Volume 3

Natural Occurrence and Physical Properties • General Chemical Properties of Halogens • Chlorine • Hydrochloric Acid • Is HCl the Strongest Acid of All Hydrogen Halides? • Oxygen Derivatives of Chlorine • Bromine and Iodine • Fluorine • Silver Halides • Halogen Compounds Around Us • Exercises

3.3.1. Natural Occurrence and Physical Properties. In the two previous sections (3.1 and 3.2) we learned about the most metallic elements, the alkali metals (Group 1) and the alkaline earth metals (Group 2). We are now moving across the periodic table all the way to the other extreme, the halogens (Group 17) that represent the most nonmetallic elements.

All five halogens, F, Cl, Br, I, and At, occur in nature. The heaviest halogen, astatine (At), is the rarest element in the Earth's crust. Astatine is radioactive, constantly emerging from other radioactive elements and quickly decaying. Fluorine and chlorine are abundant on Earth. Bromine is less so. Iodine is quite rare, but not considered to be an exotic element.

All naturally occurring halogens have the oxidation state of -1, existing mostly in the form of the highly stable anions, F-, Cl-, Br-, and I-. While the simple substances (F2, Cl2, Br2, and I2) are too reactive to occur in nature, one exception does exist. This striking isolated example is the mineral antozonite, a rare variety of fluorite (CaF2) that contains fluorine gas, F2, confined in small pores inside its crystals. Antozonite (Figure 3-19) is sometimes called "stinky rock" due to the pungent smell it gives off when crushed. The odor is from the F2 released from the pockets inside the mineral on grinding. Astonishingly, it is fluorine, the most reactive halogen and the most reactive simple substance ever that occurs naturally!!
Figure 3-19. Antozonite, the mineral that confines F2 in the pockets inside its crystals (source).

Selected physical properties of the halogens as elements and as simple substances are presented in Tables 3 and 4, respectively.

Table 3. Selected properties of the halogens as elements.
Table 4. Selected physical properties of the halogens as simple substances.
Note that a halide anion is larger in size than the parent atom (Table 3). Adding an electron to a halogen atom enhances the attractive electrostatic force between the electrons and the positively charged nucleus, which one would have expected to result in contraction. Yet the change we observe is the opposite. Why? The effect is explained by electrostatic repulsion between the electrons in the larger number. This repulsion overpowers the attraction, thereby entailing the observed increase in size.

At standard temperature and pressure, fluorine and chlorine are gases. Iodine is a solid that sublimes easily to produce deep purple vapors. It is often stated that iodine sublimes before it can be melted and that liquid iodine can be observed only if I2 is heated in a sealed tube. This is incorrect. Bromine is a dense yet volatile dark red-brown liquid, the only liquid nonmetal and one of the only two simple substances that are liquids at standard temperature and pressure. (What is the second one? Mercury.)

3.3.2. General Chemical Properties of Halogens. The electron configuration of the halogens shows that a halogen atom needs just one electron to attain the stable electron configuration of a noble gas.

F 1s2 2s2 2p5
Cl 1s2 2s2 2p6 3s2 3p5
Br 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
I 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5

There are two possibilities for a halogen atom to get an octet of valence electrons. One is to receive an electron from another atom, ion, or molecule, as exemplified by the reaction of a fluorine atom with a cesium atom (Figure 3-20, A). This video provides a unique opportunity to watch a demonstration of the spectacular reaction between fluorine, the most reactive nonmetal, with cesium, the most reactive metal.
Figure 3-20. A halogen atom attains the stable octet of electrons by either accepting an electron from an electron donor (A) or forming a covalent bond with an identical or different atom (B).

The other option for a halogen atom to stabilize is to form a covalent bond to another atom by creating a shared electron pair (Figure 3-20 B). The bonding partner can be an atom of the same halogen or a different element such as hydrogen. These considerations account for the fact that halogen molecules are diatomic.

The halogens react with a broad variety of substances, the order of reactivity being F2 > Cl2 > Br2 > I2. Of any two different halogen atoms, the one that is lighter has fewer electrons and is therefore "hungrier" for an extra one. Moreover, since a lighter halogen atom is smaller in size, its outermost shell provides a stronger electrostatic attraction to the incoming electron relative to a halogen atom that is heavier and, consequently, larger. The weakest electron acceptor of the four halogens is iodine. Fluorine is the most powerful acceptor of electrons not only among the halogens but also among all elements.

