffs, NJ, 1994; pp 720–751.