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Chemistry: Halogens


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About this Lesson

  • Type: Video Tutorial
  • Length: 13:28
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 145 MB
  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Nonmetals (12 lessons, $19.80)
Chemistry: Group 17: The Halogens (2 lessons, $3.96)

This lesson was selected from a broader, comprehensive course, Chemistry, taught by Professor Harman, Professor Yee, and Professor Sammakia. This course and others are available from Thinkwell, Inc. The full course can be found at The full course covers atoms, molecules and ions, stoichiometry, reactions in aqueous solutions, gases, thermochemistry, Modern Atomic Theory, electron configurations, periodicity, chemical bonding, molecular geometry, bonding theory, oxidation-reduction reactions, condensed phases, solution properties, kinetics, acids and bases, organic reactions, thermodynamics, nuclear chemistry, metals, nonmetals, biochemistry, organic chemistry, and more.

Dean Harman is a professor of chemistry at the University of Virginia, where he has been honored with several teaching awards. He heads Harman Research Group, which specializes in the novel organic transformations made possible by electron-rich metal centers such as Os(II), RE(I), AND W(0). He holds a Ph.D. from Stanford University.

Gordon Yee is an associate professor of chemistry at Virginia Tech in Blacksburg, VA. He received his Ph.D. from Stanford University and completed postdoctoral work at DuPont. A widely published author, Professor Yee studies molecule-based magnetism.

Tarek Sammakia is a Professor of Chemistry at the University of Colorado at Boulder where he teaches organic chemistry to undergraduate and graduate students. He received his Ph.D. from Yale University and carried out postdoctoral research at Harvard University. He has received several national awards for his work in synthetic and mechanistic organic chemistry.

