Chlorine

The Halogens are an important chemical group in the periodic table. Halogens have had many uses throughout history, including disinfecting water (Br), rocket fuel (F), and disinfecting in medicine (I). However, one of the most brutal and lethal uses of halogens in history was in World War I with the use of chlorine gas. The Germans first started using chlorine gas as a weapon in 1915. It was used in the battle of Ypres and backfired on the German troops. The wind changed and they had to retreat to avoid the gas. With the second battle of Ypres the Germans were more successful at attacking the allies. The chlorine gas created 9,000 casualties and 1,000 deaths on the Russian line.

Figure \(\PageIndex{1}\). British Soldiers in chlorine Gas Cloud

The reason that chlorine gas is so deadly is that it affects the respiratory system. It can cause swelling of the trachea and asphyxiation with prolonged exposure. During the war many soldiers would run away from the gas. This however increased the affects of the gas. A better counter measure to a chlorine gas attack was to find high ground. chlorine gas stays close to the ground (due to its higher density than air, more physical properties will be discussed later), and as a result many of the men affected were the ones in the trenches. The men who simply climbed out succumbed to far less affects than those who stayed below. Another effective treatment of the chlorine gas was to place a urine soaked towel over ones mouth. The urea in urine reacts with the chlorine in the following reaction:

\[\ce{CON2H4 + 2Cl2 -> CONH2Cl2 + 2HCl}\nonumber\]

The product formed is called dichlorourea, precipitates as a solid, and is soluble in water. Because of these properties of dichlorourea, it is not toxic and therefore was an effective treatment when chlorine gas was used against soldiers. chlorine gas started to become obsolete in WWI soon after its use began because of the ease at which it was defended against. Casualties were large, but with new knowledge of how chlorine behaved the allies were able to combat the German gas attacks.

Halogen Reactions and Properties

Chlorine has many other reactions and properties along with the ones listed above. It is readily reduced to X^- form, making it a good oxidizing agent. This qualitative aspect and others are widely shared by the other halogens, with a few exceptions due to increasing and other differences one would find moving up or down a group in the periodic table.

For example, all halogens are quite reactive, and in the natural world they always occur combined with other elements. Fluorine reacts so readily with almost any substance it contacts that chemists were not successful in isolating pure fluorine until 1886, although its existence in compounds had been known for many years. Chlorine, bromine, and iodine are progressively less reactive but still form compounds with most other elements, especially metals. A good example of this is mercury.

\[\ce{Hg (l) + Cl2 (g, l, or s) -> HgCl2 (s)}\nonumber\]

The following table lists the halogens and their basic properties. The differences between each element can be compared by moving up and down a column.

Table \(\PageIndex{1}\): Properties of chlorine, a Group VIIA Element.

Element	Symbol	Atomic Number	Density	Molar Mass	Room Temp Phase	Electron Configuration	Usual Oxidation State
Fluorine	F	9	1700 g/mL	18.998 g/mol	gas	[He]2s^22p^5	-1
Chlorine	Cl	17	3.2 g/mL	35.453 g/mol	gas	[Ne]3s^23p^5	+7, +5, +3, +1, -1
Bromine	Br	35	3.1028 g/mL	79.904 g/mol	liquid	[Ar]4s^23d^104p^5	+7, +5, +4, +3, +1, -1
Iodine	I	53	4.933 g/mL	126.9 g/mol	solid	[Kr]5s^24d^105p^5	+7, +5, +3, +1, -1

Element	Radius (pm)		Ionization Energy (MJ mol^–1)			Electronegativity	Melting Point (in °C)	Boiling Point (in °C)
	Covalent	Ionic	First	Second	Third
Fluorine	64	136	1.681	3.374	6.050	3.98	-219.62	-188.1
Chlorine	99	181	1.251	2.298	3.822	3.16	-101	-34.0
Bromine	114	195	1.140	2.103	3.470	2.96	-7.2	58.8
Iodine	139	215	1.008	1.846	3.180	2.66	113.7	184.3

