INTRODUCTION: Photochromic Sunglasses
We find it quite uncomfortable on the eyes to move rapidly in and out of brilliant sunlight and shadow, even if we are wearing sunglasses. Our eyes cannot adjust quickly enough to sudden changes of light intensity – that is, unless the sunglasses are made from photochromic glass, which compensates for changes in brightness.
Lens glass is usually made photochromic by adding tiny amounts of silver chloride and copper (I) chloride to the molten glass as it cools, which traps the crystals within the structure of the glass. When bright sunlight strikes the glass, the silver chloride decomposes to form metallic silver, which darkens the glass. This reduces the intensity of the light that reaches the eyes. At the same time, chlorine atoms are formed and they react with copper(II) ions to form copper(l) ions and chloride ions. As soon as the exposure to bright light ends, the copper(II) ions are reduced by silver atoms to reform silver chloride and copper(I) chloride. This lightens the glass and allows all the available light to reach the eyes.
In this way, spectacle lenses can be made to darken or lighten according to the intensity of the light, and so make it comfortable for people to keep their sunglasses on, whether they are in bright light or shadow.
WHAT ARE THE HALOGENS?
Halogens is the collective name given to the non-metal elements in Group 7 of the Periodic Table. The name is derived from hals (the Greek for salt) and the suffix gen (meaning producer). It was first used by the Swedish scientist Jöns Berzelius (1779-1884) to indicate that chlorine, bromine and iodine occurred in sea water. Chlorine was first prepared in 1774 by the Swedish chemist Carl Wilhelm Scheele, but he did not recognise it as an element.
Humphry Davy in 1810 recognised chlorine as an element and named it after the Greek word for green, chloros. In the same way, iodine was derived from ioides (the Greek word for violet-like) and bromine from bromos (the Greek word for stench). It was not until 1940 that astatine, a radioactive element and the halogen with the highest atomic number, was prepared artificially.
SOME USES OF THE HALOGENS AND THEIR COMPOUNDS
The halogens are commercially of great importance, although the use of many chlorine-containing compounds is controversial for environmental reasons. Chlorine and fluorine are both used in the production of polymers, such as polyvinyl chloride (PVC) and polytetrafluoroethene (PTFE), and fluorine is also used in toothpaste. Many insecticides contain chlorine, but there is concern over the use of these compounds, despite their obvious usefulness.
Bromine has a variety of applications, including the manufacture of fuel additives and of soil fumigators such as 1,2-dibromoethane and bromomethane, which kill pests found in the soil. Silver bromide and silver iodide both have applications in traditional photographic film. I will write more about the uses of halogens and their compounds later in this post.
PHYSICAL PROPERTIES OF THE HALOGENS
The halogens have physical properties that are typical of non-metals with a simple molecular structure. The halogens have relatively low melting and boiling points and are very poor conductors of heat and electricity.
A table showing some properties of four halogens
| Halogen | Melting Point/°C | Boiling Point/°C | Atomic (covalent) radius/pm | Electronegativity | Electron Configuration |
|---|---|---|---|---|---|
| Fluorine | -223 | -188 | 64 | 4 | 1s22s22p5 |
| Chlorine | -101 | -34 | 99 | 3 | 1s22s22p63s23p5 |
| Bromine | -7 | 59 | 114 | 2.8 | [Ar]3d104s24p5 |
| Iodine | 114 | 187 | 133 | 2.4 | [Kr]4d105s25p5 |
SIMPLE MOLECULAR STRUCTURE OF THE HALOGENS
All the halogens have diatomic molecules, which in the solid state are arranged in a simple molecular lattice. Halogen molecules are held in place by weak intermolecular forces known as induced dipole-induced dipole attraction. These forces result from an asymmetric distribution of electrons within each halogen molecule. This produces an instantaneous dipole, which induces dipoles in neighbouring molecules. Such a weak attractive force between molecules is easily overcome, so the elements have relatively low melting points and boiling points. The simple molecular structure has no free electrons, so a halogen, either as a solid or as a liquid, cannot conduct electricity.
All the halogens form solids with this type of intermolecular force. The larger the halogen molecule (and the more electrons), the easier it is to distort the electron cloud and increase the intermolecular forces. Hence, the melting point of the halogens increases with increasing atomic number.
Volatility measures the ease of evaporation of a substance. Liquids with weak intermolecular forces that are easy to overcome are very volatile, whereas liquids with strong intermolecular forces are not very volatile.
CHEMICAL PROPERTIES OF THE HALOGEN
Halogens react by gaining electrons to form an anion in ionic halides, or by sharing electrons to form a covalent bond in molecular halides. In both cases the halogen atom attains a noble gas electron configuration. Most metals form ionic halides and most non-metals form molecular halides. The ability to gain an electron is typical of a non-metal atom.
THE HALOGENS AS OXIDISING AGENTS
Since their atoms accept electrons, the halogens are oxidising agents and in a reaction they are reduced. Of the halogens, fluorine is the most powerful oxidising agent and astatine the least. This can be explained by the relative size of their atoms and their ability to capture an electron. The fluorine atom is the smallest, with fewer inner-shell shielding electrons, so its nucleus can have a greater attraction or ability to attract an electron. The reduction potentials, Eꝋred, of the halogens also illustrate this trend. The reaction of a fluorine molecule to give a fluoride ion has a very positive reduction potential, which indicates that this is a highly feasible process. (The more positive an electrode potential the more feasible the reaction.