About this Author

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I am going to guess that the only time you have heard the word halogen, before you starting taking chemistry, was in the context of a light bulb. And that furthermore, you probably have absolutely no clue what it refers to. Well in this lecture, you are going to find out.
What we are talking about, as chemists, when we talk about the word, halogen, are the elements in an ultimate column in the periodic table. So the next to the right most column. We have fluorine, chlorine, bromine, iodine and astatine. And we are not going to talk much about astatine, because it is radioactive and there just isn't very much. As you know, bromine is a liquid. Iodine is a solid. Fluorine and chlorine are both gases. And as you also already know they are relatively electronegative elements. They have high ionization energies and high electron infinities.
Now, they all form diatomic molecules in nature when we make them. They don't occur in nature as diatomic molecules, they occur mostly as salts and then we use chemistry to turn them into diatomic molecules that are useful and convenient toward the chemistry that we want to do. And they are all oxidizing again because they are all very electronegative. In particular fluorine is very strongly oxidizing and it has the double effect of both being very electronegative and having a very weak homolytic bond cleavage entity or low bond enthalpy, And because it has this very low bond enthalpy, and the reason is it is very small so you have the two nuclei of the chlorine really close together and you get a lot of repulsion of the lone pair, so in other wards chlorine just doesn't want to be a diatomic molecule, but it also does not want to be monatomic because then it doesn't have a complete octet.
Anyway to work with fluorine, you have to use special containers like nickel and copper, which form fluorides which passivate the layer of the metal. For this reason, almost all of industry uses chlorine, rather than fluorine, even though fluorine might be a stronger oxidizer.
Now, how do we make these things? For instance, how do we make chlorine from chloride, which is how it appears exclusively in nature? And the answer is the chlor-alkali process is certainly one way where you take sodium chloride in solution and you pass a current through it and you make sodium hydroxide and hydrogen gas and chlorine and the chlorine comes off at the anode and hydrogen comes of as a cathode and then hydroxide is also prepared. What is this good for? Well, you get sodium hydroxide, which you can sell, you get chlorine, which you can sell and you get hydrogen, which you can sell. So this is a very important process for making these materials, but this is the way that you can make chlorine. You can take bromine and do the same sort of thing. So you can take bromide, bromide is also found in the ocean so this would be starting with sea water or something slightly purer than sea water. You can also do this with bromine, but more conveniently the way that bromine is made is by taking these solutions of sodium bromide or potassium bromide or whatever and just hitting them with chlorine. Chlorine is a stronger oxidizing agent than bromine or iodine and so by taking solutions of sodium bromide or sodium iodide, you can oxidize the iodide or the bromide up to iodine or bromine. And there isn't a whole lot of iodide hanging around, but it turns out that there are certain kelps that concentration iodide and so you go to sea, you collect the kelp and than if you burn it down to make ash, then that is a reasonably high concentration of iodide. So it is commercially viable to make iodide that way.
Now you can't do this process with fluorine and the is, is that fluorine is such a powerful oxidizing agent that it will actually oxidize water to O[2] and so you will have to use non-aqueous means to get to fluorine from various fluoride sources. For instance if you electrolyze molten potassium hydrogen fluoride you can make H[2], F[2] and KF and what this molten potassium hydrogen chloride is, is potassium cations and then FHF anions where you have a lot of hydrogen bonding and that is what holds this anion together.
Now some of the applications, let's just take them as we go down the periodic table, fluorine as a gas and it is really difficult to handle, but it is really important in the use of a lot of chemicals that are important commercially. One of them are the freons. This is freon[12], actually as a little aside, the number is a code word that is totally bizarre that I think Dupont came up with that allows you to name these things, because they didn't want anyone to know what they were talking about. Anyway that is an aside. I worked at Dupont briefly so that's why I heard this story. Anyway, the beauty of the things like fluoride, this is a chlorofluorocarbon is that they are very inert and being very inert, you can put them in your refrigerator and they serve as coolants in your refrigerator. The nice thing is that if they are very inert, they don't oxidize, they don't catch on fire. That's really important if you have got a lot of it in your refrigerator. But the problem is that they are so inert, that they actually get all the way up to the stratosphere, where the ozone layer is compromised because of the presence. These react with ultraviolet light to form ozone-destroying compounds. Any so that is why we are phasing out these freons, these chlorofluorocarbons. Fluorine is also important in Teflon, which is polytetraflouroethylene. Tetrafluoroethylene is this molecule, that is tetrafluoroethylene and when you polarize tetrafluoroethylene, you make polytetrafluoroethylene. A lot of pans and pans now are coated with Teflon; it is a very good non-stick surface.
And finally, a lot of fluorine is used in fluoridation of water. And when you add fluoride to water, people drink it and eat it and it gets incorporated into the hydroxyapatite that makes up the hard enamel of your teeth to form flourapatite. Fluorapatite is more resistant to decay from acids and so people get fewer cavities as a result of fluoridating water. Interestingly enough the way that they discovered that is, that there is some places where there is naturally occurring high concentrations of fluoride and people who lived there, just didn't get cavities and so they eventually did epidemiological studies and figured out that it was the fluoride was helping people to avoid cavities.
