CONTENT IN A
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The elements fluorine, chlorine, bromine, iodine, and astatine form a group or family in the Periodic Table known as the halogens (salt formers). As a group, these elements are unique in several ways. All three states of matter are represented by halogens at room temperature and one atmosphere pressure: fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids. As elements, all exist as diatomic molecules. Because of their high reactivity, halogens are not found free in nature, but are usually produced from halide salts.

When considering the halogens, particularly during laboratory activities and demonstrations, one usually concentrates on chlorine, bromine, and iodine. There is a good reason for this. All known isotopes of naturally occurring astatine are radioactive with relatively short half-lives. It is estimated that only a few milligrams of astatine exist at the earth’s surface at any particular time. Fluorine, on the other hand, has a single stable isotope and is the second most abundant of the halogens at the earth’s surface—if the hydrosphere and lithosphere are combined; chlorine is most abundant.  But fluorine is the most reactive of all elements, reacting with all other elements except the noble gases helium, neon, and argon. Fluorine, which is prepared by electrolysis, cannot be used for reactions in water solution because it oxidizes water, but can be stored in stainless steel vessels because protective fluoride coatings are formed.

The most obvious chemical property of halogen elements is their ability to act as oxidizing agents. The relative oxidizing ability of halogens corresponds to their order in the Periodic Table, F₂ > Cl₂ > Br₂ > I₂ . This trend can be amply shown by a student laboratory activity and/or teacher demonstration and correlated to trends in atomic size, electron affinity, and electronegativity. Commercially,Cl₂ is obtained by the electrolysis of NaCl solution; Br₂ and I₂ are, in turn, produced by Cl₂ oxidation of Br^– and I ^– , respectively. For example:

Cl₂(g) + 2Br^–(aq) ---> Br₂(aq) + 2Cl^–(aq)

A major use of Cl₂ is in water treatment, where its action on pathogens has led to enormous improvement in the public’s general health. All halogens are extensively used in the synthesis of organic compounds where properties may be tailored by replacing hydrogen atoms with halogen atoms. Of particular note are Freons and halogen-containing polymers such as polyvinyl chloride (PVC) and fluorine-containing Teflons. Many of these halogenated hydrocarbons have been found to be highly toxic, potential carcinogens or, as in the case of Freons, environmentally unsound and so have been or are being removed from general use.

All halogens, with the exception of fluorine, exhibit oxidation states of –1, 0, +1, +3, +5, and +7 (F exhibits only –1 and 0). The most common state is the halide or –1 state; crystalline metallic halides are common compounds. The iodide ion is a reasonably good reducing agent and can cause the reduction of metallic ions such as Fe³⁺ to Fe²⁺ and Cu²⁺ to Cu⁺ . A clear distinction must be made between elemental halogens and halide ions during the study of “halogens.”

Oxoanions and oxyacids of halogens are themselves potent oxidizing agents that, in the form of “bleaches,” “cleaners,” or pool additives, find their way into consumer products.  Common liquid “chlorine” bleaches are not chlorine water solutions, but rather sodium or calcium hypochlorite solutions; the hypochlorite ion does the bleaching. It should be emphasized that chlorine bleaches and ammonia-containing cleansers should never be mixed, since they react to form very toxic chloramine (NH ₂ Cl).

The importance of halogens to human history is detailed later, as are links between halogen chemistry and issues of importance to our everyday lives.

This descriptive module can be used in the first year chemistry in a variety of places, either as a free-standing module taking cognizance of the related concepts and skills noted below, or selected parts of the module can be used in two or more parts of course(s) in multi-year chemistry/science curricula. For example, much of the material can be integrated into study of oxidation-reduction. Also, chemical and physical properties of halogens can be used as a part of a unit of study on the Periodic Table.

1. Halogens comprise a highly reactive family of elements that do not exist in the free state in nature but can be prepared electrochemically or chemically from halide melts or halide solutions (see Laboratory Activity 1 and Demonstration 1).

