We shall learn P- Block Elements Chemistry 12th Notes of group 17. In this article, we shall discuss halogen family which includes Fluorine, Chlorine, Bromine, Iodine and Astatine. All the important topics regarding this group elements have been discussed in a short form so that students can easily revise this chapter and score a good marks in their examinations. This part of inorganic chemistry is very important as we know that the halogens are highly reactive non- metallic elements like group 1 and 2. The elements of group 17 show similarities among themselves very much w.r.t. to the elements of other groups. Also, there is a regular gradation in their physical and chemical properties.
P-Block Elements Chemistry 12th Notes of group 17 :
(1) Occurrence: Fluorine and chlorine are fairly abundant while Bromine and iodine less so. Fluorine is present as insoluble fluorides and fluoroapatite. Sea water is the main source of chlorine bromine and Iodine as halides of them.
(2) Electronic Configuration: All the elements of this group has one electron short of the next noble gas. Hence, their outermost electronic Configuration is ns2np5.
(3) Atomic and Ionic radii: Halogen atoms have the smallest atomic radius in their respective periods due to maximum effective nuclear charge. Atomic and Ionic radii increase from Fluorine to iodine due to increase in number of shells.
(4) Ionisation enthalpy: They have little tendency to loss electrons. Thus they have very high ionisation enthalpy. Due to increase in atomic size, Ionisation enthalpy decreases down the group.
(5) Electron gain enthalpy: Halogen atoms have maximum negative electron gain enthalpy in the corresponding periods. This is due to the fact that the atoms of these elements have only one electron less than stable noble gas Configuration. Electron gain enthalpy of the elements of this group becomes less negative down the group. However, the negative gain enthalpy of Fluorine is less than that of chlorine due to small size of fluorine atoms.
(6) Electronegativity: They have very high electronegativity. The electronegativity decreases down the group. Fluorine is the most electronegative elements in the periodic table.
(7) Physical properties: Fluorine and chlorine are gases, bromine is a liquid and Iodine is solid. Their m.p. and b.p. steadily increases with atomic number. All halogens are coloured due to absorption of radiation in visible regions which results in the excitation of outer electrons to higher energy levels. F2 has yellow, Cl2 has greenish yellow, Br2 has red and I2 violet Colour. Fluorine and chlorine react with water, bromine and Iodine are sparingly soluble in water but are soluble in various organic solvents and give coloured solutions.
(8) Bond dissociation enthalpy of F2 is smaller than Cl2 and further from chlorine onwards show the expected decrease. A reason for this anomaly is the relatively large electron-electron repulsion among the lone pairs in F2 molecules.
The final order of dissociation enthalpy is:- Cl2 > Br2 > F2 > I2.
Chemical properties of group 17 elements:
(1) Oxidation state: All the halogens show -1 oxidation state. However, chlorine ,bromine and iodine show +1, +3, +5 and +7 oxidation state also. The higher oxidation state of chlorine, bromine and Iodine are seen when these elements form intrhalogens, oxides and oxaacids. Fluorine atoms has no d-orbitals in its valence shell and therefore can’t expands its octet. It shows only -1 oxidation state.
(2) chemical reactivity: All the halogens are highly reactive. They reacts with metals and nonmetals to form halides. The reactivity of the halogens decreases down the group. F2 > Cl2 > Br2 > I2.
(3) Oxidising nature: F2 is the strongest oxidising halogens and it oxidises other halide ions in solution and even in solid phase. The oxidising ability decreases down the group. The relative oxidising power of halogens can further be explained by their reaction with water. Fluorine oxidise water in oxygen whereas chlorine and bromine form corresponding hydrohalic and hypohalous acid. Iodine does not react with water.
(4) Anomalous behaviour of fluorine: The anomalous behaviour of fluorine is due to its small size, highest electronegativity, low bond dissociation enthalpy and non availability of d-orbitals in valence shell. Most of the reactions of Fluorine are exothermic. It forms only one oxaacids while other members form a number of oxaacids. Hydrogen fluoride is liquid due to strong hydrogen bond other halogen halides are gases.
Reactivity towards hydrogen: They all form hydrogen halide with hydrogen but their affinity for hydrogen decreases from fluorine to Iodine. Hydrogen halides dissolve in water to form hydrohalic acid. The acidic strength of Hydro halic acids are HF < HCl < HBr < HI.
Reactivity towards Oxygen: Halogens form many oxides with oxygen but most of them are not stable .
Fluorine forms two Oxides as OF2 and O2F2. But only OF2 is thermally stable at 298K. Both are strong fluorinating agents . O2F2 oxidises plutonium to PuF6 and this reaction is used in removing plutonium as PtF6 from spent nuclear fuel.
Chlorine forms oxides as Cl2O, ClO2, Cl2O6 and Cl2O7. They are highly reactive oxidising agent and tend to explode.
Bromine forms Br2O, BrO2 and BrO3 oxides. They are least stable and very powerful oxidising agents.