Now much more important industrially is chlorine. Chlorine can be disproportinated to hypochlorite. So this is a disproportionation where this is chlorine at zero oxidation state. That is minus one that is plus one. You dissolve chlorine in cold base and we will talk about why is has to be cold later on, and you form hypochloride, that is Clorox bleach. Sodium hypochloride is a solution of Clorox bleach, and it is used for cleaning your clothes. For getting clothes white, that aren't white. Chlorine is also dissolved in water, being a strong oxidizing agent, you can just bubble it into water and the oxidizing agent oxidizes bacteria and oxidizes whatever happens to be in the water, so basically makes the water fit for drinking. At least in Texas, we have been trying to do demonstrations where we take the Texas tap water and we add silver to try and make up a silver nitrate solution or something and it goes instantly cloudy. It's because the product of oxidation of a bacterium with chlorine gas is chloride and so there is a lot of chloride in the water in Texas. Another application of chlorine gas is in the production in titanium dioxide, TIO[2]. So the purification step in TIO[2] is, by the way TIO[2] is a white pigment. White paint has a bunch of TIO[2] in it. So in order to get it real pure, what you do is, you take crude TIO[2] and react it with carbon and chlorine to for titanium tetrachloride, in the lab we call this "tickle four." And this is a liquid, so it is distillable and then you react it with water to form titanium dioxide. Again, really important because paint, all paint, even colored paints, have white titanium dioxide as the thing that reflects light very strongly, so it provides the coverage.
And here we get now to halogen and halogen lamps. A regular light bulb has something like argon in addition to the tungsten wire. But in the case of a tungsten halogen light bulb, instead of having argon on the inside of the envelope, it has chlorine gas on the inside of the envelope, and the reason is that you have a tungsten wire and the tungsten wire gets hot and that is what gives off light, and it turns out that the hotter you run the wire, the more light that it gives off. That is just a physical fact. If you could run the wire hotter, you get more light out, so it would be a more efficient light bulb. That would be a good thing, but the problem is that when you run the wire hotter, it has a tendency to burn out. What happens is you're volatizing some of the tungsten off of it. As the wire gets thinner, the thinner the wire, the resistance goes up, so it locally gets hotter there and it gets hotter and hotter and eventually the tungsten filament burns out and your light bulb goes dead, so that is a problem. How could you run the tungsten halogen light bulb hotter and get more light? And the answer is you introduce a little bit of halogen gas into and when you introduce a little bit of halogen gas like chlorine, you react the tungsten that is boiling off of the filament with the chlorine to form tungsten hexochloride, that is WCL[6]. And so now there is some tungsten hexochloride floating around on the inside of your light bulb and what happens is that is preferentially decomposed at the hottest part in the filament. And so if you have part of your filament that is starting to get thin, just getting ready to burn out, the tungsten hexochloride that is in the rest of the light bulb, is going to go to that spot and decompose and sort of build it back up again, reinforce that part of the filament and so what it does, it allows you to run the whole thing faster, because there is this negative feedback whenever the filament starts to get thin, the tungsten hexochloride that is in the gas phase preferentially deposits more tungsten there to make it thicker again. So you can run the whole thing hotter. Turns out when you run it hotter, it is more efficient, it gives off more light and the light is whiter for reasons that won't get into, it is actually whiter as well, and so it is brighter and it is whiter and they last longer.
The commercial processes that use chlorine, the synthesis of vinyl chloride, this is CH2[2]HCL, it is used in making polyvinyl chloride, PVC, which is the white plastic tubing that you sometimes see underneath your sink for the drains. And similarly, dichloroethylene, which is the molecule, which is closely related to vinyl chloride, which is used as a dry-cleaning solvent.
Bromine has many fewer applications, but one of them is in the use of methyl bromide, which is a fumigant. It kills bugs. Grain is typically fumigated with methyl bromide. The problem is, is that methyl bromide is also an ozone depleter, just like chlorofluorocarbons are a depleter. But methyl bromide works so well, that there is a lot of resistance to getting rid of methyl bromide. I think that this brings up an interest point, that sometimes you have to consider cost benefit analysis. In chemistry, sometimes you have a molecule that works so well at what it does that you may have to live with the negative effects of its use, in the case of methyl bromide, that is definitely an example. Bromide is also used in silver bromide. In photographic film, when the film is unexposed, what it is, is it's a silver bromide in a colloidal gelatin and it is distributed out over the film and when you shine light on it, it makes some silver metal and then when it's processed it makes more silver metal and that is how you get the negative. So going from unexposed film to the negative involves making silver metal that deposits upon the film.
And finally, you probably have actually been exposed to iodine, so much as the solid, but as tincture of iodine, which is iodine dissolved in alcohol. When you go and give blood, when they are swabbing the area that they are going to take the blood out of, there is that orangish-red solution that is tincture of iodine. It is really good at killing germs because it is an oxidant. And it is just really convenient, much more convenient to dissolve a solid in ethanol, than a liquid like bromine or a gas-like chlorine. So it's just a convenient way of doing it. But it turns out that iodine is an essential element. You have to have it for function of your thyroid gland. There is a hormone called thyroxin that requires iodine in its synthesis, and typically a lot of people don't get enough iodine in their diet, so I don't know if you have ever noticed, but on salt boxes, you will typically see that there is some potassium iodide added as a necessary nutrient.
Finally silver iodide is used in high-speed photography so you may have noticed that films are getting faster and faster. Now you can shoot in much, much less light and that is because they have used silver iodide in the emulsion instead of silver bromide.
So halogens play a role in our lives. What we are going to see in the next unit is that there is more to this story yet involving things like the aqueous solution of halogens and halogen oxides.
The Nonmetals
Group 17: The Halogens
Halogens Page [1 of 3]

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