3.3.3. Chlorine. Chlorine, the most broadly used halogen, is made in industry by electrolysis of aqueous NaCl or KCl (chloralkali process, see: 3.1.4 and Figure 3-6). In the laboratory, Cl2 is usually made by reacting HCl with an oxidant that is strong enough to oxidize the chloride ion, 2 Cl- - 2 e- = Cl2. Suitable oxidants include MnO2, KClO3 (Figure 3-21), and KMnO4.
Figure 3-21. Two laboratory methods to make Cl2.

Watch a demonstration of the reaction with KClO3 here. Unless you understand Vietnamese, never mind the narrative to the video. Just watch the mixture turn yellow on addition of HCl to Berthollet's salt due to the formation of Cl2. My Vietnamese is nonexistent, but I know well the setup used for the experiment. The Cl2 produced from KClO3 and HCl in the reaction flask is passed through two bubblers on the way to the receiving flask. One bubbler contains a solution of NaCl, apparently to trap the HCl vapors. I do not know why they used aqueous NaCl, as just water would do the job. The second bubbler is filled with concentrated H2SO4 to remove moisture from the chlorine gas produced in the reaction. (You are expected to remember that concentrated sulfuric acid avidly absorbs water and is therefore widely used as a powerful dehydrating agent.)

Both reactions in Figure 3-21 are redox reactions where the oxidant (MnO2 or KClO3) oxidizes the Cl-. The oxidation state of the Mn changes from +4 to +2 in the reaction. The reaction of KClO3 with HCl was considered earlier in our course (Volume 2). In this comproportionation reaction, the Cl is both reduced (from +5 to 0) and oxidized (from -1 to 0). That is why there are two zeros, one blue and one red, above the Cl2 formula in the equation in Figure 3-21. The red zero indicates the reduction of the Cl in KClO3, whose oxidation state (+5) is in red. The blue zero is to point to the oxidation of the Cl- (from HCl), whose oxidation state (-1) is also in blue.

Chlorine (but not chloride!) is highly reactive and toxic. It is chlorine gas that was the first killing chemical warfare agent ever used (World War 1, April-May 1915). The high reactivity of Cl2 stems from its strong oxidizing ability. Chlorine readily oxidizes not only the most strongly electron-donating alkali and alkaline earth metals, but also less reactive iron, zinc, copper, and even highly chemically inert gold. Examples of reactions of Cl2 with some metals are presented in Figure 3-22. Note that metals that can form compounds in different oxidation states are converted to the higher oxidation state chloride in the reaction with Cl2. For example, iron reacts with Cl2 to give FeCl3, not FeCl2. Likewise, CuCl2, not CuCl, is formed in the reaction of metallic copper with Cl2.
Figure 3-22. Chlorine readily reacts with various metals to produce metal chlorides. For metals that can exist in two oxidation states, such as Fe and Cu, it is the higher oxidation state chloride that is formed in the reaction with Cl2.

Here you can watch good video demonstrations of some reactions of Cl2. (I, for one, prefer to watch this video with the sound off, but this is a matter of taste.) One of the experiments in this video is the reaction of chlorine with phosphorus. Chlorine does react with many nonmetals that have a lower electronegativity, such as phosphorus. It is correctly stated in the video clip that the reaction of Cl2 with P furnishes a mixture of PCl3 and PCl5. First, PCl3 is formed, which then reacts with extra Cl2 to give PCl5. While the formation of PCl3 is an irreversible process, the ongoing reaction with Cl2 is reversible (Figure 3-23).
Figure 3-23. Reaction of phosphorus with chlorine.

Chlorine readily displaces less electronegative bromine and iodine from bromide and iodide salts (Figure 3-24). Obviously, however, chlorine cannot and does not displace more electronegative fluorine from a fluoride salt.
Figure 3-24. Chlorine displaces bromine and iodine from bromide and chloride salts.

