How to Make Quick and Easy Disinfectant from Water and Salt

Table of Contents

Introduction: DIY disinfectant - can this be real?

While researching the best ways for maintaining hygiene in passive hydroponics, somebody suggested to try using hypochlorous acid ($\ce{HClO}$, or equally common $\ce{HOCl}$) as an alternative to e.g. hydrogen peroxide ($\ce{H2O2}$). $\ce{HClO}$ has a short shelf life (days rather than weeks) and therefore was not considered to be commercially viable until recently, however there is by now a lot of evidence that it has very good properties as a safe general-purpose disinfectant. Furthermore, and this makes it especially interesting to me, it can be easily made at home using nothing but a cheap USB dongle, water and salt.

The instructions that came with my 5$ device were a bit minimalistic and the whole thing looked somewhat sketchy and questionable at first. I was wondering whether this is really safe and whether it can work at all. According to the instructions, you just add around 2g of regular kitchen salt into 100 mL of clean water, let the device run for around 10 minutes, and then you supposedly have a hypochlorous acid solution.

It turns out that these instructions do work and will give you some solution with disinfecting properties, but without any adjustments and controls you do not really know what exactly and how much you actually produced.

In this post I am going to unpack what is going on and complete the picture. I have to confess that my knowledge of chemistry does not reach beyond first semester of university, and even this was a while ago, but luckily you do not need more than that to understand the whole system.

Theory: Some simple redox chemistry

Those USB-powered dongles are often sold as “generators” for either hypochlorous acid or bleach. They are actually simple electrolysis systems consisting of a cathode, an anode and some resistors limiting electron flow from anode to cathode, thus ensuring that the reaction stays within controlled and safe parameters. What they actually are doing is driving anodic chloride oxidation, producing chlorine at the anode surface, which immediately hydrolyzes in water. The main reactions happening in this electrolysis system are what this section is about.

By dissolving salt in water, we have $\ce{H2O + NaCl -> H2O + Na+ + Cl-}$, i.e. the ions from the dissolved $\ce{NaCl}$ can freely move in the water and participate in reactions. Assuming an idealized setup without any additives or contaminants in the water and salt, and no ions from corroding electrodes, then this is all we have to consider in our system. While the $\ce{Na+}$ ions will remain passive observers, you will see that the $\ce{Cl-}$ ions are running the show.

Once we start the USB electrolysis device, at the cathode water is gaining electrons and therefore is being reduced, producing some hydrogen gas in the process: $\ce{2H2O + 2e- -> H2 + OH-}$. The hydrogen quickly bubbles out of the system, leaving the $\ce{OH-}$ ions behind and thus making the solution more alkaline, i.e. the pH increases in the local area around the cathode.

At the anode there is more going on. We have some amount of oxygen being produced: $\ce{2 H2O -> O2 + 4 H+ + 4 e-}$. However, this reaction is not dominating (as long as chloride is available) and any created oxygen, like the hydrogen at the cathode, quickly forms bubbles and leaves the system. Until here all we have seen is simply the electrolysis of water. Now with the added salt, the main action at the anode is that chloride is being oxidized: $\ce{2 Cl- -> Cl2 + 2e-}$. The electrons are taken up by the anode, whereas the oxidized chloride ions form chlorine.

Recall that water is always in an equilibrium $\ce{2 H2O <-> H3O+ + OH- }$ and the pH value is just the negative decimal logarithm of the $\ce{H+}$ concentration in the water, which is the same as the concentration of hydronium cations $(\ce{H3O+})$. There are two reversible reactions that will happen now:

  • $\ce{ Cl2 + H2O <=> H+ + Cl- + HClO }$
  • $\ce{ HClO <=> H+ + ClO- }$

In which direction the balance is shifting depends primarily on the pH around the chlorine molecules:

Chloride species as a function of pH

As you can see, formation of $\ce{HClO}$ is strongly preferred at pH between 4 and 6. In more acidic environments, production of chlorine is favored, which then also will start to form bubbles and escape as $\ce{Cl2}$ gas (which is very toxic, so we definitely want to avoid this). In alkaline environments, the hypochlorite form $\ce{ClO-}$ is preferred.

