WARNING! This is not a beginner procedure. Bromine can kill you. It can destroy your lungs. It can literally dissolve your flesh. Handle with care.
If you choose to attempt any of the experiments or procedures described on this site, you do so entirely at your own risk. The uses for the glassware and lab supplies shown in this article are not claimed to be the proper or safe uses. They are in fact rather dangerous uses which are performed by a professional under carefully-controlled circumstances.
While there's a certain allure to having around for no reason a viciously corrosive, red, fuming liquid that emits choking fumes with a chlorine-like odor, it turns out that our lab occasionally has use for bromine water (e.g., reactions that decolorize bromine water as a diagnostic test; bromine as oxidizer; research for articles such as this one). Instead of purchasing bromine in the pure state and being stuck with more than we'd probably use in a year, the author did the following procedure to make a small amount for experiments-- and to demonstrate the phenomenon to our readers.
Some readers have correctly noted that there would have been better methods for producing dry bromine. The goal here wasn't so much the making of bromine as it was the making of bromine by electrolysis. Nevertheless, distillation of the dark liquid taken from the anolyte chamber would have yielded bromine of roughly 98% Br2; we may do this experiment in a future article just to show that the concept works. Please be advised that distilling bromine, however, mobilizes the toxic fumes en masse and must be conducted in a working fume cupboard. Such work also has to be done with all-glass or PTFE apparatus; bromine attacks cork and rubber, and at elevated temperatures this can be violent.
This article is not a recommendation that the reader attempt the experiment; it is simply an account of what the author did. If you don't have a fume hood handy, don't work with bromine. A small whiff of it will leave your lungs inflamed for hours and could give you an asthma attack; a serious inhalation can cause your lungs to secrete fluid until you drown.
Quoting the MSDS for the compound:
10 ppm Immediately Dangerous to Life or Health.1000 ppm is only a 0.1% concentration of bromine in air.
Of course, would-be chemophobes should note that bromine is not really any more toxic than chlorine gas. Both have an IDLH concentration of 10 ppm, and both must be treated with respect. (Incidentally, ozone gas is more toxic than either, having an IDLH of only 5 ppm!)
Another safety concern in this experiment is the generation of hydrogen gas during electrolysis. Not only is hydrogen a fire and explosion hazard in itself, but bromine and hydrogen can also react to cause ignition or explosion. This is usually a problem only at high temperatures, but keep them separate just to be safe. Most of the time the bromine is under water, and the hydrogen dissipates, getting sucked up into the fume hood. However, it's conceivable that an accident could happen if the gases were to build up. It's just one more reason why this reaction should not be carried out by beginners.
Safety goggles and face shield;
Three layers of vinyl gloves on each hand*
Heavy lab coat and rubberized apron;
Electrolysis cell (as shown & described below)
Laboratory power supply
* Bromine will destroy vinyl with prolonged exposure. The author wore the gloves to give temporary protection from accidental contact. The idea was to provide enough time to get to the sink, where bromine could be washed off with copious water and the gloves removed.
Methods & Observations:
A plastic tumbler was obtained; a piece of polycarbonate sheet was then carefully scored, broken, and sanded to make it fit the tumbler. A hole was melted in the center of this polycarbonate sheet by pressing a heated copper tube through it.
Instead of gluing the partition in with industrial adhesive, it seemed better to wedge the partition in with paper. After all, the partition wasn't supposed to exclude ions completely. Wedging it in allowed removal if necessary.
The hole in the center of the partition was stuffed with three layers of paper towel (1-ply, generic brand) before wedging the partition into the tumbler. When wet with electrolyte, this paper plug (and the paper around the edges) would be analogous to a "salt bridge". Note: Pure bromine should not be allowed to contact cellulose or other combustibles, but it was not deemed a problem in the submerged environment of this experiment.
