Lecture 2. Physical properties of the Halogens

Some chemical and physical properties of the halogens are summarized in the Table below.
It can be seen that there is a regular increase in many of the properties of the halogens proceeding down group 17 from fluorine to iodine. This includes their melting points, boiling points, intensity of their colour, the radius of the corresponding halide ion, and the density of the element. On the other hand, there is a regular decrease in the first ionization energy as we go down this group. As a result, there is a regular increase in the ability to form high oxidation states and a decrease in the oxidizing strength of the halogens from fluorine to iodine.

The amount of energy required for excitation depends upon the size of the atom. Fluorine is the smallest element in the group and the force of attraction between the nucleus and the outer electrons is very large. As a result, it requires a large excitation energy and absorbs violet light (high energy) and so appears pale yellow. On the other hand, iodine needs significantly less excitation energy and absorbs yellow light of low energy. Thus it appears dark violet. Using similar arguments, it is possible to explain the greenish yellow color of chlorine and the reddish brown color of bromine.

The halogens show a variety of colours when dissolved in different solvents. Solutions of iodine can be bright violet in CCl₄, pink or reddish brown in aromatic hydrocarbons and deep brown in alcohols for example. This can be explained by weak donor-acceptor interaction and complex formation. The presence of charge-transfer bands further supports this since they are thought to be derived from interaction with the HOMO σ_u* orbital.

The X-ray structure of some of these have been obtained and often the intense colour can be used for characterisation and determination such as the bright blue colour of iodine in the presence of starch. In the case of the solid formed between dibromine and benzene, the structure is shown below and a new charge transfer band occurs at 292 nm. The Br-Br bond length is essentially unchanged from that of dibromine (228 pm).

In a study of the reaction of dibromine with substituted phosphines in diethyl ether, all but one showed a tetrahedral arrangement where one bromine was linked to the phosphorus.[3]

The X-ray study of the triethylphosphine was interpreted as [Et₃PBr]Br where the Br-Br separation was 330 pm. This is considerably longer than the 228 pm found above and was taken to mean that the compound was ionic In the case of the tri(perfluorophenyl)phosphine however the structure showed both bromines linked to give a trigonal bipyramid arrangement with D3 symmetry. Why (C₆F₅)₃PBr₂ was the only R₃PBr₂ compound that adopted trigonal bipyramidal geometry was reasoned to be due to the very low basicity of the parent tertiary phosphine.

Intermolecular forces are the attractive forces between molecules without which all substances would be gases. The various types of these interactions span large differences in energy and for the halogens and interhalogens are generally quite small. The dispersion forces involved in these cases are called London forces (after Fritz Wolfgang London, 1900-1954). They are derived from momentary oscillations of electron charge in atoms and hence are present between all particles (atoms, ions and molecules).

To picture how this occurs, compare the situation 1) where the electrons are evenly distributed and then consider 2) an instantaneous dipole that would arise from an uneven distribution of electrons on one side of the nucleus. When two molecules are close together, the instantaneous dipole of one molecule can induce a dipole in the second molecule. This results in synchronised motion of the electrons and an attraction between them. 3) Multiply this effect over numerous molecules and the overall result is that the attraction keeps these molecules together, and for diiodine is sufficient to make this a solid.

On average the electron cloud for molecules can be considered to be spherical in shape. When two non-polar molecules approach, attractions or repulsions between the electrons and nuclei can lead to distortions in their electron clouds (i.e. dipoles are induced). When more molecules interact these induced dipoles lead to intermolecular attraction.

The changes seen in the variation of MP and BP for the dihalogens and binary interhalogens van be attributed to the increase in the London dispersion forces of attraction between the molecules. In general they increase with increasing atomic number.

1. Enthalpy of Atomisation,              2. Δ_EAH₁,              3. Δ_hydH°(X^-,g)

The redox potential, E°, X₂/2X^-, measures a free-energy change, usually dominated by the ΔH term. The values in the Table show that there is a decrease in oxidising strength proceeding down the group (2.87, 1.36, 1.09, 0.54 V). This can be explained by comparing the steps shown above.

1) Atomisation of the dihalide is the energy required to break the molecule into atoms
½ X₂(g) → X(g)
note that only F₂ and Cl₂ are gases in their natural state so the energies associated with atomisation of Br₂ and I₂ requires converting the liquid or solid to gas first.

2) Δ_EAH₁ is the energy liberated when the atom is converted into a negative ion and is related to the Electron Affinity
X(g) + e^- → X^-(g)

Addition of an electron to the small F atom is accompanied by larger e^-/e^- repulsion than is found for the larger Cl, Br or I atoms. This would suggest that the process for F should be less exothermic than for Cl and not fit the trend that shows a general decrease going down the group.

3) Δ_hydH°(X^-,g) is the energy liberated on the hydration of the ion, the Hydration energy.
X^-(g) + H₂O → X^-(aq)

This shows a very negative energy change for the fluoride compared to the others in the group. This comes about because of two main factors: the high hydration energy and the low atomisation energy. For F₂ 2) is less than for Cl₂ but since the energy needed to break the F-F bond is also less and the hydration more, the total energy drop is much greater. In spite of their lower atomisation energies, Br₂ and I₂ are weaker oxidising agents than Cl₂ and this is due to their smaller Δ_EAH₁ and smaller Δ_hydH°.

It can be seen that the Δ_EAH₁ value for fluorine is in between those for chlorine and bromine and so this value alone does not provide a good explanation for the observed variation.

Each of the halogens is able to oxidise any of the heavier halogens situated below it in the group. They can oxidise hydrogen and nonmetals such as:

In water, the halogens disproportionate according to:

3Cl₂(g) + 6OH^-(aq) → ClO₃^-(aq) + 5Cl^-(aq) + 3H₂O(l)

The trend seen for the complete removal of an electron from the gaseous halogen atoms is that fluorine has the highest IE₁ and iodine the lowest. To overcome the attractive force of the nucleus means that energy is required so that the Ionisation Energies are all positive. The variation with size can be explained since as the size increases it take less energy to remove an electron. This inverse relationship is seen for all the groups, not just group 17. As the distance from the nucleus to the outermost electrons increases, the attraction decreases so that those electrons are easier to remove. The high value of IE₁ for Fluorine is such that it does not exhibit any positive oxidation states, whereas Cl, Br and I can exist as high as 7.

Fluorine is the most electronegative element in the periodic table and exists in all its compounds in either the -1 or 0 oxidation state. Chlorine, bromine, and iodine however can be found in a range of oxidation states including: +1, +3, +5, and +7, as shown below.

In general, odd numbered groups (like group 17) form odd-numbered oxidation states and this can be explained since all stable molecules contain paired electrons. (Free radicals are obviously much more reactive). When covalent bonds are formed or broken two electrons are involved so the oxidation state changes by 2.

References:

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