Already covered in the section on alkali metals, chlorine reacts readily with alkali metals with the general form of:

\[\ce{2Na + Cl2 -> 2NaCl}\nonumber\]

Compounds of an alkali metal and a halogen, such as sodium chloride, potassium fluoride, lithium bromide, or cesium iodide, have closely related properties. (All taste salty, for example.) They belong to a general category called salts, all of whose members are similar to ordinary table salt, sodium chloride. The term halogen is derived from Greek words meaning “salt former.” Chlorine was perhaps the most widely used salt in ancient Greece. Salt held high importance in ancient society as a food preserver and a addition to foods for taste. Without salt, meats and other foods would have spoiled in mere days. Salt allowed them to be stored for months. This was especially important when armies were on long military campaigns. Had the Greek soldiers not been able to preserve food with salt they would not have been able to conquer as much territory. The salt used by the Greeks, and most everyone in present times uses NaCl as salt, which comes from the reaction of sodium and chlorine shown above.

Chlorine also reacts with alkaline-earth metals in the general reaction:

\[\ce{M + Cl2 -> MCl2}\nonumber\]

M = Be, Mg, Ca, Sr, Ba, or Ra

Another vigorous reaction occurs when certain compounds containing carbon and hydrogen contact the halogens. Turpentine, C₁₀H₁₆, reacts quite violently. In the case of chlorine the equation is

\[\ce{C10H16 (l) +8Cl2 (g) -> 10C (s) + 16HCl (g)}\nonumber\]

The violent reaction is due to α-pinene in turpentine. The relief of ring strain is highly exothermic. This temperature increase causes the sublimation.

Chlorine also reacts directly with hydrogen, yielding the hydrogen chloride:

\[\ce{H2 + Cl2 -> 2HCl}\nonumber\]

This compound is a gas, water soluble, and, a strong acid in aqueous solution. It is conveniently prepared in the laboratory by acidifying the appropriate sodium chloride:

\[\ce{NaCl (s) + H3O+ (aq) ->[\Delta] Na+ (aq) + H2O (l) + HCl (g)}\nonumber\]

The production and eventual use of HCl is an important one in history. Alchemists, scientists who wanted to find the philosophers stone, discovered in the 15th century that when Nitric acid was combined with hydorchloric acid, gold was dissolved. Still to this day that combination is one of the only substances that will dissolve gold. The alchemists called their new creation aqua regia. The reason that the combination of the two acids works is because Nitric acid is a strong oxidizer and when it comes in contact with gold it creates very few gold ions, Au³⁺, which can be picked off by the Cl^- ions of HCl to form chloroaurate anions (AuCl₄^-). The ions are dissolved in solution, and then more oxidation can occur until all of the gold is dissolved. The reaction mechanism is as follows:

\[\ce{Au (s) + 3NO3- (aq) + 6H+ (aq) -> Au(3+) (aq) + 3NO2 (g) + 3H2O (l)}\nonumber\]

\[\ce{Au(3+) (aq) + 4Cl- (aq) -> AuCl4- (aq)}\nonumber\]

Aqua regia is still used today in many gold processing factories as a way to get rid of impurities in the metal.

The relative oxidizing strengths of chlorine can be illustrated nicely in the laboratory. If, for example, a solution of Cl₂ in H₂O is combined with a solution of NaI, the dark color of I₂ can be observed, showing that the Cl₂ has oxidized the I^–:

\[\ce{Cl2 (aq) + 2I- (aq) -> 2Cl- (aq) + I2 (aq)}\nonumber\]

This very reaction is shown in the following video:

The video starts out with four solutions. The experimental solution is on the far left, and contain Cl₂ in water, which is covered by a layer of hexane, a nonpolar solvent which is immiscible with H₂O. The three other solutions, from left to right are a Cl₂ solution, a Br₂ solution, and an I₂ solution. When a solution with iodide ions is added to the experimental solution, nonpolar I₂ molecules are formed. They concentrate in the hexane layer, and a beautiful violet color can be observed, the same as I₂ solution. From such experiments it can be shown that the strongest oxidizing agent is F₂ (at the top of the group). The weakest oxidizing agent, I₂, does not react with any of the halide ions. Chlorine is the second highest oxidizing agent, behind fluorine.