In our system the anode and cathode are close together and the water is continuously mixed due to diffusion and also due to moving gas bubbles. Recall that at the cathode the pH is increasing due to the removal of hydrogen and production of additional $\ce{OH-}$ ions. Due to continuous mixing and diffusion, the chlorine produced at the anode will rapidly react into $\ce{HClO}$ and $\ce{ClO-}$ due to the overall non-acidic environment.

If you started with roughly neutral salty water, in the end you will be left with a somewhat alkaline solution, so at pH $\gt$ 8 you will end up with mostly $\ce{ClO-}$, which is just chlorine bleach. Sometimes, as we also have sodium ions ($\ce{Na+}$) around, for purely formal bookkeeping, they are grouped and accounted for together as sodium hypochlorite, $\ce{NaClO}$, which is the main ingredient of typical chlorine-based cleaning agents you can find in any supermarket.

At the same time, at pH $\lt$ 10, a part of the chloride ions will be bound in the form of $\ce{HClO}$, the hypochlorous acid and our desired disinfectant. As a disinfectant it is several times more effective than bleach at the same concentrations, because it is reacting faster and unlike the negatively charged bleach ions can permeate membranes and thus enter cells.

There are more side reactions and transient molecules that can be and are produced in negligible amounts, but there is one of special interest, because it is persistent: the chlorate ion $\ce{ClO3-}$. The two main reasons for chlorate formation in our system are:

  1. over-oxidation of $\ce{ClO-}$ at the anode during electrolysis
  2. decay of bleach during storage: $\ce{3 ClO- -> ClO3- + 2 Cl-}$

The first pathway is triggered by too high salt concentration (because it increases conductivity in the solution), but also by properties of the electrodes and the whole electrolysis setup (which we cannot influence directly and thus we have to trust the engineering of our USB device). The second pathway is favored at high pH and high $\ce{ClO-}$ concentrations and will happen slowly and naturally over time.

Chlorate, once formed, will stay and accumulate in the solution as an undesired byproduct. There are regulations for chlorate concentration in drinking water due to risks by chronic exposure, as it is known to interfere with iodine uptake. Chlorate formation during storage is usually more important than during short and controlled electrolysis runs, which is why self-made non-stabilized bleach should be used fresh and should not be stored longer than a few weeks.

$\ce{HClO}$ has an even shorter shelf life because it is chemically reactive and decomposes via multiple pathways, including oxygen evolution via $\ce{2HClO -> O2 + 2Cl- + 2H+}$, photochemical decay, and reactions with trace impurities. This is why $\ce{HClO}$ should be used fresh, ideally the same day.

Practice: Make Yourself Some Disinfectant (Or Bleach)

What You Need

  • USB dongle for NaCl electrolysis
  • free chlorine test strips (resolving between ~10-1000 ppm)
  • pH meter (or pH test strips, but they are much less precise)
  • sufficiently pure kitchen salt ($\ce{NaCl}$):
    • rock salt > sea salt (typically fewer undesired minerals than sea salt)
    • if possible, avoid additives like iodine or anticaking agents
  • sufficiently pure water ($\ce{H2O}$), such as:
    • distilled or reverse osmosis water (optimal)
    • water from an electric air dehumidifier (low-EC, so quite pure)[1]
    • tap water (works, but is suboptimal; at least boil + cool once)
  • for HClO: an acid to lower pH, such as:
  • for NaClO: a base to raise pH, such as:
    • baking soda (sodium bicarbonate, $\ce{NaHCO3}$)
    • washing soda (sodium carbonate, $\ce{Na2CO3}$)

Run the Electrolysis

Once you sourced the purest ingredients you can easily access, you can run the electrolysis following the simple instructions that come with the device, i.e. dissolve 1-2g of salt in 100mL water and run the device for at most 10-15 minutes.

Adding more salt will not automatically increase free available chlorine (FAC) yield at fixed device current limit and runtime, it will just increase corrosion of the electrodes and risk of chlorate formation due to increased conductivity and current density. Adding not enough salt on the other hand will make the current flow and thus the electrolysis unstable. For that reason you should stay in the recommended window for best results.

Longer running times provide diminishing returns while posing risks similar to high salt concentration, so again it is better to run the process only as long as necessary to reach the target concentration, which you can easily check using a little piece of the free chlorine test strips after turning off the device.

You can double the amount of salty water and the running time to get twice as much solution with about the same FAC concentration. This can quickly become impractical, unless you want to produce a lot of less concentrated solution. The FAC production rate is limited by the USB dongle, not the volume of salty water, so by scaling up the amount of salty water at a fixed running time you are simply diluting the result.