3.3.4. Hydrochloric Acid. Pure HCl is a gas that is easily soluble in water. Aqueous solutions of HCl are called hydrochloric (or muriatic) acid. Produced on a very large scale, HCl is the most common mineral acid and one of the cheapest commodity chemicals. Being a strong acid, HCl readily reacts with metals (Figure 3-25) as well as with bases and basic and amphoteric oxides (Figure 3-26). Note that the reaction of HCl with iron gives rise to FeCl2, not FeCl3 (Figure 3-25). This contrasts with the previously considered reaction of Fe with Cl2 that produces FeCl3, not FeCl2 (Figure 3-22).
Figure 3-25. Reactions of HCl with selected metals.
Figure 3-26. Reactions of HCl with basic and amphoteric hydroxides and oxides.

Now let us recall that HCl is a strong electrolyte and convert the regular equations presented in Figure 3-26 to ionic ones (Figure 3-27). The ionic equations show that the role of HCl in all of these transformations is to provide the hydrogen ion (H+) for the reaction, whereas the anion, Cl-, is just a spectator. Therefore, any acid of reasonable strength should be able to effect similar transformations. This is indeed the case. Other strong acids such as H2SO4, HBr, and HNO3 react with basic and amphoteric oxides and hydroxides similarly.
Figure 3-27. Converting the regular molecular equations shown in Figure 23 to ionic ones.

3.3.5. Is HCl the Strongest Acid of All Hydrogen Halides? Although aqueous HCl is one of the strongest mineral acids, aqueous HBr and HI are even stronger. In contrast, aqueous HF is a weak acid. The order of acidity in water HF < HCl < HBr < HI can be rationalized in terms of the difference in size of the halogen atoms and, as a consequence, the hydrogen-halogen bond distance r (Figure 3-28). The longer the H-X bond (X = F, Cl, Br, I), the weaker it is and, as a consequence, the more inclined the HX molecule is to dissociate. Note that, in general, the ability (or willingness) of an acid to dissociate is a very complex issue. The explanation provided for the order of acidity of hydrogen halides in this course is a simplified one.
Figure 3-28. A rationale for the order of acidity HF < HCl < HBr < HI.

3.3.6. Oxygen Derivatives of Chlorine. Chlorine compounds containing a chlorine-oxygen bond are not uncommon. One such compound is a broadly used household chemical, laundry bleach, which is used for cleaning and stain removal. Laundry bleach is a 5-8% solution of sodium hypochlorite (NaOCl), a salt of hypochlorous acid, HClO. Besides HClO, three more chlorine oxoacids (oxygen-containing acids) are known (Table 5).

Table 5. Oxoacids of chlorine.
Pure oxoacids of chlorine are unstable and decompose violently. Their solutions, however, are much more stable and so are some of their isolated salts (see below). The strength of chlorine oxoacids as electrolytes varies in a broad range. While both HClO and HClO2 are weak acids, HClO3 is a strong acid (on a par with HNO3), and HClO4 is exceptionally strong, even stronger than H2SO4, HCl, and HNO3.

In all of the oxoacids of chlorine, the Cl atom has a positive oxidation state because it is bonded to one or more O atoms that are more electronegative than Cl. Consequently, all oxoadcids of chlorine and their salts are oxidizers.

Hypochlorous acid is slowly formed in solutions of Cl2 in water. Salts of HClO, however, are quickly produced on treatment of solutions of a base with chlorine at ambient temperature (Figure 3-29). These are disporportionation reactions, in which one-half of the Cl2 (oxidation state 0) is oxidized to HClO or its salt (oxidation state of Cl = +1) and the other half is reduced to Cl- (oxidation state of Cl = -1).
Figure 3-29. The formation of hypochlorous acid and its salts from Cl2.

Calcium hypochlorite, Ca(ClO)2, is a stable salt and a key active ingredient of bleaching powders. In contrast, pure solid NaOCl is unstable and can decompose explosively on friction or heating. Laundry bleach, however, can be stored without much decomposition for months. The bleaching effect deals with the oxidation of colored organic compounds and dyes by NaOCl. Important: acids decompose bleach solutions to toxic and corrosive chlorine gas (Cl2). Being a (weak) solution of acetic acid, vinegar should not be mixed with bleach solutions or powders to avoid the formation of Cl2.

