The carbon tetrahalides are generally prepared by the direct (thermal) reaction of carbon with the appropriate halogen, Equation 1; however, specific syntheses are possible for each derivative.

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In addition to the direct reaction of fluorine with carbon, CF₄ can be prepared from SiC, Equation 2. The SiF₄ side product is removed by passing the reaction mixture through NaOH solution, in which SiF₄ reacts to form silicate. The difference in reactivity of SiF₄ and CF₄ is attributable to the lack of an energetically accessible five-coordinate intermediate required for the associative mechanism.

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Carbon tetrabromide can be obtained by bromination of CH₄ with HBr or Br₂, or by the reaction of CCl₄ with AlBr₃, Equation 3. Carbon tetraiodide (CI4) can be made by the Lewis acid catalyzed halogen exchange reaction, Equation 4.

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CF₄ is very stable. In fact, it is so stable that it does not even react with molten sodium. In contrast to CF₄, carbon tetrachloride (CCl₄) reacts readily with alkali metals (K and Na) or other strong reducing agents (e.g., F₂, Al, Ba, Be, and Zn). While CCl₄ is thermodynamically unstable with respect to hydrolysis, it is kinetically stable, and thus finds extensive use as a solvent. Photolysis can result in the transfer of a chloride radical to various substrates. It is also used in the conversion of metal oxides to the chlorides. Carbon tetrabromide (CBr₄) is insoluble in water and other polar solvents, but soluble in benzene. Carbon tetraiodide (CI₄) decomposes thermally, Equation 5.

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The decreasing stability of CX₄, from fluorine to iodine, is directly related to the C-X bond energy (Table 2).

Organic compounds that contain only carbon and a halogen are called halocarbons, and these include fluorocarbons and chlorocarbons. The easiest route to fluorocarbons involves the reaction of a hydrocarbon with a high valent fluoride (e.g., CoF₃) or the reaction of a chlorocarbon with SbF₃. In general, chlorocarbons with sp³ carbon atoms are more stable than those with sp² carbon centers. The exception to this is aromatic compounds such as C₆Cl₆.

The physical properties of fluorocarbons range from inert to toxic. Thus, poly(tetrafluoroethylene), (C₂F₄)_n, known by either its acronym (PTFE) or its trade name (Teflon), is chemically inert and has a low coefficient of friction (Table 3). As a consequence its uses include coatings on armor-piercing bullets (to stop the wear on the gun barrel), laboratory containers and magnetic stirrers, tubing for corrosive chemicals, and thread seal tape in plumbing applications (plumbers tape). A summary of the physical properties of PTFE is given in Table 3. PTFE is synthesized by the emulsion polymerization of tetrafluoroethylene monomer under pressure through free radical catalyst, Equation 6.

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In contrast with PTFE, octafluoroisobutylene, (CF₃)₂C=CF₂, is highly toxic, while perfluorodecahydronapthalene (C₁₀F₈, Figure 1) is used as a blood substitute component.

Figure 1: Structure of perfluorodecahydronapthalene.

Mixed halides are an important class of halocarbon compound. They are synthesized by halide exchange, Equation 7. The high cost of SbF₃ means that the reaction is generally run with an excess of the chloride.

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The ordinary name for mixed carbon halide is halon or Freon, although Freon is actually a Du Pont trademark. A list of selected Freon compounds are given in Table 4. Halons are non-toxic, non-flammable, and have no odor. However, it is their very lack of reactivity that has caused a problem.

Environmental impact of chlorofluorcarbon compounds (CFCs)

Chlorofluorcarbon compounds (CFCs) are very stable and are not degraded in the environment. As a consequence they are transported to the stratosphere where they decomposed upon photolysis, Equation 8. The resulting chloride radical is a catalyst for the decomposition of ozone, Equation 9, as well as a catalyst for the reaction of ozone with molecular oxygen, Equation 10.

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The widespread use of CFCs as refrigerants and propellants meant that by 1986 there were 2.5 billion pounds of CFC being liberated to the atmosphere. This was equivalent to ¹/₂. lb per person on the planet. Since the ozone layer provides the vital protection to life on the Earth’s surface from high energy UV radiation the release of CFC (along with other chemicals) caused a dramatic change in the ozone layer, including the increase in the polar hole in the ozone layer. As a result of the EU called for a complete ban of CFCs (which was followed by other countries). In their place new chemicals with similar refrigerant properties were developed. These compounds contained C-H bonds (e.g., C₂HCl₂F₃ and C₂H₃Cl₂F) that are readily broken in the lower atmosphere, thus limiting the transport to the stratosphere.

All the carbonyl halides (X₂C=O, X = F, Cl, Br, I) are known (Table 1). Phosgene (Cl₂C=O) was first synthesized by John Davy (Figure 2) in 1812 by exposing a mixture of carbon monoxide and chlorine to sunlight, Equation 11. He named it phosgene from the Greek, phos (light) and gene (born), in reference to use of light to promote the reaction. The fluoride is also prepared by the reaction of carbon monoxide with the halogen, while the bromide is prepared by the partial hydrolysis of CBr₄ with sulfuric acid.

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Figure 2: British doctor and chemist John Davy FRS (1790 – 1868) was also the brother of the noted chemist Sir Humphry Davy.

The synthesis of isocyanates from alkyl or aryl amines illustrates the electrophilic character of phosgene and its ability to introduce the equivalent of "CO²⁺", Equation 12. This reaction is conducted in the presence of a base such as pyridine that absorbs the hydrogen chloride. Phosgene may also be used to produce acyl chlorides from carboxylic acids, Equation 13. However, thionyl chloride is more commonly and more safely used in this reaction.

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Phosgene as a weapon of war

Phosgene is a toxic gas with the smell of “new-mown hay” and was used in chemical warfare during the First World War (Figure 3) where it was a more potent weapon than chlorine. While chlorine was potentially deadly it caused the victim to violently cough and choke (the bodies natural defense to limiting inhalation), in contrast, phosgene caused much less coughing with the result that more of it was inhaled. Phosgene often had a delayed effect; apparently healthy soldiers were taken down with phosgene gas poisoning up to 48 hours after inhalation. A fatal dose of phosgene eventually led to shallow breathing and retching, pulse up to 120, an ashen face and the discharge of four pints of yellow liquid from the lungs each hour for the 48 of the drowning spasms.

Figure 3: Australian infantry from the 45^th Battalion, Australian 4^th Division wearing Small Box Respirators (SBR) as protection of chemical warfare agents. Photograph taken at Zonnebeke, on the Ypres sector, 27^th September 1917. Copyright: Australian War Memorial.

Although phosgene’s boiling point (7.6 °C) meant that is was a vapor, the so-called "white star" mixture of phosgene and chlorine was commonly used on the Somme, because the chlorine supplied the necessary vapor with which to carry the phosgene. A summary of the casualties inflicted by chemicval warefare agents during the Great War is shown in Table 5.

Table 5: Casualties from gas attacks during the First World War (including chlorine, phosgene, and mustard gas). British Empire includes troops from United Kingdom, Australia, Canada, India, New Zealand, and South Africa.

Country	Total casualties	Deaths
Russia	419,340	56,000
Germany	200,000	9,000
France	190,000	8,000
British Empire	188,706	8,109
Austria-Hungary	100,000	3,000
USA	72,807	1,462
Italy	60,000	4,627
Others	10,000	1,000