Key Questions

1. Where and in what form are halogens found in nature? [Halogens are found as halide salts, often in solution, either in oceans or in subterranean brines, or as solid deposits on land. Insoluble calcium fluoride, and soluble alkali metal chlorides, bromides, and iodides are most common. Iodine has been found as the iodate in some calcium minerals, but known deposits have been largely used up. Iodine from seaweed was also a common source at one time.]

2. Why aren’t these elements found in free form in nature? [Halogens are too reactive to be found in their free or elemental form in nature.]

3. What processes are employed to obtain halide ions from natural sources? [Calcium fluoride is dissolved in sulfuric acid to provide HF, which can be used as a source of metal fluoride salts by adding a metal hydroxide or carbonate to the HF. The soluble halide salts are typically dissolved and reprecipitated.]

4. What processes are employed to obtain free halogens from natural sources?  [Fluorine and chlorine are typically obtained through electrochemical oxidation of their halides, either in aqueous media or in a melt. Bromine and iodine are typically made from oxidation of bromide or iodide, respectively, by chlorine.]

5. What are some properties of halogen elements? [Fluorine and chlorine are yellow-green gases, bromine is a deep red liquid, and iodine is a very deep violet (essentially black) lustrous solid, etc. For further elaboration, see an inorganic textbook that emphasizes descriptive chemistry, such as Greenwood and Earnshaw, Chemistry of the Elements.]

6. How do chemical properties of halogen elements vary with their positions in the Periodic Table? [For example, as the family is surveyed from fluorine to iodine, ionization energies decrease, electron affinities maximize at chlorine, but electronegativity decreases from fluorine on, radii increase, electrode potentials and oxidizing powers decrease. Also see Question 5.]

7. What are some uses of halogens and halogen compounds in industry and in consumer products? [Uses have been summarized in Content in a Nutshell and in Links/Connections.]
 

1. HF(aq) is a very reactive, weak acid. This does, at first glance, seem to be a contradiction. It is not. The great attraction that fluorine has for most other atoms is the cause of HF(aq) being a weak acid. This high attraction of fluorine for most other atoms gives rise to a very strong H—F bond that is not broken when HF dissolves in water. The designation of HF as a weak acid does not mean that it is unreactive, only that it does not ionize completely in water. Indeed, periodicity is confirmed by the great attraction between H and F—an attraction that is greater than H—Cl > H—Br > H—I (see Demonstration 3).

2. Iodide is a reactant in Option B of the Halogen Reactivity demonstration (Demonstration 1), and is thus possibly not “stable.” This discrepant event is really nothing more than confirmation that Br ^– is more stable to oxidation than is I ^– and that I₂ is more stable to reduction than Br₂ —confirmation of periodic trends illustrated in the laboratory activities.
 

1. Halogens can be regarded as very “electron greedy” species. Their high electronegativities point out the fact that, in a covalent bond, they attract the shared pair of electrons aggressively. Indeed, if at all possible, a halide anion will form as the halogen seeks the bonding pair of electrons for itself. Family relationships involving younger siblings can provide a suitable analogy.

2. The Picture in the Mind diagram (Figure 9) of a square dance is another good analogy for halogen atom reactivity.

3. The use of an analogy of striving for perfection (complete energy level or stable octet) as the reason for the high “desire” of halogens for one more electron can be quite useful.
 

Thief: The chlorine atom attracts an electron from the bromide ion to complete its electronic octet (as the chloride ion) and, in so doing, leaves bromine one electron short.  It is analogous to a thief at a square dance who completes his square at the expense of someone else.

Other Activities

1. Home halogen hunt . This event can take many forms. Ask students to hunt at home (or at work) for places where halogens are used. This can be accomplished by reading labels or by following a list of possible uses compiled from a source such as an encyclopedia. You can process a wide variety of student responses as a class discussion. Point out that, in almost all cases, the reason for halogen use is its high reactivity or high stability depending upon which form is present.

2. Pervasiveness of human-introduced halogens in the environment . There are altogether too many examples of PCBs, insecticides, herbicides, plastics, and chlorofluorocarbons present in the environment. It is useful to note that the very reason these products are used is the same as the reason they plague the environment—their high stability in certain forms. A good background source is Environmental Chemistry by Lucy Pryde or Chemistry in Perspective by Mohrig and Child.

3. Chlorofluorocarbons and the ozone layer . This is another  contemporary case of persistence that lends itself to a good discussion—not only of the chemistry involved, but also of the possibility that environmental impact considerations should be placed ahead of personal convenience and comfort.

4. Use of fluoride as a preventive measure for tooth decay . Many
communities have placed this question on the ballot with passion evident on both sides of the issue. Here is a vehicle to inform your community regarding this issue. An excellent background source is a fluoridation article in the August 1, 1988 issue of Chemical and Engineering News.