Chlorous acid, HClO2, is the least stable of all oxoacids of chlorine. Its sodium salt (NaClO2), however, is rather stable and is used commercially to make chlorine dioxide, ClO2. Chlorine dioxide is a clean powerful oxidant employed in the disinfection of municipal drinking water and in wood pulp bleaching. Take warning: ClO2 solutions under the names Miracle Mineral Supplement, MMS, or CD have been falsely promoted and quite broadly marketed as a universal cure for many diseases ranging from common colds and flu to hepatitis and cancer. Read about the pseudomedical uses of ClO2 in this article.

Chloric acid, HClO3, is mostly known for its potassium salt, KClO3 (Berthollet's salt), a key ingredient of safety match heads and an essential component of certain types of pyrotechnics.

Perchloric acid, HClO4, the strongest acid discussed in our course, finds an important application in the production of liquid crystal displays. The largest quantities of perchloric acid, however, are manufactured to make its ammonium salt, NH4ClO4, for use as rocket fuel.

3.3.7. Bromine and Iodine. Both bromine and iodine are much less abundant than chlorine and fluorine. Bromine is produced by the treatment of brine that is rich in bromide ions with Cl2 (Figure 3-30). Iodine is produced similarly from iodide-rich brines (Figure 3-30) as well as from Chilean caliche rock and seaweed. Note that the equations presented in Figure 3-30 are the net ionic forms of the regular equations considered above (Figure 3-24). Basic chemical properties of bromine and iodine are similar to those of chlorine.
Figure 3-30. Chlorine gas displaces Br2 and I2 from bromide- and iodide-rich brines.
Digression. The discovery of iodine in 1811 is credited to Bernard Courtois, a French scientist. There is little doubt that Courtois discovered iodine serendipitously by inadvertently mixing seaweed ashes with concentrated sulfuric acid. Interestingly, however, some books and articles provide a fascinating detail of the discovery, suggesting that iodine was actually discovered not by Courtois himself but by his cat! "Iodine was discovered by Bernard Courtois or, to be more accurate, by his cat... It is said that his cat once pushed over a vessel containing sea-weed. The liquid mixed with something that had been spilt on the floor, and violet vapours appeared."

If this is a true story, what could that "something" be?

Seaweed contains iodine in the form of iodide, I-. The oxidation state of iodine of the iodide anion is -1, whereas the oxidation state of the iodine atom in the product of the reaction, I2, is 0. Clearly, in the accidental discovery, the iodide ion was oxidized to iodine:

2 I- - 2 e- = I2

This prompts the conclusion that the spilled "something" contained an oxidant capable of abstracting an electron from the iodide-anion in the seaweed. What might that oxidant have been?

In principle, any other halogen can oxidize iodide to iodine (Figures 3-24 and 3-30). In 1811, however, Br2 and F2 were not discovered yet. Chlorine was already known, but pure Cl2 is a gas that cannot be spilled like a liquid. A solution of Cl2 in water could, but on addition of an iodide source to aqueous Cl2 no violet vapors of I2 are given off. The resultant solution just turns dark due to the formation of I2, as shown in this demonstration.

So, what was it "that had been spilt on the floor" in Courtois's laboratory? I propose that it was oil of vitriol, the way they called concentrated sulfuric acid back in the old days. Concentrated H2SO4 was a common reagent then already, and it does react energetically with iodide salts to produce I2, as can be seen in this video. If performed on a sufficiently large scale, this exothermic reaction might generate enough heat for the I2 produced to (partially) sublime in the form of purple vapors. How does concentrated H2SO4 oxidize iodide to iodine? What element is reduced? There is little doubt that the sulfur atom that has the oxidation state of +6 accepts electrons from the iodide. In principle, the sulfur in H2SO4 can be reduced to SO2 (oxidation state +4), S (oxidation state 0), or H2S (oxidation state -2):

