Group 17

General
Properties of
Halogens
 Introduction
Halogens form diatomic molecules (of the form X 2, where X denotes a halogen atom) in
their elemental states. The bonds in these diatomic molecules are non-polar covalent
single bonds. However, halogens readily combine with most elements and are never seen
uncombined in nature. As a general rule, fluorine is the most reactive halogen and
astatine is the least reactive. All halogens form Group 1 salts with similar properties. In
these compounds, halogens are present as halide anions with charge of -1 (e.g. Cl -, Br-,
etc.). Replacing the -ine ending with an -ide ending indicates the presence of halide
anions; for example, Cl- is named "chloride." In addition, halogens act as oxidizing agents
—they exhibit the property to oxidize metals. Therefore, most of the chemical reactions
that involve halogens are oxidation-reduction reactions in aqueous solution. The halogens
often form single bonds, when in the -1 oxidation state, with carbon or nitrogen in
organic compounds. When a halogen atom is substituted for a covalently-bonded
hydrogen atom in an organic compound, the prefix halo- can be used in a general sense,
or the prefixes fluoro-, chloro-, bromo-, or iodo- can be used for specific halogen
substitutions. Halogen elements can cross-link to form diatomic molecules with polar
covalent single bonds.

Chlorine (Cl2) was the first halogen to be discovered in 1774, followed by iodine (I 2),
bromine (Br2), fluorine (F2), and astatine (At, discovered last in 1940). The name
"halogen" is derived from the Greek roots hal- ("salt") and -gen ("to form"). Together
these words combine to mean "salt former", referencing the fact that halogens form salts
when they react with metals. Halite is the mineral name for rock salt, a natural mineral
consisting essentially of sodium chloride (NaCl). Lastly, the halogens are also relevant in
daily life, whether it be the fluoride that goes in toothpaste, the chlorine that disinfects
drinking water, or the iodine that facilitates the production of thyroid hormones in one's
body.
 Elements
Fluorine - Fluorine has an atomic number of 9 and is denoted by the symbol F.
Elemental fluorine was first discovered in 1886 by isolating it from hydrofluoric acid.
Fluorine exists as a diatomic molecule in its free state (F 2) and is the most abundant
halogen found in the Earth's crust. Fluorine is the most electronegative element in the
periodic table. It appears as a pale yellow gas at room temperature. Fluorine also has a
relatively small atomic radius. Its oxidation state is always -1 except in its elemental,
diatomic state (in which its oxidation state is zero). Fluorine is extremely reactive and
reacts directly with all elements except helium (He), neon (Ne) and argon (Ar). In H 2O
solution, hydrofluoric acid (HF) is a weak acid. Although fluorine is highly
electronegative, its electronegativity does not determine its acidity; HF is a weak acid due
to the fact that the fluoride ion is basic (pH>7). In addition, fluorine produces very
powerful oxidants. For example, fluorine can react with the noble gas xenon and form the
strong oxidizing agent Xenon Difluoride (XeF 2). There are many uses for fluorine, which
will be discussed in Part VI of this article.

Chlorine - Chlorine has the atomic number 17 and the chemical symbol Cl. Chlorine was
discovered in 1774 by extracting it from hydrochloric acid. In its elemental state, it forms
the diatomic molecule Cl2. Chlorine exhibits multiple oxidation states, such as -1, +1, 3,
5, and 7. At room temperature it appears as a light green gas. Since the bond that forms
between the two chlorine atoms is weak, the Cl 2 molecule is very reactive. Chlorine reacts
with metals to produce salts called chlorides. Chloride ions are the most abundant ions
that dissolve in the ocean. Chlorine also has two isotopes: 35Cl and 37Cl. Sodium chloride
is the most prevalent compound of the chlorides.

Bromine - Bromine has an atomic number of 35 with a symbol of Br. It was first
discovered in 1826. In its elemental form, it is the diatomic molecule Br 2. At room
temperature, bromine is a reddish- brown liquid. Its oxidation states vary from -1, +1, 3, 4
and 5. Bromine is more reactive than iodine, but not as reactive as chlorine. Also,
bromine has two isotopes: 79Br and 81Br. Bromine consists of bromide salts, which have
been found in the sea. The world production of bromide has increased significantly over
the years, due to its access and longer existence. Like all of the other halogens, bromine
is an oxidizing agent, and is very toxic.

