B.sc 1st year Book
(Page 15)

Group VII A (17) Elements: F, Cl, Br, I, At Halogen family

Elements coming just before noble gases in the periodic table are known as “halogens“. The term halogen is derived from two Greek words ‘Halo’ and ‘genes’ which means salt and producers respectively because their compounds with metal behave as salts like table salt. (NaCl). The halogen group (Group 17) comprises the elements fluorine, chlorine, bromine, iodine, and artificially produced astatine (At) all belong to the Halogen family.
  1. Electronic configuration of halogen family :

All elements of group 17(VIIA) have seven electrons in their outermost shell thus the valence shell electronic configuration of these elements in the ground state is ns2np5 as shown below:

Table 4.14: The electronic configuration of group VIIA elements

Element Symbol Atomic No. Electronic configuration with an inert gas core
Fluorine F 9 [He]2s2,2p5
Chlorine Cl 17 [Ne]3 s2,3p5
Bromine Br 35 [Ar]3 d6,4 s2,4p5
Iodine I 53 [Kr]4 d10,5 s2,5p5
Astatine Al 85 [Xe]41414,5 d10,6 s2,6p5

 

1- Physical state of halogen family :

Halogens(Halogen family) exist as covalent diatomic molecules (X2) under ordinary conditions. These discrete X2 molecules are held together by weak van der Waals forces which explain the volatile nature of these elements. Fluorine and chlorine are gases with pale yellow and greenish-yellow colors respectively. Bromine is a deep reddish-brown liquid with high vapor pressure. Iodine is a lustrous greyish black crystalline solid which sublimes readily, when heated it gives deep violet vapor.

2- Atomic and Ionic radii of halogen family :

Atomic radii of these elements are larger than corresponding noble gases and smaller than the elements of group VIA. This property increases on moving down the group with an increase of atomic number. Thus, fluorine has a smaller atomic size than the next member of this group. The actual order of the atomic size for the element of group VIIA is :
F<Cl<Br<1<Al
Similarly, ionic radii of halide ions (X−1 increases from F− to F).
3- Metallic property of halogen family :
Although all the halogens are nonmetals, their metallic properties increase on passing from top to bottom in the group i.e, in the order F<Cl<Br<1.
General studies of s and p block elements.
Some important physical properties of the elements of group
4- Electron affinity of halogen family :
Halogens have a great tendency to accept electrons. Hence their electron affinity values are very large. Chlorine has the highest electron affinity value among all the elements. The actual order of electron affinity values for the elements of group VII A is.
F<Cl>Br>I>Al or Cl>F>Br>1>At
Reason :
Because of the very small atomic size of fluorine the incoming electron experiences: a greater repulsive force due to electron-electron repulsion from the already present electrons than in chlorine. In addition, the size of F-ion is almost double that of F-atom. In other words, due to the entry of one electron, the remaining electrons of 2p-orbitals are moved away from the nucleus, which is possible by the absorption of energy. Thus, a part of the energy released by adding an electron to a fluorine atom is utilized in expanding its size. This results, in giving a lower electron activity value. Whereas in other halogens increase in size is relatively less. Hence in the formation of X – ion, less energy is absorbed.
5- Electronegativity:

Since the halogens have the least atomic size and greater nuclear charge in their respective periods, they have the maximum tendency to attract the electrons towards them and hence have the maximum electronegativities in their respective periods. Fluorine has the highest electronegativity value amongst all known elements mentioned in the periodic table. In group VII A this property decreases gradually on moving down the group due to the increase in the atomic size. The gradual decrease of electronegativity values down the group indicates a gradual decrease in the nonmetallic character of halogens from fluorine to iodine. Consequently, fluorine shows only nonmetallic characters white the last element Iodine (i) shows both non-metallic and metallic characteristics.