(a) 2 KI + 3 H2SO4 = I2 + SO2 + 2 KHSO4 + 2 H2O reduction to SO2

(b) 6 KI + 7 H2SO4 = 3 I2 + S + 6 KHSO4 + 4 H2O reduction to S

(c) 8 KI + 9 H2SO4 = 4 I2 + H2S + 8 KHSO4 + 4 H2O reduction to H2S

In the same demonstration, it is stated that the H2SO4 is reduced by the iodide to SO2, reaction (a). A commenter on the video questions that SO2 is produced and suggests that H2S is formed instead, reaction (c). I contest both outcomes because neither SO2 nor H2S can coexist with I2. Both SO2 and H2S reduce iodine to iodide. Interestingly, many websites teach reaction (a) or reaction (c) on one of their pages yet contradict themselves on another page by stating that I2 is reduced to I- by SO2 or H2S. Such a discrepancy can be found even in a serious modern scientific monograph on iodine, which states that iodide reduces H2SO4 to SO32- (equivalent of SO2) on page 563, yet also describes, on page 227, the reverse redox process, the reduction of I2 with SO2 to give iodide and H2SO4! I believe that equation (b) above portrays the most plausible outcome for the reaction between iodide and concentrated H2SO4. Unlike SO2 and H2S, sulfur can coexist with iodine. The only way of knowing for sure, though, is to either find a report on a thorough study of the reaction or investigate it experimentally.
3.3.8. Fluorine. The most electronegative element, fluorine, forms the simple substance F2 that is probably the most reactive chemical ever. Fluorine gas reacts with all metals and most nonmetals. Many reactions involving F2 are violent and highly exothermic, bursting into flames or even resulting in explosions. Watch this unique demonstration of spontaneous combustion of steel wool, charcoal, cotton wool, sulfur, and iodine on exposure to F2. Can you imagine a burning brick? Watch this video showing a brick burning on contact with fluorine gas. Even water burns in fluorine.

2 F2 + 2 H2O = 4 НF + O2

This reaction also gives rise to some quantities of F2O and other products.

3.3.9. Silver Halides. Chloride, bromide, and iodide ions react with the silver cation, Ag+, to give insoluble silver halides (Figure 3-31). This reaction is most characteristic of Cl-, Br-, and I-, and has long been used to test aqueous solutions for the presence of these halide ions. You can watch AgCl precipitate on mixing solutions of NaCl and AgNO3 here. Unlike AgCl, AgBr, and AgI, silver fluoride (AgF) is water-soluble.
Figure 3-31. Chloride, bromide, and iodide (but not fluoride) form insoluble silver halides on treatment with a soluble Ag+ salt.

Both AgCl and AgBr are white, although AgBr is light creamy white. Silver iodide, AgI, is pale yellow. Silver fluoride, AgF, is light brown as a solid, but its aqueous solutions are colorless. Figure 3-32 displays photos of freshly precipitated AgI, AgBr, and AgCl.
Figure 3-32. Freshly precipitated (left to right) pale yellow AgI, creamy white AgBr, and white AgCl (source).

Silver halides play the central role in sliver-based photography. Like many other silver salts, silver halides are photosensitive, decomposing to silver metal under sunlight. Although bulk Ag metal is silver in color, in a finely dispersed form silver and other metals are black. The areas on the film that are exposed to more light darken stronger due to the formation of Ag metal from the silver halide in the sensitive layer. Watch this excellent video demonstrating how AgCl is precipitated and subsequently used in imaging.

3.3.10. Halogen Compounds Around Us. We encounter various derivatives of halogens on a daily basis. Below are selected examples of halogen compounds that we use most frequently.

Fluorine. Fluorine in the form of fluoride is critical for healthy teeth, especially in children. Many brands of toothpastes contain fluorine, usually in the form of sodium fluoride (NaF) or sodium monofluorophosphate (Na2PO3F). Tin difluoride (SnF2) is used in some toothpastes for sensitive teeth. In many areas, fluoride is added to public water supplies, usually in the form of NaF, fluorosilicic acid (H2SiF6), or its sodium salt (Na2SiF6). The addition of a fluoride source to municipal water, called water fluoridation, is strictly controlled to meet the standards of the World Health Organization, 0.5-1.5 mg of fluoride per liter. In some countries, including China, India, and many European states, water is not artificially fluoridated for one reason or another. In some others, such as Spain, only a fraction of the population receives fluoridated water. In Ireland, water fluoridation is mandatory. In Switzerland, a fluoride source is added to table salt. About 50-70% of all residents of the U.S.A. and Canada receive fluoridated water.

