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ChemWiki: The Dynamic Chemistry E-textbook > Inorganic Chemistry > Descriptive Chemistry > Main Group Elements > Group 17: The Halogens

Group 17: The Halogens

Figure 1.1 A visual to place the halogens relative to the periodic table. (Not drawn to scale.)

The halogens can be found on the left-hand side of the noble gases. These five toxic, non-metallic elements make up Group 17 of the periodic table and consist of: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Although astatine is radioactive and only has short-lived isotopes, it behaves similar to iodine and is often included in the halogen group. Since the halogen elements have seven valence electrons, they only require one additional electron to form a full octet. This characteristic makes them more reactive than other non-metal groups.

 

Introduction

Halogens form diatomic molecules (X2 where X refers to a halogen atom) when they are in their pure 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 of thumb, fluorine is the most reactive halogen and astatine is the least reactive. All halogens form sodium salts that have similar properties. Halogens are present as halide anions with a negative one charge (i.e. Cl-, Br-, etc.). Replacing the -ine ending with an -ide ending forms the names for these halide anions. (For example, Cl- becomes chloride.) In addition, halogens serve as oxidizing agents, because 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 sometimes 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- for specific halogen substitutions. Halogen elements could cross-bond with each to form diatomic molecules with polar covalent single bonds.

Chlorine (Cl2) was the first halogen to be discovered in 1774, followed by iodine (I2), bromine (Br2), fluorine (F2), and astatine (At), which was discovered last in 1940. The name "halogen" is derived from the Greek roots hal- (which means "salt") and -gen (which means "to form"). Together these words combine to mean "salt former", which is in reference to the fact that halogens form salts when they react with metals. Halite is the mineral name for rock salt, which is a natural mineral consisting essentially of sodium chloride (NaCl). Lastly, the halogens are also relevant to real-life, whether it be the fluoride that goes in toothpaste, the chlorine that disinfects drinking water, or the iodine that is responsible for the production of thyroid hormones in one's body.

Figure 1.2 Shows all of the halogen elements in respect to its mass and atomic numbers.

Key:
Mass
number
 
Element
 
Atomic Number

 

18.998

F

9

35.453

Cl

17

79.904

Br

35

126.90

I

53

[210]

At

85

Elements

  1. Fluorine- Fluorine has an atomic number of 9 and a symbol of F. Elemental fluorine was first discovered in 1886 by isolating it from hydrofluoric acid. Fluorine exists as a diatomic molecule in its free state (F2) and is the most abundant halogen found in the Earth's crust. Fluorine is the most electronegative element out of all of the elements 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, with the exception of its free state (where the oxidation state is zero). Fluorine is extremely reactive and reacts directly with all elements except helium (He), neon (Ne) and argon (Ar). In H2O solution, hydrofluoric acid (HF) is a weak acid. Although fluorine is highly electronegative, its electronegativity does not quite determine its acidity. HF is a weak acid due to the fact that the fluoride ion is basic (pH of >7). In addition, fluorine produces very powerful oxidants. For example, fluorine can react with the noble gas xenon and form the strong agent, Xenon Difluoride (XeF2). There are many uses for fluorine, which will be discussed later on in Part VI.
  2. 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 Cl2 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. 
  3. 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 Br2. 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.
  4. 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, I2, 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. 
  5. 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 4s4p5

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

Halogen Melting Point (˚C) Boiling Point (˚C)
Fluorine -220 -188
Chlorine -101 -35
Bromine -7.2 58.8
Iodine 114 184
Astatine 302 337

Atomic Radii: increases down the group

The size of the nucleus increases down a group because there is a higher number of protons and neutrons. Also, more energy levels are added on after passing each period. This results in a bigger orbital, and therefore a bigger radius.

F < Cl < Br < I < At

Table 1.3 Atomic Radii of Halogens

Halogen Covalent Radius (pm) Ionic (X-) radius (pm)
Fluorine 71 133
Chlorine 99 181
Bromine 114 196
Iodine 133 220
Astatine 150  

Ionization Energy: decreases down the group

If the outer valence electrons are not near the nucleus, it will not take as much energy to remove them. Therefore, the energy required to pull off the outermost electron will not be 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 should decrease down the group.