Iodine - Iodine has the atomic number 53 and symbol I. Iodine has oxidation states -1,
+1, 5 and 7. Iodine exists as a diatomic molecule, I 2, in its elemental state. At room
temperature, it appears as a violet solid. Iodine has one stable isotope: 127I. It was first
discovered in 1811 through the use of seaweed and sulfuric acid. Currently, iodide ions
can be isolated in seawater. Although iodine is not very soluble in water, the solubility
may increase if particular iodides are mixed in the solution. Iodine has many important
roles in life, including thyroid hormone production. This will be discussed in Part VI of
the text.

Astatine - Astatine is a radioactive element with an atomic number of 85 and symbol At.
Its possible oxidation states include: -1, +1, 3, 5 and 7. It is the only halogen that is not a
diatomic molecule and it appears as a black, metallic solid at room temperature. Astatine
is a very rare element, so there is not that much known about this element. In addition,
astatine has a very short radioactive half-life, no longer than a couple of hours. It was
discovered in 1940 by synthesis. Also, it is thought that astatine is similar to iodine.
However, these two elements are assumed to differ by their metallic character.

Table 1.1: Electron configurations of the halogens.

Halogen Electronic Configuration

Fluorine 1s2 2s2 2p5

Chlorine [Ne]3s2 3p5

Bromine [Ar]3d10 4s2 4p5

Iodine [Kr]4d10 5s2 5p5

Astatine [Xe]4f14 5d10 6s2 6p5

Periodic Trends
The periodic trends observed in the halogen group:

Melting and Boiling Points (increases down

the group)
The melting and boiling points increase down the group because of the van der Waals
forces. The size of the molecules increases down the group. This increase in size means
an increase in the strength of the van der Waals forces.

\[F < Cl < Br < I < At\]

 Table 1.2: Melting and Boiling Points of Halogens

Haloge Melting Point Boiling Point

n (˚C) (˚C)

Fluorine -220 -188

Chlorine -101 -35

Bromine -7.2 58.8

Iodine 114 184

Astatine 302 337

Atomic Radius (increases down the group)

The size of the nucleus increases down a group (F < Cl < Br < I < At) because the
numbers of protons and neutrons increase. In addition, more energy levels are added with
each period. This results in a larger orbital, and therefore a longer atomic radius.

Table 1.3: Atomic Radii of Halogens

Haloge Covalent Ionic (X-)

n Radius (pm) radius (pm)

Fluorine 71 133

Chlorine 99 181

Bromine 114 196

Iodine 133 220

 Table 1.3: Atomic Radii of Halogens

Haloge Covalent Ionic (X-)

n Radius (pm) radius (pm)

Astatine 150

Ionization Energy (decreases down the

group)
If the outer valence electrons are not near the nucleus, it does not take as much energy to
remove them. Therefore, the energy required to pull off the outermost electron is not as
high for the elements at the bottom of the group since there are more energy levels. Also,
the high ionization energy makes the element appear non-metallic. Iodine and astatine
display metallic properties, so ionization energy decreases down the group (At < I < Br <
Cl < F).

Table 1.4 Ionization Energy of

Halogens

Haloge First Ionization

n Energy (kJ/mol)

Fluorine 1681

Chlorine 1251

Bromine 1140

Iodine 1008

Astatine 890±40

Electronegativity (decreases down the

group)
The number of valence electrons in an atom increases down the group due to the increase
in energy levels at progressively lower levels. The electrons are progressively further
from the nucleus; therefore, the nucleus and the electrons are not as attracted to each
other. An increase in shielding is observed. Electronegativity therefore decreases down
the group (At < I < Br < Cl < F).

Table 1.5: Electronegativity of

Halogens

Haloge Electronegativi
n ty

Fluorine 4.0

Chlorine 3.0

Bromine 2.8

Iodine 2.5

Astatine 2.2

Electron Affinity (decreases down the

group)
Since the atomic size increases down the group, electron affinity generally decreases (At
< I < Br < F < Cl). An electron will not be as attracted to the nucleus, resulting in a low
electron affinity. However, fluorine has a lower electron affinity than chlorine. This can
be explained by the small size of fluorine, compared to chlorine.
 Table 1.6: Electron Affinity of
Halogens

Haloge Electron Affinity

n (kJ/mol)

Fluorine -328.0

Chlorine -349.0

Bromine -324.6

Iodine -295.2

Astatine -270.1

Reactivity of Elements (decreases down

the group)
The reactivities of the halogens decrease down the group ( At < I < Br < Cl < F). This is
due to the fact that atomic radius increases in size with an increase of electronic energy
levels. This lessens the attraction for valence electrons of other atoms, decreasing
reactivity. This decrease also occurs because electronegativity decreases down a group;
therefore, there is less electron "pulling." In addition, there is a decrease in oxidizing
ability down the group.

