Electrolysis is a chemical process that uses an electric current to drive a non-spontaneous chemical reaction. This method converts a simple saline solution, or brine, into sodium hypochlorite, the active ingredient in bleach, using a low-voltage electrical charge. The final product is a dilute form of bleach that can be used for household sanitation and cleaning.
The Underlying Chemical Reaction
The foundation of this process is the dissociation of salt (sodium chloride) in water into sodium and chloride ions, creating an electrically conductive electrolyte. When a direct electrical current is introduced through two submerged electrodes, distinct reactions occur at the terminals. At the positively charged anode, chloride ions lose electrons (oxidation) to form chlorine gas (\(\text{Cl}_2\)).
Simultaneously, at the negatively charged cathode, water molecules gain electrons (reduction), forming hydrogen gas (\(\text{H}_2\)) and hydroxide ions. Sodium ions remain in the solution, combining with the hydroxide ions to create sodium hydroxide. Because the cell is undivided, the chlorine gas produced at the anode dissolves and reacts immediately with the sodium hydroxide.
This secondary reaction forms the desired product, sodium hypochlorite (\(\text{NaOCl}\)), the active ingredient in bleach. The reaction is \(\text{Cl}_2 + 2\text{NaOH} \to \text{NaOCl} + \text{NaCl} + \text{H}_2\text{O}\). The temperature must be kept relatively low, ideally below \(60^\circ \text{C}\), because higher temperatures can cause the sodium hypochlorite to decompose into less effective products. The continuous flow of current drives these transformations.
Necessary Materials and Setup
The physical setup requires a power source, two electrodes, a container, and the chemical inputs. A low-voltage direct current (DC) power source, typically \(2\) to \(5\) volts, is needed. This low voltage is sufficient to drive the reaction while limiting unwanted side reactions.
The electrodes must be inert and corrosion-resistant to withstand the reactive environment. While commercial systems use expensive coated titanium, a common DIY choice is high-purity carbon or graphite rods, often salvaged from zinc-carbon batteries, which function acceptably for small-scale production. Avoid using stainless steel electrodes, as they can corrode and release toxic hexavalent chromium into the solution.
The electrolyte solution requires non-iodized salt, such as canning or pickling salt, because impurities in standard table salt can interfere with the reaction. Distilled water is recommended for mixing the brine to ensure purity and prevent mineral deposits from scaling the electrodes. The entire process takes place in a non-metallic container, which must be kept open to the air to allow for gas dissipation.
Step-by-Step Electrolysis Procedure
The first step is preparing the brine solution at an optimal concentration to ensure good electrical conductivity and efficient hypochlorite production. A common ratio for a homemade batch is a \(10\%\) salt solution (\(100\) grams of salt dissolved in \(1\) liter of water). Precise measurement is necessary to achieve a predictable outcome and ensure the solution is saturated enough for the reaction to favor chlorine production over oxygen.
Once the brine is prepared, the electrodes are submerged into the solution, ensuring they do not touch each other to prevent a short circuit. The electrodes are then connected to the DC power source, with the positive terminal connected to the anode and the negative to the cathode. Current should be applied at a consistent low voltage to begin the electrolysis.
The reaction is visibly confirmed by the steady production of bubbles at both electrode surfaces: chlorine gas at the anode and hydrogen gas at the cathode. The reaction must be run for a controlled period, typically between one to three hours, depending on the current applied and the desired concentration. A longer run time increases the concentration of sodium hypochlorite, but also increases the risk of product decomposition and side reactions.
After the current is turned off, the electrodes are carefully removed, and the solution should be allowed to cool before being transferred to a storage container. The final product is the entire resulting solution, containing sodium hypochlorite and remaining salt. No complex filtration is typically required, but the solution should be handled with caution before use.
Safety, Storage, and Product Concentration
The electrolysis of brine generates two hazards: toxic chlorine gas (\(\text{Cl}_2\)) and flammable hydrogen gas (\(\text{H}_2\)). Due to the continuous production of these gases, the process must only be conducted in a location with exceptional ventilation, such as outdoors or under a certified chemical fume hood. The container should never be sealed during operation, as this could allow the explosive hydrogen gas to accumulate.
Protective equipment, including chemical-resistant gloves and safety goggles, must be worn to prevent contact with the corrosive electrolyte and the final alkaline product. The resulting hypochlorite solution is unstable and degrades quickly when exposed to light and heat. For safe storage, the homemade bleach must be immediately transferred to an opaque, sealed container and kept in a cool, dark place, away from children and pets.
The concentration of the final homemade solution is significantly lower than commercial household bleach, which is typically \(5\%\) to \(6\%\) sodium hypochlorite. Small-scale electrolytic generators typically produce a solution concentration in the range of \(0.05\%\) to \(0.8\%\). This makes it suitable for general surface cleaning and sanitizing rather than heavy-duty laundry bleaching. Never mix this homemade hypochlorite solution with any other cleaning agents, particularly ammonia or acidic compounds, as this can release toxic fumes.