At < I < Br < Cl < F

Table 1.4 Ionization Energy of Halogens

Halogen First Ionization Energy (kJ/mol)
Fluorine 1681
Chlorine 1251
Bromine 1140
Iodine 1008
Astatine 890±40

Electronegativity: decreases down the group

The number of valence electrons increases due to the increase in energy levels as the elements progress down the group. The electrons are not as near to the nucleus anymore. Therefore, the nucleus and the electrons are not as attracted to each other as much. An increase in shielding is observed. Electronegativity will therefore decrease down the group.

At < I < Br < Cl < F

Table 1.5 Electronegativity of Halogens

Halogen Electronegativity
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 will decrease. 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.

At < I < Br < F < Cl

Table 1.6 Electron Affinity of Halogens

Halogen Electron Affinity (kJ/mol)
Fluorine -328.0
Chlorine -349.0
Bromine -324.6
Iodine -295.2
Astatine -270.1

Reactivity of Elementsdecreases down the group

Down a group, the atomic radius gets bigger, and there is an increase in the amount of energy levels. This results in less attraction of the valence electrons. It also decreases because electronegativity decreases down a group, which means that there will be less interactions with the electrons in terms of "pulling". Also, since there is a decrease in oxidizing ability down a group, the reactivity of the elements will decrease as well.

 At < I < Br < Cl < F

Hydrogen Halides and Halogen Oxoacids

Hydrogen Halides

Halide- Halogen reacts with another element [which does not exceed that halogen's electronegativity] to form a binary compound. More relevant to look at the hydrogen halides; a halogen reacts with hydrogen to form a halide "HX" (where X refers to a halogen element):

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

These hydrogen halide compounds can readily dissolve in water to form hydrohalic (hydrofluoric, hydrochloric, hydrobromic, hydroiodic) acids in a range of concentrations. Acidic strength of the hydrogen halides:

  • Based on the following reaction: HX (aq) + H2O (l) X(aq) + H3O+ (aq) (where X is a halogen)
  • all hydrogen halides are strong acids, except HF
  • HF < HCl < HBr < HI [HF is the weakest acid]

Hydrofluoric acid can etch glass, as well as certain inorganic fluorides over a long period of time.

  • Question: Why is HF the weakest acid when fluorine has such a high electronegativity?
  • Answer: If the H-X bond is strong, the resulting acid will be weak. A strong bond is determined by a short bond length and a large bond dissociation energy. Out of all of the following hydrogen halides, HF has the shortest bond length and biggest bond dissociation energy.

Halogen Oxoacids 

A halogen oxoacid is an acid which has hydrogen, oxygen, and halogen atoms. Halogen oxoacids demonstrate that the acidity can be found by analyzing their structures. The halogen oxoacids consist of the following:

  • 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

When finding the acidity of the oxoacid, the length of the bond is no longer a factor. This is due to the fact that for each acid, there is a similar foundation: a hydrogen atom is bonded to an oxygen atom. Electronegativity becomes the key factor in figuring out the acidity of the oxoacid. Acidic strength will increase depending on the number of oxygen atoms bound to the central atom.

States of Matter at Room Temperature

Table 1.7 States of Matter and Appearance of Halogens

States of Matter (at Room Temperature)

Halogen

 Appearance

Solid

 Iodine

Violet

 

 Astatine

Black/Metallic [Assumed]

Liquid

 Bromine

Reddish-Brown

Gas

 Fluorine

Pale Yellow-Brown

 

Chlorine

Pale Green

Explanation for Their Appearance

The color of the halogens is based on the absorption of visible light by the molecules, which leads to excitation. Fluorine absorbs violet light, and therefore appears light yellow. Iodine, on the other hand, absorbs yellow light and appears violet (color wheel). This shows that the color of the halogens becomes darker down the group:

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 violet (iodine), i.e. black, based on the periodic trend for colors observed in the halogens.

V. Oxidation States of Halogens in Compounds

Rule: Halogens have an oxidation state of -1. Although if the halogen is bonded to oxygen or to another halogen that is more electronegative, its oxidation state will not be -1.

Example  

  • Example 1: Iodine Chloride (ICl):
    • Chlorine will have 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 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 will be +7.

Another Exception: If fluorine exists in its elemental form (F2), its oxidation state is zero.


Table 1.8 Oxidation States of Halogens

Halogen Oxidation States in 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

 

Example  

Question: Why does fluorine always have an oxidation state of-1 in its compounds?