Fluorine always has an oxidation number of -1 in its compounds, since a fluorine atom either gains one electron to form an ion or contributes one electron to a shared pair of electrons in a covalent bond. Fluorine oxidises other substances by gaining electrons from them.
Displacement reactions
Since fluorine is the most reactive halogen, in theory it can react with the halide ion of any of the other halogens. Fluorine becomes the fluoride ion and the free halogen (chlorine, bromine or iodine) is formed from the halide ion. This is called a displacement reaction. Experiments are not normally carried out with fluorine because its extraordinary reactivity makes it extremely dangerous. For example, if inhaled, fluorine can seriously damage the respiratory tract.
In aqueous solution, chlorine displaces bromide ions to form bromine:
CI2(aq) + 2KBr(aq) → 2KCl(aq) + Br2(aq)
During this reaction, the colour of the aqueous mixture becomes orange, which indicates that the element bromine has been produced. Chlorine will also displace iodide ions to form iodine, and again an orange-brown solution is formed.
CI2(aq) + 2I-(aq) → 2Cl-(aq) + I2(aq)
In the same way, aqueous bromine displaces aqueous iodide ion to form iodine. (The presence of iodine in a solution can be confirmed by the addition of starch solution. A dark blue coloration is produced). In general, the presence of the displaced halogen can be demonstrated easily if hexane is added after the displacement reaction, since bromine will preferentially dissolve in the hexane to give an orange-red solution, while iodine will give a purple solution.
The displacement reaction is an example of a redox reaction. This is because the halogen that reacts is reduced because it gains electrons to form halide ions, and halide ions are oxidised by the loss of electrons to form the halogen.
Manufacture of Bromine
The concentration of bromide ion in normal sea water is between 65 to 70 ppm. Therefore, a large quantity of sea water has to be processed to make significant quantities of bromine. Inland seas, such as the Dead Sea, have considerably higher concentrations of bromide ion and provide a much better feedstock.
Chlorine is used to liberate bromine by a displacement reaction. It is bubbled through the acidified sea water and the bromine formed is removed as a gas from the water by blowing air through it:
CI2(aq) + 2Br-(aq) → Br2 (aq) + 2Cl-
The bromine vapour formed is difficult to handle and so is converted back into bromide ion by reaction with sulfur dioxide and water:
Br2(aq) + SO2(g) + H2O(l) → H2SO4(aq) + 2HBr(aq)
This makes a much more concentrated solution of bromide ion, from which it is easier to produce high-purity bromine. A second displacement reaction with chlorine yields bromine vapour, which can be condensed and then purified by distillation.
REACTIONS OF HALOGENS WITH ELEMENTS
I have already described some of the reactions of metals and non-metals with chlorine in some of my previous posts. In every case the halogen oxidises the element. Fluorine will react the most violently, and astatine the least.
Reaction of halogens with non-metals
Almost all non-metals react with fluorine and chlorine to form covalent halides. The halogen atom shares electrons to make a single covalent bond. Very often, the halide contains the element in its highest possible oxidation state. In the reaction between phosphorus and chlorine, the reaction can be controlled to make either phosphorus (III) chloride or phosphorus (V) chloride:
P4(s) + 6Cl2(s) → 4PCl3(l)
P4(s) + 10Cl2(s) → 4PCl5(l)
Reactions of the halogens with metals
Most metals react with the halogens to form ionic halides. This reaction is necessarily a redox reaction, in which the metal is oxidised and the non-metal is reduced. Electrons are transferred to halogen atoms from metal atoms. The reaction of a metal with the halogens varies with the halogen. For instance, the reaction between a metal and fluorine is faster, more exothermic and more violent than that between the same metal and iodine.
Since halogens are good oxidising agents, if the metal has more than one oxidation state then in the halide formed the metal is often in one of its higher oxidation states. So, when chlorine is passed over hot iron wire it forms iron (III) chloride, FeCl3, rather than iron (II) chloride. Similarly, copper and chlorine form copper(II) chloride, CuCl2, rather than copper (I) chloride.
REFERENCES
https://www.opto-reseau.com/en/blog/benefits-of-photochromic-lenses
https://en.wikipedia.org/wiki/Photochromic_lens
https://www.allaboutvision.com/lenses/photochromic.htm
https://www.iloencyclopaedia.org/component/k2/item/1047-halogens-and-their-compounds
https://www.britannica.com/science/halogen#:~:text=What%20are%20some%20uses%20of,supplies%20to%20prevent%20tooth%20decay.
https://courses.lumenlearning.com/boundless-chemistry/chapter/halogens/
https://opentextbc.ca/chemistry/chapter/18-11-occurrence-preparation-and-properties-of-halogens/
https://edu.rsc.org/resources/reactions-of-halogens-as-aqueous-solutions/733.article
https://en.wikipedia.org/wiki/Halogen
https://chemed.chem.purdue.edu/genchem/topicreview/bp/ch10/group7.php
https://www.bbc.co.uk/bitesize/guides/ztjy6yc/revision/2
https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Modules_and_Websites_