Chlorine Trifluoride

Introduction

Chlorine Trifluoride (ClF3) is highly combustible and explosive. During the time of World War II, the Nazis took an interest in ClF3 to make powerful bombs. However, the production was reduced once they came to know about its volatile and reactive nature. That is the main reason, after the war, it was never used in combat. ClF3 is banned under the Chemical Weapons Convention because of its destructive nature and even its production has a limit up to 30 tons.

What is Chlorine Trifluoride?

In the 1930s, two scientists Otto Ruff and H. Krug, came up with a liquid compound that is more reactive than Fluorine, they isolated the Chlorine Trifluoride $\mathrm{(ClF_{3})}$. It can be defined as an interhalogen compound that is represented by the chemical formula of $\mathrm{ClF_{3}}$. It is also known by the name Chlorine Fluoride and Trifluoro–λ3–Chlorane. This gas can cause lung damage if anyone gets close to that. Instead of that, it can bring irritation to the skin, eyes, and mucous membrane. $\mathrm{ClF_{3}}$ is considered one of the most powerful oxidizers, which can catch the fire, when comes in contact with any kind of inflammable materials.

Structure of Chlorine Trifluoride

Figure 1 − Chlorinetrifluoride

Chlorine-trifluoride, Public domain, via Wikimedia Common

In the terms of molecular geometry $\mathrm{ClF_{3}}$ has two long and one short bond, and because of those bonds, it takes the shape of a T. In a $\mathrm{ClF_{3}}$ molecule the Chlorine atom stands in the center. It has five regional densities they are three bonds and two lone pairs. $\mathrm{ClF_{3}}$ structure follows the VSEPR theory, because the lone pairs of the electrons present in the two equatorial positions, are Trigonal Bipyramid. They are arranged at 175 degrees of F-Cl-F bond angle. The bond of Cl-F is consistent because of the hypervalent bonding. In a quartz vessel, $\mathrm{ClF_{3}}$ is considered stable up to 180°C, above that, it is decomposed to its constituent elements and that happened because of the free radical mechanism.

Properties of Chlorine Trifluoride

Figure 2 − Electron structure of Chlorine Trifluoride

$\mathrm{ClF_{3}}$ is such a chemical compound that cannot be found as free in nature. As a compound, it is not only dangerous but also highly flammable. It has a density of 1.77g/cm3 and a molecular mass of 92.448g/mol. The boiling and melting point of $\mathrm{ClF_{3}}$ is 11.75°C and - 76.34°C, respectively. In the structure of $\mathrm{ClF_{3}}$, 4 atoms are present and it also has one covalently-bonded unit, moreover, it can be dissolved in water. In the gaseous form or as vapor, $\mathrm{ClF_{3}}$ can be decomposed to $\mathrm{ClF\:,\:ClOF\:,\:ClO_{2}F\:,\:ClO_{3}F\:,\:Cl_{2}\:and\:HF}$. The last three are the most significant amongst all of them. However, the result is always dependent on the availability of the water.

Preparation of Chlorine Trifluoride

ClF3 can be prepared by any direct action of the Chlorine and the Fluorine gases. The reaction can be represented as −

$$\mathrm{Cl_{2}\:+\:3F_{2}\:\rightarrow\:2ClF_{3}}$$

On the other hand, it can be made by the reaction between ClF and Chlorine gases. It can be present in an equation form and that is −

$$\mathrm{ClF_{3}\:+\:F_{2}\:\rightarrow\:ClF_{3}}$$

Use of Chlorine Trifluoride

• $\mathrm{ClF_{3}}$ can be used as a fluorinating agent.

• It can be used as the igniter and propellant in rockets.

• $\mathrm{ClF_{3}}$ is considered the most crucial component in the making of nuclear fuel

• It plays a significant role as a pyrolysis inhibitor for the fluoro carbon polymers.

• It is used to convert the uranium into gaseous hexafluoride uranium

• It is used for cleaning the chemical vapor deposition in the semiconductor industries.

• With the reaction of Phosphorus, $\mathrm{ClF_{3}}$ produces $\mathrm{PCl_{3}}$, which is Phosphorus Trichloride and $\mathrm{PF_{5}}$, which is Phosphorus Pentafluoride.

Facts about Chlorine Trifluoride

There are many important aspects of $\mathrm{ClF_{3}}$. If $\mathrm{ClF_{3}}$ comes in contact with any kind of element, it goes through the process of evaporation and turns into toxic gas. For the laboratory experiments, if the exposure of 400 ppm of $\mathrm{ClF_{3}}$ is for thirty minutes it can cause death for rats. Other than that, $\mathrm{ClF_{3}}$ has a total number of 28 valence electrons and the bond is between the Cl and its surrounding Fluorine atoms. Moreover, in the process of decomposition, $\mathrm{ClF_{3}}$ produces the hydrofluoric and hydrochloric acid, which comes in a steam form.

Conclusion

$\mathrm{ClF_{3}}$ is one of the most dangerous compounds to use, and it is considered dangerous because of its flammable nature. It can set fire even with some of the most inflammable materials, such as sand, glass, and even asbestos. The most dangerous part is that it can cause a fire with burnt elements such as ash. However, because of its multipurpose use, it is still considered useful.

FAQs

1. What is the main use of Chlorine Trifluoride?

The main use of Chlorine Trifluoride, is used for rocket fuels and can be used in the processing of nuclear reactor fuels.

2. What is the reason behind Chlorine Trifluoride’s flammable nature?

The reason behind Chlorine Trifluoride’s flammable nature is the oxidizing. It is more oxidizing than the oxygen itself. Not only it is flammable, but it is considered an extremely effective explosive because of its oxidizing nature.

3. What is the common name used for Chlorine Trifluoride?

Chlorine Trifluoride is also known as Chlorine Fluoride, which is represented as the $\mathrm{ClF_{3}}$.

4. How is Chlorine Trifluoride formed?

Chlorine Trifluoride is formed by the cation Chloride in the middle and three Fluoride anions is bound with that.

5. What is the color of Chlorine Trifluoride?

The gaseous form of Chlorine Trifluoride appears colorless, whereas the liquid form is green and has a pungent odor.