  • Answer: Electronegativity increases across a period, and decreases down a group. Therefore, fluorine has the highest electronegativity out of all of the elements because of its position on the periodic table. Its electron configuration is 1s 2s2 2p5. If fluorine gains one more electron, the outermost p orbitals will all be filled (resulting in a full octet). Since fluorine has a high electronegativity, it can easily pull off that one desired electron from a nearby atom. Fluorine would then have a similar electron configuration as a noble gas (eight valence electrons) since all of its outermost orbitals would be filled. This also makes fluorine a stable element, which is preferred.


Applications of Halogens

Fluorine: Although fluorine is very reactive, it still serves many purposes in the real world. 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 prior to the discovery that they were depleting the ozone layer in the atmosphere, especially in polar regions of the earth. They were used as refrigerants, as well as propellants in aerosols. Hydrochlorofluorocarbons (HFCs) were then used as an alternative. Fluoride is used in other compounds to produce many results. For example, it is used in toothpaste and water to help reduce tooth decay. Fluorine also exists in clay, which is used to make some ceramics. Fluorine is associated with generating nuclear power as well. In addition, fluorine 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

Compound Uses
Na3AlF6 Manufacture of aluminum
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 uses in modern-day life. For example, chlorine is used to disinfect drinking water as well as swimming pools. Compounds which generate chlorine in water are also used, such as sodium hypochlorite (NaClO) in bleach. Hydrochloric acid (a solution of hydrogen chloride, HCl, dissolved in water), sometimes referred to as muriatic acid, is a very commonly used acid in industry and laboratories. Chlorine is also present in polyvinyl chloride (PVC), and a couple other polymers. PVC is used in wire insulation, pipes, as well as in electronics. In addition, chlorine is very useful in relation to its use in medicine. For example, medicinal products that contain chlorine help fight infections, allergies, and diabetes. Various compounds used in medicine are used as the neutralized hydrochloride form. Chlorine is also used to help sterilize the machines in hospitals, limiting the growth of infections in patients. Chlorine is used in agriculture as well. For example, there are many pesticides which contain chlorine. DDT (dichlorodiphenyltrichloroethane) was used as an agricultural insecticide, but use was discontinued.

Bromine: Bromine is often used in flame retardants because of its ability to resist fire. They are also found in the pesticide, methyl bromide, which can help store crops and eliminate the spread of bacteria. However, the excessive use of methyl bromide has been shown to destroy the ozone layer. Due to this realization, this pesticide was eventually banned because of its negative effects on the environment. Bromine is involved in gasoline production as well. Other uses of bromine include: the production of photography film, the content in fire extinguishers, as well as in particular drugs to treat pneumonia and Alzheimer's disease.

Iodine: Iodine is important in the proper functioning of the thyroid gland in one's 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 (kills germs). Solutions used to clean open wounds will most likely contain iodine in them. Iodine is also found in disinfectant sprays. In addition, silver iodide is important for photography development.

Astatine: Since astatine is radioactive, and not common, there are no proven uses for this halogen element. However, there is speculation that this element could possibly aid iodine in regulating the thyroid hormones. Also, astatine 211 has been used in mice to aid the study of cancer.

VII. Outside Links

Practice Problems

1. Why does fluorine always have an oxidation state of -1 in its compounds?

(please highlight to see text) Answer: Electronegativity increases across a period, and decreases down a group. Therefore, fluorine has the highest electronegativity out of all of the elements. Since fluorine has seven valence electrons, it only needs one more electron to acheive a noble gas configuration (eight valence electrons). Therefore, it will be more likely to pull off an electron from a nearby atom.

2.  Find the oxidation state of the halogen in each problem:
a. HOCl  
b. KIO3
c. F2
 

Answers: 

  • a. +1 (Hydrogen has an oxidation state of +1, and oxygen has an oxidation state of -2. Therefore, chlorine must have an oxidation state of +1 so that the total charge can be zero)  
  • b. +5  (Potassium's oxidation state is +1. Oxygen has an oxidation state of -2, so for this compound it is -6 (-2 charge x 3 atoms= -6). Since the total oxidation state has to be zero, iodine's oxidation state must be +5).  
  • c. 0 (Elemental forms always have an oxidation state of 0.)

3. What are three uses of chlorine?
 

Answer: disinfecting water, pesticides, and medicinal products

4. Which element(s) exist(s) as a solid in room temperature?
 

Answer: iodine and astatine

5. Do the following increase or decrease down the group of halogens?
a. boiling point and melting point
b. electronegativity
c. ionization energy

 
 Answer: a. increases b. decreases c. decreases

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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