Now what is a good target concentration? Such a solution is not intended for long-term storage, so there is no point in concentrating it further to counteract decay. On the contrary, reactions such as chlorate formation are favored at higher concentrations, so it is better to keep them as high as necessary and as low as possible.

For most household and personal disinfection uses, concentrations of a few hundred ppm are more than sufficient, many applications require way less. Therefore roughly 200-300ppm of free chlorine is a good target[3] for the stock solution (which can be diluted further, depending on the use-case). With our equipment, it is anyway neither possible to have more precision than that, nor a way to reach significantly more without harming the device or risking other problems (such as chlorate formation).

The electrolysis process itself should be very unspectacular - all you should see ever happening is small or larger bubbles forming around the USB dongle, nothing more. The solution should remain fully clear the whole time and there should also be no strong smell at an arm’s length distance. If anything changes except for some bubbles coming from the electrodes, you should turn the device off immediately. In any case make sure to do all of this in a well-ventilated room.

Under no circumstances you should add any acid before the electrolysis is completed. Doing so you risk formation of chlorine gas, which is preferred by the system in an acidic environment. This is probably the most safety-critical caveat to keep in mind. Apart from this it is quite unlikely that something can go seriously wrong, assuming that you tried to avoid adding impurities to your solution that could facilitate unwanted side reactions and contaminate the result.

Pick Your Poison: Dialing in the pH

Use your pH meter to see where you stand with respect to acidity. Now is the point where you decide whether you want to make bleach or disinfectant. As discussed in the theory section above, in the solution we produced the bleach is most likely dominating.

For bleach ($\ce{ClO-}$) with maximized purity and shelf life, you can stabilize it in your solution by carefully adding some soda (both baking and washing soda will do) in small amounts and stirring until you reach a pH in the range between 9-10.5. Going above that is not helpful. On the contrary, it increases the likelihood of side reactions such as chlorate formation, which accelerates decay.

For disinfectant ($\ce{HClO}$), very carefully titrate an acid while stirring continuously and watching the pH meter until you reach a pH in the range between 5-6. This will maximize the $\ce{HClO}$ concentration in the solution. If you have no experience with this process, prepare a very diluted acid solution so that each drop changes the pH by at most ~0.1 in 100mL of pure water before titrating the real solution. Remember that bleach and acid should never be mixed in an uncontrolled way, as they can react violently and produce toxic chlorine gas ($\ce{Cl2}$).

Concerning the choice of the acid, phosphoric acid is the most predictable, stable and chemically robust in terms of side reactions in solution. If the HClO is not for immediate use, this is the only acceptable choice out of the three considered here. Both citric acid and acetic acid are organic acids that react with and consume some of the HClO over time, thereby reducing the effective HClO concentration and pH stability. Out of these two, citric acid is the more robust choice and is still very accessible in supermarkets as a cleaning agent.[4]

At the produced concentrations, both bleach and HClO will smell like a “swimming pool” when you hold your nose over it, but neither should be aggressive or irritating. HClO will smell slightly more “biting” at the same concentration. Pay attention to the smell just before and after titration to learn the difference. Any sudden sharp or irritating smell during acid titration is a warning sign indicating $\ce{Cl2}$ formation.

Shelf Life and Storage

If you do not use it immediately or have produced too much, you can store the solution in a cool and dark place, preferably in a HDPE plastic container. Make sure that the container is not closed airtight to avoid pressure build-up due to slow gas formation.

Glass bottles are acceptable for short-term storage of low-concentration HClO, but are not ideal for bleach, especially at higher concentrations. If glass is used, avoid metal caps to prevent corrosion and contamination.

You can use a HClO solution within hours up to a few days at useful disinfecting strength and a bleach solution within weeks to a few months after production, depending on creation and storage conditions.

Technically, you can also store the solution as alkaline bleach and acidify on-demand to $\ce{HClO}$ for its better disinfectant properties. However, the bleach will be already in the process of decay and may contain increasing amounts of by-products such as chlorate. Such $\ce{HClO}$ should not be used for any sensitive applications with direct contact to humans, animals or plants.

