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24 Schroeder - Chemistry of halogen family
Added by Mary Blackwell, last edited by Holly Schroeder on Nov 02, 2007

General Chemical Properties of the Halogens

One of the most important characteristics of the halogen family is that all the elements are oxidizing agents. In this oxidization, a halogen, whose typical oxidation number is 0, reduces to -1. In this state, halogens combine with other elements to form halides. We commonly know these halides as fluoride, chloride, bromide, iodide and astatide. Astatide is less well known than the other halogens because it has some metallic properties while the other halogens are nonmetals.

 Oxidizing properties of halogen elements increase from the bottom of the periodic family to top. i.e. fluorine has the highest level of oxidizing properties of the halogens. This corresponds to the fact that fluorides are usually more stable than corresponding chlorides, bromides, or iodides. Despite this, fluorine is not easily prepared in the free state, while iodine easily is.

The halogens, in group 7a of the periodic table, each have seven valence electrons in their outermost shells. These seven electrons are in the s and p orbitals. Potentially, a halogen could hold one more electron in its p-orbital, giving it the same stable configuration of the following noble gas in period 8. This explains why halogens are such strong oxidizing agents; it is easy to gain an electron when it results in becoming such a stable element. Halogen elements exist in their salts as halide ions, which are relatively large ions, and extremely stable.

At room temperature and atmospheric pressure, 1 atm, the halogen elements exist in their free states as diatomic molecules, which are molecules composed of two atoms each. Fluorine atoms are held together by sigma bonds, while chlorine and the halogens following it use their d-orbitals to create pi bonds, which are stronger than sigma bonds. 

General Physical Properties and Trends

The great reactivity of fluorine stems from the low energy of the fluorine sigma bond. The chlorine bond and pi bonds of the other halogens are of considerably higher energies.
Fluorine and Chlorine are gases at room temperature. Bromine is a red-brown liquid at room temperature, and iodine is a dark violet crystal solid.
The electron affinities for the halogen elements are all high and show only slight differences from one another.
Fluorine is the most electronegative of all elements, and there is a decrease in electronegativity within the family of the halogen elements from fluorine to iodine.
The ionization potentials of the halogens are generally high, but they decrease with decreasing atomic number.

The halogen elements all form compounds with hydrogen---the hydrogen halides. The energy of the hydrogen-halogen bond increases strongly from iodide to fluoride.


Fluorine reacts with nearly all other elements at room temperature. There is only one stable isotope of the element---fluorine-19. Because fluorine is the most electronegative of the elements, atomic groupings rich in fluorine are negatively charged. The ionization potential of fluorine is very high, at 420 kilocalories per gram, giving a standard heat formation for fluoride ion of 420 kilocalories per mole. This means that it is very easy for fluorine to gain an electron, which we know is favored because of the stabilization acquired in doing so. Since fluorine is the most powerfully oxidizing element, no other substance can reduce the fluoride ion to the free element This is why elemental fluorine is not found in its free state in nature.


Chlorine gas is easily liquefied by cooling or by pressures of a few atmospheres at room temperature. Chlorine has a high electronegativity and a high electron affinity. Chlorine reacts with many elements of both metals and nonmetals to give chlorides. However, chlorine is relatively inert toward carbon, nitrogen, and oxygen. The ionization potential of chlorine is high. In compounds, chlorine is capable of attaining the oxidation numbers of +1, +3, +4, +5, +6, and +7.

As a result of this, chlorine is capable of converting several oxides into chlorides. An example is the conversion of iron(III) oxide to the corresponding chloride:

2Fe~2~O~3~ + 6Cl~2~ ----> 4FeCl~3~ + 3O~2~

Chlorine is moderately soluble in water, yielding chlorine water. Chlorine water loses its efficiency as an oxidizing agent on standing, because hypochlorous acid gradually decomposes. The reaction of chlorine with alkaline solutions yields salts of oxo-acids.


Like the other halogens, bromine exists as a diatomic molecule at standard temperature and pressure. Bromine dissolves in aqueous alkali hydroxide solutions, giving bromides, hypobromites, or bromates, depending on the temperature. Bromine can be extracted from water by organic solvents such as carbon tetrachloride, in which it is very soluble.

The electron affinity of bromine is high and similar chlorine's. However, it is a less powerful oxidizing agent than chlorine. Though it shares many of the same properties as chlorine, a metal-bromine bond is weaker than the same metal-chlorine bond. 

Bromine combines violently with the alkali metals, but not as violently with other metals. Again, similar to chlorine, the ionization potential of bromine is high. Bromine can make compounds where its oxidation number is +1, +3, +4, +5, or +7. 


The ionization potential of the iodine atom is considerably smaller than that of the lighter halogen atoms, but the electron affinity of the iodine atom is not much different. Alkali iodides react with compounds containing iodine with the oxidation number +1, such as iodine bromide, as in the following equation:

KI~3~ +IBr ---> KBr + 2I~2~

In this reaction and those similar to this, alkali iodides may be regarded as Lewis bases.

Because of the reaction with Lewis bases, iodine must be regarded as a weak Lewis acid. Additionally, Iodine reacts also with iodide ions, because iodide ions can also act as Lewis bases. This property allows iodine to become significantly more soluble in water with the presence of an iodide. Iodine is a weaker oxidizing agent than bromine because the iodide ion is a weaker Lewis base than the bromide ion. The oxidizing properties of iodine are considerably weaker than those of the other halogens, as explained by the decreasing trend in oxidation strength from fluorine to iodine. Iodine combines directly with many elements. Silver and aluminum are easily converted into the respective iodides. Phosphorus unites readily with iodine.


The astatide ion is formed by reduction of the element, using, for example, zinc and acid. The astatide ion is a stronger reducing agent than the iodide ion, but free iodine is a stronger oxidizing agent than astatine.




+Chemical Principles: The Quest for Insight, Third Edition.+ W.H. Freeman and Company, 2005.

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