The anode was graphite; the cathode was copper. Both electrodes were held in place with wooden clothespins. These are useful in the laboratory; just don't re-use them for clothes or they will cause mysterious holes in fabric, not to mention possibly poisoning you.
A saturated solution of KBr in distilled water was prepared at room temperature. A moderately strong solution of KOH was also prepared in another beaker.
The KBr solution was poured into the anode chamber at precisely the same time the KOH was poured into the cathode chamber. The reason behind pouring them at the same time was that there wouldn't be an imbalance of pressure that could dislodge the paper barrier from the center hole.
A 12-volt power supply (current limiting) was hooked up. Immediately, the cathode chamber began to fizz with hydrogen bubbles. The current draw was between 3 and 4 amps. Within less than a minute, red-brown wisps of bromine began to diffuse away from the graphite anode.
Below: Most of the bromine formed directly in line with the center hole / membrane, indicating the cell was working as designed. Note the bromine sliding off the graphite anode and down to the bottom of the cell. The dark material in the bromine wisp appears to be graphite particles, carried away by erosion of the anode.
Unfortunately, the paper barrier around the edges of the partition wasn't tight enough. Bromine began to find its way into the catholyte chamber, especially at the bottom where there'd been a notch filed in the polycarbonate partition. The notch was intended to make room for the raised bump inside the bottom of the plastic tumbler; evidently it made too much room.
What's very interesting is that there was an apparent phase separation between the Br2-KBr-H2O and the KOH-H2O solutions. It's readily visible in the following photo:
What you're seeing is not a low level of aqueous solution; it's two distinct layers. The bottom layer itself then has two regions of its own: a bromine-rich layer (orange) and a bromine-poor layer (pale yellow, almost colorless). The entirely colorless region above is the KOH solution. The solutions are not spontaneously mixing.
It was apparent that bromine wasn't mixing completely with the water. Because of bromine's density, it poured off the graphite anode as it formed and sank to the bottom of the cell.
The electrolysis was stopped for a moment, since an idea dawned on the writer. The graphite anode was removed and its end was placed into a 10-mL micro beaker. This anode-beaker assembly was carefully lowered into the anolyte and the graphite anode was then re-clamped. When electrolysis resumed, any bromine formed would sink into the micro beaker instead of mixing with the rest of the cell contents. It worked:
At this point it was easy enough to pipette off the bromine and save it. It is not certain how much water the bromine had acquired by the time it sank into the beaker, but it was obviously concentrated enough to emit dense, red-brown fumes of Br2 when the liquid was placed into a micro flask:
Above: Pure bromine is very dark; this sample contains significant water. Colorimetric analysis would be one way to determine concentration of Br2 relative to water in this sample, though it would be necessary to have a bromine standard (i.e., known concentration) when making the curve.The experiment was run just long enough to produce only a small amount of bromine. If this procedure is adapted for [advanced] classroom demonstration, make only enough bromine to demonstrate the principle. Be careful! Of course, you must set up the apparatus under a working fume hood. You don't want too much bromine hanging around. It is best to make only the amount needed for the experiments at hand
The electrolysis of bromides proceeds according to the half-reactions:
2Br- (aq.) <---> Br2 (aq.) + 2e- -1.0873 volts
2Br- (aq.) <---> Br2 (liquid) + 2e- -1.066 volts
According to the voltages (Handbook of Chemistry and Physics, 1989), we might guess that the formation of liquid Br2 is slightly favored over aqueous Br2, since the former's oxidation potential is slightly less negative. One might guess bromine to have somewhat more affinity for other bromine molecules than for water, at least at room temperature.
Of course, at the cathode there is reduction of water to hydrogen gas:
2H2O + 2e- <---> H2 (gas) + 2OH- (aq.) -0.8277 volts,
though this doesn't bear directly on the formation of bromine. The KOH itself is just an electrolyte; sulfate could have been used just as easily. However, the author thought it might be handy to use the KOH in case he might later want to form potassium hypobromite, with a view to performing some bleaching experiments (and also because BrO- and HBrO are just plain interesting).