Some Applications of HClO

There is some overlap where both $\ce{HClO}$ and $\ce{ClO-}$ can be used, but as a rule of thumb, you should strongly prefer $\ce{HClO}$ for any use-case with direct contact to living beings, probably you should only use bleach where $\ce{HClO}$ does not make more sense, and I struggle to come up with any realistic household scenario. For that reason, I will only consider $\ce{HClO}$ in this section.

My favorite HClO survey lists various examples of reported concentrations that have been successfully used in various hygienic and clinical use-cases. Quoting directly from the paper:

HOCl has been shown to inactivate a variety of viruses including coronaviruses in less than 1 minute. At a concentration of 200 ppm, HOCl is effective in decontaminating inert surfaces carrying noroviruses and other enteric viruses in a 1-minute contact time. When diluted 10-fold, HOCl solutions at 20 ppm were still effective in decontaminating environmental surfaces carrying viruses in a 10-minute contact time.

If it is good enough for hospitals, it certainly is good enough for use at home.[5] Beside the application as a surface disinfectant, the same paper further reports about successful use of HClO as wound disinfectant, hand disinfectant and even mouth rinse at concentrations between 50-200 ppm, and similar recommendations can be found in this document.

From this we can conclude that at those concentrations produced in our setup (i.e. around 200-300 ppm), there is no danger from skin contact with the HClO solution, if correctly produced. I would still wear gloves to avoid contact if I am not sure that it is below 200 ppm. I will probably try to optimize my electrolysis times to reliably land around 100-200 ppm without the need for control by test strips and further dilution.

For drinking water concentrations below 5 ppm are considered effective and safe, and WHO papers suggest that the water will taste horrible long before the allowed chlorine content is above safe limits:

[…] the guideline value is 5 mg/litre (rounded figure). It should be noted, however, that this value is conservative, as no adverse effect level was identified in this study. Most individuals are able to taste chlorine or its by-products (e.g. chloramines) at concentrations below 5 mg/litre, and some at levels as low as 0.3 mg/litre.

For applications in hydroponics, this site suggests using 200-500 ppm (soak + rinse) for cleaning the growing medium and 2-4 ppm continuously to prevent algae and fungi, assuming a re-circulating system. Another site states more generally:

While hypochlorous is safe to spray on plants and add to their water sources, it is important to monitor the chlorine levels. Chlorine concentrations of over 4 ppm can have detrimental effects on the plant system, and long-term exposure of greater than 2 ppm hypochlorous can also be toxic to some plants. We always recommend testing hypochlorous on a small sample of plants before applying it to your entire crop.

However, there is a paper indicating that especially in combination with ammonium ($\ce{NH4+}$) even concentrations of 0.5 ppm can be problematic, and especially in passive systems without circulation I’d rather stay below the stated concentrations.

So it turns out that typically used conservative concentrations for drinking water of 0.5-2 ppm are already on the higher end when considered for plant irrigation, among other things because nutrients for plants consist (unsurprisingly) very different substances than what you would typically find in water intended for human consumption.

Addition of concentrations of 2-10 ppm in water that is intended for later consumption by plants or humans should be considered to be a shock treatment in case of suspected contamination, assuming clear water, except for at most minor biofilm. Water that is visibly contaminated should be filtered or disposed anyway.

Conclusion

In the beginning I was asking - can this be real? The answer is clear: Yes, it is indeed, and I think that it’s amazing that you can produce such a useful chemical as $\ce{HClO}$ at home cheaply and easily!

This little project certainly sparked my interest in exploring more practical DIY chemistry at home. Maybe at some point I will try to replace the USB dongle with a system out of a battery pack, resistors and suitable titanium electrodes and also experiment with other simple and easily controllable reactions with useful products.

Until then, stay safe and healthy!

  1. Just in case this was not obvious: Note that water from the dehumidifier is fine for electrolysis, but not for drinking, unless you treat it appropriately (filter, boil, etc.) ↩

  2. If you do hydroponics, you probably have it already for your pH-down, and this is exactly what I use. ↩

  3. Note that the test strips cannot distinguish between $\ce{HClO}$ and $\ce{ClO-}$, so until here all instructions apply equally. ↩

  4. My favorite HClO survey actually suggests adding vinegar in its HClO recipe. This is certainly better than no acidification at all and is the most easily available choice, however it is chemically inferior to the other options. ↩

  5. Note that HOCl will kill many, but not all possible pathogens, for example this report states that tuberculosis is only reliably destroyed above 1000 ppm. ↩