Hydrogen Halides and Halogen Oxoacids

Hydrogen Halides
A halide is formed when a halogen reacts with another, less electronegative element to
form a binary compound. Hydrogen, for example, reacts with halogens to form halides of
the form HX:

 Hydrogen Fluoride: HF
 Hydrogen Chloride: HCl
 Hydrogen Bromide: HBr
  Hydrogen Iodide: HI

Hydrogen halides readily dissolve in water to form hydrohalic

(hydrofluoric, hydrochloric, hydrobromic, hydroiodic) acids. The properties of these
acids are given below:

 The acids are formed by the following reaction: HX (aq) + H 2O (l) → X- (aq) +
H3O+ (aq)
 All hydrogen halides form strong acids, except HF
 The acidity of the hydrohalic acids increases as follows: HF < HCl < HBr < HI

Hydrofluoric acid can etch glass and certain inorganic fluorides over a long period of
time.

Table 1.7: States of Matter and Appearance of

Halogens

States of Matter (at Room Haloge Appearanc

Temperature) n e

Solid Iodine Violet

Astatine Black/Metallic [Assumed]

Liquid Bromine Reddish-Brown

Gas Fluorine Pale Yellow-Brown

Chlorine Pale Green

It may seem counterintuitive to say that HF is the weakest hydrohalic acid because
fluorine has the highest electronegativity. However, the H-F bond is very strong; if the H-
X bond is strong, the resulting acid is weak. A strong bond is determined by a short bond
length and a large bond dissociation energy. Of all the hydrogen halides, HF has the
shortest bond length and largest bond dissociation energy.

Halogen Oxoacids
A halogen oxoacid is an acid with hydrogen, oxygen, and halogen atoms. The acidity of
an oxoacid can be determined through analysis of the compound's structure. The halogen
oxoacids are given below:

 Hypochlorous Acid: HOCl

 Chlorous Acid: HClO2
 Chloric Acid: HClO3
 Perchloric Acid: HClO4
 Hypobromous Acid: HOBr
 Bromic Acid: HBrO3
 Perbromic Acid: HBrO4
 Hypoiodous Acid: HOI
 Iodic Acid: HIO3
 Metaperiodic Acid: HIO4; H5IO6

In each of these acids, the proton is bonded to an oxygen atom; therefore, comparing
proton bond lengths is not useful in this case. Instead, electronegativity is the dominant
factor in the oxoacid's acidity. Acidic strength increases with more oxygen atoms bound
to the central atom.

States of Matter at Room Temperature

Explanation for Appearance
The halogens' colors are results of the absorption of visible light by the molecules, which
causes electronic excitation. Fluorine absorbs violet light, and therefore appears light
yellow. Iodine, on the other hand, absorbs yellow light and appears violet (yellow and
violet are complementary colors, which can be determined using a color wheel). The
colors of the halogens grow darker down the group:

 Fluorine → pale yellow/brown

o http://www.daviddarling.info/images/fluorine.jpg
 Chlorine → pale green
o http://amazingrust.com/Experiments/how_to/Images/Chlorine_gas.jpg
 Bromine → red-brown
o http://www.crscientific.com/brominecell4.jpg
 Iodine → violet
o http://genchem.chem.wisc.edu/lab/PTL...ments/I/I.jpeg
 Astatine* → black/metallic
 o http://www4.msu.ac.th/satit/studentP...t/astatine.jpg

In closed containers, liquid bromine and solid iodine are in equilibrium with their vapors,
which can often be seen as colored gases. Although the color for astatine is unknown, it is
assumed that astatine must be darker than iodine's violet (i.e. black) based on the
preceding trend.

Oxidation States of Halogens in

Compounds
As a general rule, halogens usually have an oxidation state of -1. However, if the halogen
is bonded to oxygen or to another halogen, it can adopt different states: the -2 rule for
oxygen takes precedence over this rule; in the case of two different halogens bonded
together, the more electronegative atom takes precedence and adopts the -1 oxidation
state.

EXAMPLE 1.1: IODINE CHLORIDE (ICL)

Chlorine has an oxidation state of -1, and iodine will have an oxidation of +1. Chlorine is
more electronegative than iodine, therefore giving it the -1 oxidation state.

EXAMPLE 1.2: PERBROMIC ACID (HBRO4)

Oxygen has a total oxidation state of -8 (-2 charge x 4 atoms= -8 total charge). Hydrogen
has a total oxidation state of +1. Adding both of these values together, the total oxidation
state of the compound so far is -7. Since the final oxidation state of the compound must
be 0, bromine's oxidation state is +7.