6- ionization energy:
As halogens are placed on the extreme right-hand side of the periodic table, their ionization energy values become very large because of increased nuclear charge and decreased atomic size in their respective periods. The high values of l.E. indicate that these elements have very little tendency to lose electrons. On going down the group from F to 1, the ionization energy decrease with the increase in atomic size.
Thus, the last member iodine has a comparatively very low value of ionization energy and hence has a tendency to lose one or more electrons to form tho cations like 11,13+,15+. But fluorine can never form its cation.
7- Oxidation states :
The fluorine atom hers no vacant d-orbitals in its valance shell and can not, therefore, have any exciting sales. Hence, it can not have any higher oxidation states. Also, fluorine is most of the electronegative element which is why only one oxidation state (−1) is exhibited by it.
Whereas rest elements of group VII A exhibit more than one oxidation state due to the presence of empty d-orbitals in their valence shell as shown below:
Element Oxidation number
F ; −1
Cl ; −1 +1 +3 +4 +5 +6 +7
Br ; −1 +1 +3 +4 +5 +6
I ; −1 +1 +3 +5 +7

 

The higher oxidation states of halogens will be stabilized only when the other elements can bring them to properly excited states. Oxygen and fluorine are such elements because of their small atomic sizes they have few steric repulsions e.g. CiO2−(+3), BrO3−(+5)F5(+5), IF (+7), etc. In heavier halogens, because of the small energy difference between np and nd levels electrons from the np level can more easily be transferred to and level than in chlorine. Moreover, Br and I of bigger atomic sizes can attach more F-atoms than Cl. This is the reason why I can form stable IF while Cl can not give such a compound: In the case of Br the penultimate shell I.e. ( n−1) shell, is weakly screened which is why the energy required to promote an electron from the s-orbital to the vacant d-orbitals is markedly higher than in case of chlorine. This results in Br showing an inability to attain the +7 oxidation state.
The oxidation states of +4 and +6 occur in oxides and oxyhalides of chlorine and bromine only.
8- Reactivity and oxidizing power:

Halogens are very reactive elements and their reactivity decreases with an increase in atomic size. Fluorine is the most reactive of all the known elements. It is only the halogen that forms binary compounds with noble gases such as XeF2, and XeF4. The reactivity of

halogens can be explained on the basis of their low dissociation energy.
Halogen F2 Cl2 Br2 I2
Dissociation energy ( Kcal mole −1) :38.0 57.0 53.4 51.0

 

Since halogens are highly electronegative and have the highest values of electron affinity, they have a tendency to accept electrons from their elements more strongly and act as strong oxidizing agents. The greater the accepting property, the greater will be the oxidizing power. Fluorine being the most electronegative element is also the most powerful oxidizing agent. This property decreases very fast as we move from F2 to I2. It is evident from their standard electrode oxidation potential values given below:

Elements F2 Cl2 Br2 I2
Std. oxidation electrode potential (eV) (E22ax) +2.87 +1.36 +1.09 +0.54

 

Note: The very high oxidation potential value for fluorine is due to very high solvation energy.

9- Hydration energy of halide ions (X−):

Hydration energy is defined as the amount of energy released in converting gaseous halide ion (Xg→ into hydrated halide ion (Xaq−1).
X(g)−+H2O(1)⟶X((aq) −+Hydration energy 
The hydration energies of halide ions decrease as the size of X− ions, increases. Thus, F−– ion has the highest value of hydration energy and the I ion has the lowest value as clear from the table (4.15).

10- Heat of dissociation of X2 molecules or (dissociation energy);

The heat of dissociation, (DHdis) of X2 molecule may be defined as the amount of energy required to dissociate X2 molecule in the gaseous state into gaseous X-atom”.or “the amount of energy required to break X−X bond in X2 molecule to get free X-atoms:
X2(g)+ Energy ⟶2X(g) or X2(g)+ DH dis ⟶2X(g)
The heat of dissociation of X2 molecule is very low as shown above.
It is evident from the DHdis values of halogen molecules (X2) that the F2 molecule has the lowest value of DH-dis as compared to the Cl2 molecule and as we move down the group, DHdis decreases from Cl2 to I2. The lower value of heat of dissociation of F2 molecule than Cl2 molecule is due to the following reasons:
0 FF−F bond distance in F2 molecule is smaller than C−Cl bond distance in Cl2 molecule which results in repulsion between non-bonding electrons is very large in F2 molecule. This makes the F-atoms in the F2 molecule repel each other and promote the dissociation of the F2 molecule into F-atoms.
(ii) X−X bond in Cl2, Br2, and I2 molecules is stronger than the F−F bond in the F2 molecule. This is due to the possibility of the existence of multiple bonding in X2 molecules other than F2.
11- Hydrogen halides:

