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Petri Dish Electrolysis Activity

By Bruce Wilson
Product Manager, Chemistry

petri dishIntroduction

This simple microscale electrolysis activity yields 2 pure diatomic gases, hydrogen and oxygen, from water in a petri dish. With the addition of Bogen universal indicator solution to the reaction, pH color changes help students understand what is happening at each electrode in the petri dish. You may use the activity to introduce reduction-oxidation reactions and some of the basic terms of electrochemistry, such as electrode potential, anode, and cathode. To highlight real-world applications of electrolysis, discuss the industrial production of chlorine from saltwater. Also discuss current research into the practicality of using water electrolysis as a means of producing hydrogen fuel.

Content standards

Grades 9–12

  • Structure and properties of matter
  • Chemical reactions



Wear safety eyewear.

Preparation and procedure

  1. Put distilled water into a petri dish (lid or base) until the bottom is covered.
  2. With a graduated pipet, add 0.5 mL of Bogen universal indicator solution to the water and swirl to mix.
  3. Add a small spatula of the electrolyte, sodium sulfate, to the water and swirl to dissolve.
  4. Clip wire leads to the 9-V battery terminals. To each lead’s other end, clip a pencil lead or paper clip.
  5. Put the tip of each pencil lead or paper clip in the solution. What do you observe?


Electrolysis involves passing a charge through a solution to drive a reaction that is typically nonspontaneous. When the charge is applied to water that contains an electrolyte, the electrical energy causes the water to decompose into the pure diatomic gases of its constituents, hydrogen and oxygen. The balanced chemical equation for this reaction is:

Energy + 2H2O(l) ¿ 2H2(g) + O2(g)

The above reaction is a reduction–oxidation reaction. Hydrogen is reduced in the reaction (from +1 to 0), and oxygen is oxidized (from –2 to 0). Here are the two half-reactions and the energy changes associated with them, expressed as standard electrode potentials (Eo):

Reduction     4H2O(l) + 4e– ¿ 2H2(g) + 4OH–(aq) Eored = 0.00 V
Oxidation     2H2O(l) ¿ O2(g) + 4H+(aq) + 4e– Eoox = –1.23 V

Adding these half-reaction equations gives the overall equation. Adding the standard reduction and oxidation potentials gives a negative value of –1.23 V. A negative value indicates that the reaction is not spontaneous and will occur only when energy is added to the system.

In the experiment, reduction of hydrogen occurs at the cathode (electrode connected to the negative battery terminal), while oxidation of oxygen occurs at the anode (electrode connected to the positive battery terminal). An electrolyte is needed to carry the charge through the solution. Due to the low self-ionization of water, a separate, inert electrolyte is needed. The electrolyte in this activity is sodium sulfate, Na2SO4. This salt dissolves in water to form sodium and sulfate ions.

The addition of Bogen universal indicator solution lets students monitor the reaction’s progress. When the experiment begins, the solution is green because the sodium sulfate solution has a pH of 7 and Bogen's color at this pH is green. As the reaction progresses, violet develops at the cathode as hydroxide ions form and the solution becomes basic. At the anode, yellow or orange develops as hydrogen ions form and the solution becomes acidic. In addition, bubbles of hydrogen gas form at the cathode, and bubbles of oxygen gas form at the anode. If the hydrogen and oxygen gas were collected at each electrode, the volumes would be 2:1 respectively, as indicated by the molar ratios of the balanced equation.

Multiple colors appear as hydrogen ions and hydroxide ions spread through solution, creating pockets of locally different pH. Bogen universal indicator contains the acid-base indicators methyl red, bromthymol blue, and phenolphthalein to create a range of colors from pH 4 to 10. Refer to the color chart provided with the solution and to the table below to identify the pH values.

pH Color
4 Red
5 Orange
6 Yellow
7 Green
8 Light Blue
9 Dark Blue
10 Violet

Tip: Although most water-soluble salts make good electrolytes, chloride salts (e.g., sodium chloride and potassium chloride) generate small amounts of chlorine gas at the anode (because the oxidation of chloride ion is energetically competitive with the oxidation of oxygen). The chlorine gas may corrode the electrode. For best results, choose inert electrolytes such as sulfate salts.

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