One third exception to the rule is this: if a halogen exists in its elemental form (X 2), its
oxidation state is zero.
 Table 1.8: Oxidation States of Halogens

Haloge Oxidation States in

n Compounds

Fluorine (always) -1*

Chlorine -1, +1, +3, +5, +7

Bromine -1, +1, +3, +4, +5

Iodine -1, +1,+5, +7

Astatine -1, +1, +3, +5, +7

Applications of Halogens
Fluorine: Although fluorine is very reactive, it serves many industrial purposes. For
example, it is a key component of the plastic polytetrafluoroethylene (called Teflon-
TFE by the DuPont company) and certain other polymers, often referred to as
fluoropolymers. Chlorofluorocarbons (CFCs) are organic chemicals that were used as
refrigerants and propellants in aerosols before growing concerns about their possible
environmental impact led to their discontinued use. Hydrochlorofluorocarbons (HFCs)
are now used instead. Fluoride is also added to toothpaste and drinking water to help
reduce tooth decay. Fluorine also exists in the clay used in some ceramics. Fluorine is
associated with generating nuclear power as well. In addition, it is used to produce
fluoroquinolones, which are antibiotics. Below is a list of some of fluorine's important
inorganic compounds.

Table 1.9: Important Inorganic Compounds of Fluorine

Compou Uses
nd

Na3AlF6 Manufacture of aluminum

 Table 1.9: Important Inorganic Compounds of Fluorine

Compou Uses
nd

BF3 Catalyst

CaF2 Optical components, manufacture of HF,

metallurgical flux

ClF3 Fluorinating agent, reprocessing nuclear fuels

HF Manufacture of F2, AlF3, Na3AlF6, and fluorocarbons

LiF Ceramics manufacture, welding, and soldering

NaF Fluoridating water, dental prophylaxis, insecticide

SF6 Insulating gas for high-voltage electrical

equipment

SnF2 Manufacture of toothpaste

UF6 Manufacture of uranium fuel for nuclear reactors

Chlorine: Chlorine has many industrial uses. It is used to disinfect drinking water and
swimming pools. Sodium hypochlorite (NaClO) is the main component of bleach.
Hydrochloric acid, sometimes called muriatic acid, is a commonly used acid in industry
and laboratories. Chlorine is also present in polyvinyl chloride (PVC), and several other
polymers. PVC is used in wire insulation, pipes, and electronics. In addition, chlorine is
very useful in the pharmaceutical industry. Medicinal products containing chlorine are
used to treat infections, allergies, and diabetes. The neutralized form of hydrochloride is a
component of many medications. Chlorine is also used to sterilize hospital machinery and
limit infection growth. In agriculture, chlorine is a component of many commercial
pesticides: DDT (dichlorodiphenyltrichloroethane) was used as an agricultural
insecticide, but its use was discontinued.

Bromine: Bromine is used in flame retardants because of its fire-resistant properties. It

also found in the pesticide methyl bromide, which facilitates the storage of crops and
eliminates the spread of bacteria. However, the excessive use of methyl bromide has been
discontinued due to its impact on the ozone layer. Bromine is involved in gasoline
production as well. Other uses of bromine include the production of photography film,
the content in fire extinguishers, and drugs treating pneumonia and Alzheimer's disease.

Iodine: Iodine is important in the proper functioning of the thyroid gland of the body. If
the body does not receive adequate iodine, a goiter (enlarged thyroid gland) will form.
Table salt now contains iodine to help promote proper functioning of the thyroid
hormones. Iodine is also used as an antiseptic. Solutions used to clean open wounds
likely contain iodine, and it is commonly found in disinfectant sprays. In addition, silver
iodide is important for photography development.

Astatine: Because astatine is radioactive and rare, there are no proven uses for this
halogen element. However, there is speculation that this element could aid iodine in
regulating the thyroid hormones. Also, 211At has been used in mice to aid the study of
cancer.

References
1. Hill, Graham, and John Holman. Chemistry in Context. 5th ed. United Kingdom:
Nelson Thornes, 2000. 224-25.
2. Petrucci, Ralph H. Genereal Chemistry: Principles and Modern Applications. 9th
Ed. New Jersey: Pearson Education Inc, 2007. 920-928.
3. Verma, N.K., B. Kapila, and S.K. Khanna. Comprehensive Chemistry XII. New
Delhi: Laxmi Publications, 2007. 718-30.
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2. Atomic and Physical Properties of Halogens
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