2. Free halogen atoms (seven outer valence electrons) are one electron short of an eight-electron noble gas configuration. The atoms dimerize (2X ---> X 2 ) or act as very strong oxidants (X + e ^– ---> X ^– ) to complete the octet of electrons.

3. Diatomic halogens are fairly strong oxidants (X ₂ + 2e^– ---> 2X^– ; with F₂ > Cl₂ > Br₂ > I₂ > At₂ ). The reverse reaction can be obtained electrochemically or by stronger oxidizing agents. These reactive elements have characteristic properties, e.g., at room temperature they exist as gases (fluorine and chlorine), a liquid (bromine), and solids (iodine and astatine) with colors ranging from light and medium yellow-greens for fluorine and chlorine through deep red for bromine to almost black with a metallic lusterforiodine—although its vapor is a rich violet color (see Laboratory Activity 2 and Demonstrations 1b, 4).

4. Halide ions (X ^– ) have noble gas electron configurations and are found in nature as stable crystalline solids (salts), that, with few exceptions, are soluble in water, are colorless unless colored by the cation component, and are used to form strong acids (except fluoride).

5. Halogens, except fluorine, form oxoanions (XO_n^– ) that are highly reactive and can serve as bleaching agents and disinfectants similar to the elements themselves. Oxyhalogen ions provide another way for electron octet formation.  These anions formally contain halogen atoms in positive oxidation states, which have even fewer valence electrons per halogen atom than in the free element. This is consistent with the fact that they are even stronger oxidants than are halogens.

6. Organic halogen compounds are rare in nature, although many useful and stable ones have been synthesized. They can be solids, liquids, or gases; some form very stable polymers. Halogens in these organics are covalently bonded
to carbon, making organic halogen compounds inert, except in the presence of irradiation, heat, or catalysts. Unfortunately, many are highly toxic, and the stability of others is causing environmental problems.

7. Halogen elements provide trends that can be extended elsewhere in the Periodic Table. These include, but are not limited to, trends in oxidizing strength, atomic and ionic sizes, completion of the octet of electrons by sharing electrons or forming anions or oxoanions, and trends in electronegativities (see Laboratory Activity 2 and Demonstration 2).

1. Redox (Oxidation-reduction concepts can be taught concurrently with this descriptive module.)

2. Acids and bases

3. Atomic structure, including ions

4. Bonding (electron-pair sharing and ionic aggregates in salts—the gas-to-solid trend can also be used in considering van der Waals or London dispersion forces)

5. Formula writing (nomenclature)

6. Equation writing and balancing

7. Electrochemistry (specifically, electron donation at the cathode and electron removal at the anode)

8. Solutions (one species dispersed in another as well as salts dissolved as solvated ions)

9. Periodicity

1. Ability to read carefully and follow instructions precisely.

2. Dexterity to handle and shake potentially hazardous solutions in small test tubes and to handle a hot test-tube.

3. Arithmetic skills necessary for writing formulas and balancing equations.

4. Stoichiometry problems and quantitative laboratory activities should be added if this material is deferred until late in the year. (Representative sources for such quantitative material are listed at the close of this module.)

Following their study of halogens, students should be able to:

1. identify halogens as a distinct family of elements.

2. describe the highly reactive nature of halogens.

3. describe how to prepare halogens electrochemically or chemically from halide melts of halide solutions.

4. identify the electron configuration of a halogen.

5. use electron configurations of halogen elements to explain why they dimerize and act as oxidants.

6. list halogens in order of increasing strength as oxidizing agents.

7. identify the characteristic properties of halogens.

8. recognize that halogens, except for fluorine, form oxyanions that are highly reactive and serve as bleaching agents and disinfectants.

9. demonstrate, with a Lewis diagram, how a halogen can complete its octet by forming an oxoanion in which the halogen has a positive oxidation state.

10. identify halogen-containing organic compounds.

11. explain, using the fact that halo-organic compounds are stable and thus desirable in applications, why halo-organic compounds pose an environmental problem.

12. use halogens as the basis for a confirmation of the periodic law.