Iodine forms oxides as I2O4, I2O5, I2O7. They are insoluble solids and decompose on heating. I2O5 is a very good oxidising agent and used in the estimation of CO.
Reactivity towards metals: Halogens react with metals to form metal halides. The Ionic characters of the halides decreases in the order MF > MCl > MBr > MI. Where M is a monovalent metal. If a metal exhibit more than one oxidation state, the halides in higher oxidation state will be more covalent than the one in lower oxidation state.
Reactivity of halogens towards other halogens: Halogens combine amongst themselves to form a number of compounds known as interhalogens. They are XX’1, XX’3, XX’5 and XX’7 where X is a larger size halogen and X’ are smaller size halogens.
Structure of Halogens Oxaacids
Preparation of Chlorine:
(1) By heating manganese dioxide with conc. HCl
MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O
In place of HCl, a mixture of common salt and conc. H2SO4 can be used.
4NaCl + MnO2 + 4H2SO4 → MnCl2 + 4NaHSO4 + 2H2O + Cl2
(2) By the action of HCl on KMnO4
2KMnO4 + 16HCl → 2KCl + 2MnCl2 + 8H2O + 5Cl2
(3) By Deacon process, hydrogen chloride gas is allowed to oxidise by atmospheric oxygen in the presence of CuCl2 catalyst at 723K.
4HCl + O2 → 2Cl2 + 2H2O
(4) Chlorine is also obtained by electrolysis of brine solution (conc. NaCl solution). In which chlorine is liberated at anode.
Properties of chlorine:
(1) It is a greenish yellow gas with pungent and suffocating smell.
(2) It is about 2.5 times heavier than air and can be liquified easily into greenish yellow liquid which boils at 239K. It is soluble in water.
(3) Chlorine reacts with a number of metals and nonmetals to form chlorides.
2Na + Cl2 → 2NaCl
2FeCl2 + 3Cl2 → 2FeCl3
P4 + 6Cl2 → 4PCl3
S8 + 4Cl2 → 4S2Cl2
(4) It has great affinity for hydrogen. It reacts with compounds having hydrogen to form HCl
H2S + Cl2 → 2HCl + S
C10H16 + 8 Cl2 → 16HCl + 10C
(5) With excess of ammonia, it gives nitrogen and ammonium chloride whereas with excess chlorine nitrogen trichloride (explosive) is formed.
8NH3 + 3Cl2 → 6NH4Cl + N2
NH3 + 3Cl2 → NCl3 + 3HCl
(6) With cold and dilute alkalies chlorine produces a mixture of chloride and hypochlorite but with hot and conc. alkalies it gives chloride and chlorate.
2NaOH + Cl2 → NaCl + NaOCl + H2O
6NaOH + 3Cl2 → 5NaCl + NaOCl3 + 3H2O
With dry slaked lime, it gives bleaching powder
2Ca(OH)2 + 2Cl2 → CaOCl2 + CaCl2 + 2H2O
(7) chlorine oxidises ferrous to ferric, sulphite to sulphate, sulphur dioxide to sulphuric acid and Iodine to Iodic acid.
2FeSO4 + H2SO4 + Cl2 → Fe2(SO4)3 + 2HCl
Na2SO3 + Cl2 + H2O → Na2SO4 + 2HCl
SO2 + 2H2O + Cl2 → H2SO4 +2HCl
I2 + 6H2O + 5Cl2 → 2HIO3 + 10HCl
(8) It is a powerful bleaching agent due to oxidation behaviour.
Cl2 + H2O → 2HCl + O
Coloured substance + O → Colourless substance.
Preparation of Hydrogen chloride:
In laboratory, it is prepared by heating Sodium chloride with conc. Sulphuric acid.
NaCl + H2SO4 → NaHSO4 + HCl
NaHSO4 + NaCl → Na2SO4 + HCl
Properties of Hydrogen chloride:
(1) It is a colourless and pungent smelling gas. It is easily liquified to a colourless liquid (b.p.189K) and freezes to a white crystalline solid (f.p. 159K). It is extremely soluble in water and ionised
HCl(g) + H2O (l) → H3O+(aq) + Cl–(aq) Ka= 107
Its aqueous solution is called hydrochloric acid.
(2) It reacts with ammonia and gives white fumes of NH4Cl
NH3 + HCl → NH4Cl
(3) when three parts of Conc.HCl and one part of conc. HNO3 are mixed, aqua regia is formed which is used for dissolving noble metals as Gold and platinum.
Au+ + 4H+ + NO3– + 4Cl– → AuCl4– + NO + 2H2O
3Pt +16 H+ + 4NO3– + 18Cl– → 3PtCl62- + 4NO + 8H2O
(4) Hydrochloric acid decomposes salts of weaker acids.
Na2CO3 + 2HCl → 2NaCl + H2O + CO2
NaHCO3 + HCl → NaCl + H2O + CO2
Na2SO3 + 2HCl → 2NaCl + H2O + SO2