Chlorine is also capable of oxidizing water, but it does so very slowly. Instead the reaction

\[\ce{Cl2 + 2H2O <-> H3O+ +Cl- + HOCl}\nonumber\]

goes partway to completion. Hypochlorous acid, HOCl, is a weak acid. In basic solution the halogen is completely consumed, producing the hypohalite anion:

\[\ce{Cl2 + 2OH- -> Cl- + H2O + OCl-}\nonumber\]

Since hypochlorite, OCl^–, could also be supplied from an ionic compound such as NaOCl, the latter is often used to chlorinate swimming pools.

Hypohalite ions disproportionate in aqueous solution:

\[\ce{3OCl- ->2Cl + ClO3-}\nonumber\]

This reaction is rather slow for hypochlorite unless the temperature is above 75°C, but OBr^– and OI^– are consumed immediately at room temperature. Chlorate, ClO₃^–, bromate, BrO₃^–, and iodate, IO₃^–, salts can be precipitated from such solutions. All are good oxidizing agents. Potassium chlorate, KClO₃, decomposes, giving O₂ when heated in the presence of a catalyst:

\[\text{2KClO}_3 \xrightarrow [\text{MnO2 catalyst}]{\Delta} \text{2KCl} + \text{3O}_2\nonumber\]

This is a standard laboratory reaction for making O₂.

If KClO₂ is heated without a catalyst, potassium perchlorate, KClO₄, may be formed. Perchlorates oxidize organic matter rapidly and often uncontrollably. They are notorious for exploding unexpectedly and should be handled with great care.

One other interesting group of compounds chlorine is a part of is the interhalogens, in which one halogen bonds to another. Some interhalogens, such as BrCl, are diatomic, but the larger halogen atoms have room for several smaller ones around them. Thus a compound such as ClF₃ can be synthesized. The following video showcases a reaction which involves some interhalogens:

The video begins with a test tube containing a layer of KI aqueous solution on top of CCl₄ below it. Chlorine is bubbled through the KI layer. As seen in the video on oxidizing strength of the halogens, Cl₂ reacts with I^- to form iodine, according to the reaction:

\[\ce{2I- (aq) + Cl2 (aq) -> I2 (aq) + 2Cl- (aq)}\nonumber\]

A brown triiodide ion is also formed in the aqueous layer, according to the reaction:

\[\ce{I- (aq) + I2 (aq) -> I(3-) (aq)}\nonumber\]

A purple solution begins to form in the CCl₄ layer, as iodine dissolves in it. The iodine in the aqueous layer also reacts with the excess Cl₂ to form the red ICl, according to the following reaction:

\[\ce{I2 (aq) + Cl2 (aq) -> 2ICl (aq)}\nonumber\]

The final reaction takes place as more Cl₂ is added, which reacts with ICl, to form the yellow ICl₃. This reaction causes the aqueous solution to decolorize. This goes according to the reaction:

\[\ce{ICl (aq) + Cl2 (aq) -> ICl3 (aq)}\nonumber\]

At the end of the video, the layers have decolorized, with a red portion in the CCl₄ which is, due to its color, most likely remaining ICl.

From ChemPRIME: 12.7: Group VIIA: Halogens

References

1. www.jstor.org/stable/93011?seq=2
2. en.Wikipedia.org/wiki/Poison_gas_in_WWI
3. Encyclopedia Britannica 1911, Alchemy
4. http://www.americanchemistry.com/s_chlorine/sec_content.asp?CID=1166&DID=4476&CTYPEID=109

Contributors and Attributions

• Ed Vitz (Kutztown University), John W. Moore (UW-Madison), Justin Shorb (Hope College), Xavier Prat-Resina (University of Minnesota Rochester), Tim Wendorff, and Adam Hahn.