Halogens combine with hydrogen to form covalent molecular species HX, known as hydrogen halides. The acid strength increases from HF to HI. This property depends upon how much easier the acid is able to give protons from the solvent molecules. In HF because very low dissociation energy and great ionic character of the bond, intermolecular hydrogen bonding occur in it. As a result, its bond energy is more than other H−X bords of hydrogen. halides. Hence H−F bond breaks with very difficulty and HF behaves as a weak acid. The acid strength will increase in the order:

HF < HCl < HBr < HI
It is evident from their pKa values.
Haloacida HF HCl HBr HI
pKa value: 32 −7.0 9.5 10.0

 

12- Halides :
When halogen molecules such as F2, Cl2, Br2, and I2 react with an aqueous solution of the alkalies, give halides (X−1, hypo-halide (XO−), or X− and halides (XO3−)depending upon the condition required.
2F2+2NaOH(2)⟶2NaF+OF2+H₂O
Cl2+2NaOH (ail) ⟶ColdNaCl+NaClO+H2O
Cl2+6NaOH (canc) ⟶Hot=5NaCl+NaClO3+H2O
The bond energy of binary halides is proportional to bond order and partial charge on halogen and is inversely proportional to bond length. ionic halides having a large negative charge on halogen atoms are less volatile and have high melting points. This is the reason why metal halides are emotions while non-metal halides are liquid or gases. In metal halides, metal theories are more ionic and have high m.p. and boiling points because of more electronegative difference in M and F. While metal Ioides are least ionic and more covalent Thus, for sodium halides, the decreasing order of ionic character is:
NaF  > NaCl > NaBr > NaI
13- Oxides :
Halogens form many binary compounds with oxygen but most of them are unstable. Fluorine forms two binary compounds with oxygen, OF2 and O2F2 which are properly called ‘oxygen fluorides‘ because of the higher electronegativity of fluorine, Both are strong reducing agents. Other halogens show many oxidation states and hence form many oxides in which the oxidation state of halogen varies from +1 to +7. as shown below :
Table 4.16: Oxides of VIIA group elements
Elements Oxidation states
−1 +1 +4 +5 +6 +7 others
F OF2 O2F2
Cl Cl2O ClO2 Cl2O2 Cl2O7
Br Br2O BrO2 BrO3 Br3O8

 

Mostly halogen oxides are gases or low melting solids, weak van der Waal forces exist between their molecules. They are strong oxidizing agents. H is because-1 oxidation of halogen is most stable as compared to +1 to +7.

In addition, these oxides are more acidic, their acidic strength increases with the increase of the oxidation state of the halogen atom. The dissociation energy of the X−O bond is small, hence halogen oxides are unstable compounds, and ClO2 dissociates with an explosion.

14- Oxy acids :

Fluorine forms only one oxy acid HOF known as fluoric acid or hypofluorous acid because of its high electronegativity. On the other hand, since the remaining halogens are less electronegative than oxygen, they form several oxyacids. Most of them can not be isolated in pure form. They are stable only in aqueous solution or in the form of their salts. The main oxyacids of halogens are given in table 4.17.

The strength of oxy acids increases appreciably with increased halogen as evident from their pKa values, shown below.

Oxy acids: HOClHClO2HClO3HClO4 
pKa value: 7.52.0−1.2−10.0 

The strength of halide acids with the same oxidation state of the halogens decreases in the sequence: >Br>Cl.

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