The remarkable nonstick cookware employs materials based on fluoropolymers such as Teflon®. Teflon®, a unique material, has many applications ranging from car and aircraft building to nuclear industry and computer manufacturing. Fluorine-containing organic compounds often exhibit biological activity and are broadly used as active ingredients of pharmaceuticals and agrochemicals.

Chlorine. Table salt is NaCl. Most bleaching agents contain sodium or calcium hypochlorite as the main active ingredient. Potassium chloride is a broadly used fertilizer and a supplement to treat low blood pressure. Polyvinyl chloride (PVC) is one of the top three most broadly used polymers in the world. PVC is employed to manufacture water and sewage pipes, electric cables, and many materials for construction. Phonograph records are conventionally called "vinyl records" or just "vinyl" because they are made of PVC.

Chlorination of municipal water aims at killing bacteria to make tap water safe for drinking. To disinfect water in the chlorination process, Cl2 gas or NaClO are used in strictly controlled small amounts. Water chlorination is particularly important to prevent the spread of such serious water-borne diseases as cholera, typhoid, and dysentery.

Chlorine compounds are also active ingredients of many crop protection agents and pharmaceuticals. Two highly important cancer drugs Cisplatin and Mitotane are derivatives of chlorine.

Bromine. The main use of bromine is in the production of flame retardants. Brominated flame retardants are blended into flammable polymers and resins to make them nonflammable. The thus made materials are used to make various components for electronics (plastic covers, printed circuit boards, connectors, etc.) and kitchen appliances, as well as carpets, pillows, and upholstery. The use of bromine as well as chlorine and iodine in silver photography has reduced considerably with the spread of digital imaging techniques and devices.

Iodine. Being an essential element for life, iodine is of special importance for the thyroid gland. The thyroid gland is responsible for the production of hormones that regulate many vital body functions. Iodine deficiency in the body causes intellectual and developmental disabilities. To avoid this problem, small quantities of iodine compounds are added to table salt to make iodized salt. In some countries such as Argentina, Canada, and South Africa, all table salt for human consumption must be iodized. One of the oldest and still broadly used external antiseptics is tincture of iodine, a solution of iodine and KI or NaI in aqueous alcohol.

3.3.11. Exercises.

1. Write the electronic configuration of halogen atoms and explain why the halogens (a) are oxidizers and (b) exist in the form of diatomic molecules. [Answer: See 3.3.2]

2. The oxidizing ability of the halogens increases in the order F2 < Cl2 < Br2 < I2. True or false? Answer

3. All elemental halogens, F2, Cl2, Br2, and I2 are naturally occurring substances. True or false? Answer

4. How can laundry bleach be made from NaCl and water? Answer

5. Write balanced chemical equations for the reactions between (a) KBr and Cl2; (b) NaF and Cl2; (c) Al and Br2; (d) NaI and Br2; (e) CuO and HBr; (f) NaOH and Cl2; (g) BaCl2 and AgNO3; (h) Ca and I2 (i) Fe and Cl2; (j) Fe and HCl. Answer

6. If fluoridated water contains 0.84 mg of NaF per liter, what is the molar concentration of NaF in the water? Answer

7. It is often said that a more reactive halogen displaces a less reactive halogen from its salts. For example, in the reaction 2 KI + Cl2 = 2 KCl + I2, more reactive Cl2 displaces less reactive I2 from the iodide salt. There is a reaction, however, where iodine displaces chlorine from its salt: 2 KClO3 + I2 = 2 KIO3 + Cl2. Does the occurrence of this reaction make sense, or should it be viewed as a discrepancy? Answer

8. All hydrogen halides in water are strong acids. True or false? Answer

9. If no precipitate is formed on addition of AgNO3 to a water sample, the water is free from halide anions. True or false? Answer

10. Adding vinegar to laundry bleach is safe and highly recommended. True or false? Answer

11. Chlorine gas can be obtained in the reaction of HCl with (a) Br2; (b) KClO3; (c) K2CO3; (d) HF; (e) MnO2. Answer

12. Name HClO, HClO2, HClO3, HClO4 and their anions. Answer