Through this demonstration we've not only explored the properties of bromine, but we've also uncovered some interesting questions worthy of additional research. In no particular order:
1.) To what extent does aqueous bromine react with the cellulose of the paper barrier? What are the products? Do they dissolve in the solution or remain bound to the paper? (hint: can dilute, aqueous bromine cleave polysaccharide bonds or not? What groups on the sugar molecules are attacked by bromine?)
2.) Did we correctly determine all the half-reactions that are actually taking place in our 'bromine cell'?
As it turns out, bromine forms the tribromide ion (Br3-) in the presence of bromide. The species are in equilibrium; it is not a one-way reaction.
Br2 (aq) + Br- (aq) <-------> Br3- (aq)
This is analogous to the behavior of iodine in a solution of iodide.
There will undoubtedly be some tribromide in our bromine water. Distillation or solvent extraction could separate this.
Some of the tribromide also reacts with bromine to form pentrabromide ion, Br5- (Popov, 1967). The equilibrium constant for this is on the order of 1/10th that of tribromide (1967).
Aqueous bromine also reacts to some extent with water, forming small amount of hypobromous acid (HOBr, sometimes written as HBrO). It appears the chemistry of bromine in solution is fairly complex, especially when some bromide is present.
3.) How might we get bromine from a solution of seawater containing a majority of chlorides but only a small percentage of bromides? Is electrolysis practical for this? Based on the oxidation potentials, which comes off first-- chlorine or bromine?
4.) What is the vapor pressure curve for bromine over its liquid range? How about for water solutions? Will cooling it cause the fuming to stop? We can find from most any reference that the vapor pressure of Br2 is a substantial 175 mm Hg at 20°C, for example.
5.) Does bromine form an actual hydrate whose formula can be determined? Cady (1985) suggests that it does, though it's not a whole number.
6.) What solvents are used in the lab to extract bromine, and to what degree does bromine react with each of them at room temperature? What are bromine's partition coefficients for common systems of water and hydrophobic solvent? CCl4 is a good solvent for bromine, as is CHCl3. Qualitatively speaking, we can see that most of the red-brown Br2 will end up in the CCl4 layer when the solvent is shaken together with water.
Elemental bromine is not a lingering environmental toxin in the manner of the heavy metals; it is so reactive that it will form bromides and organobromine compounds. The problem is, some of the latter may also be toxic (consider PBB's, polybrominated biphenyls). Bromine is vicious stuff during the time it exists as a free element; this must be regarded as a hazardous waste for disposal purposes (there are, however, chemical methods to recycle bromine back into relatively harmless bromides; we may cover the specifics in another article at some future time). Don't make bromine if you aren't equipped to handle corrosive, toxic vapors and liquids.
The laboratory preparation of bromine is a very interesting demonstration, in part because of the dramatic color change. There is nothing else quite like bromine (actually, NO2 is vaguely reminiscent, but that's not a liquid at room temp, and it's not much 'safer' than bromine anyway). Br2 has a disposition only a chemist could love. On the other hand, the safe and responsible handling of chemicals should be no cause for irrational fear. Bromine and the other halogens are necessary parts of our complex world. Learning more about them is something that can benefit us all.
A final, related note... Don't, by the way, ever attempt to prepare fluorine. F2 is so much more reactive than the other three halogens that its preparation requires inert vessels such as platinum. Fluorine reacts violently with air, water, oxygen, glass... almost everything! While there is some similarity between the preparations of bromine, chlorine, and sometimes iodine, this does NOT extend to fluorine.
Cady, G.H. Journal of Physical Chemistry 89: 3302 (1985)
CRC Handbook of Chemistry and Physics, 69th Edition. Boca Raton, Florida: CRC Press, 1989.
Popov, A.I. In Halogen Chemistry, Volume 1, ed. V. Guttman. New York: Academic Press, 1967.
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