Group 17 Elements - Halogen Family

GROUP 17 ELEMENTS - HALOGEN FAMILY  

 Fluorine, chlorine, bromine, iodine, astatine and tennessine are members of Group 17. These are collectively known as the halogens.. The halogens are highly reactive non-metallic elements. Like Groups 1 and 2, the elements of Group 17 show great similarity amongst themselves. That much similarity is not found in the elements of other groups of the periodic table. Also, there is a regular gradation in their physical and chemical properties. Astatine and tennessine are radioactive elements.

 

Occurrence of group 17 elements     

Fluorine and chlorine are fairly abundant while bromine and iodine less so. Fluorine is present mainly as insoluble fluorides (fluorspar CaF2, cryolite Na3AlF6 and fluoroapatite 3Ca3(PO4)2.CaF2) and small quantities are present in soil, river water plants and bones and teeth of animals. Sea water contains chlorides, bromides and iodides of sodium, potassium, magnesium and calcium, but is mainly sodium chloride solution (2.5% by mass). The deposits of dried up seas contain these compounds, e.g., sodium chloride and carnallite, KCl.MgCl2.6H2O.

Certain forms of marine life contain iodine in their systems; various seaweeds, for example, contain upto 0.5% of iodine and Chile saltpetre contains upto 0.2% of sodium iodate.

 

Tennessine is a synthetic radioactive element. Its symbol is Ts, atomic number 117, atomic mass 294 and electronic configuration [Rn] 5f146d107s27p5. Only very small amount of the element could be prepared. Also its half life is in milliseconds only.

 

Electronic Configuration of group 17 elements    

All these elements have seven electrons in their outermost shell (ns2np5) which is one electron short of the next noble gas.

 

Atomic and Ionic Radii of group 17 elements        

The halogens have the smallest atomic radii in their respective periods due to maximum effective nuclear charge. The atomic radius of fluorine like the other elements of second period is extremely small. Atomic and ionic radii increase from fluorine to iodine due to increasing number of quantum shells.

 

Ionisation enthalpy of group 17 elements  

They have little tendency to lose electron. Thus they have very high ionisation enthalpy. Due to increase in atomic size, ionisation enthalpy decreases down the group.

 

Electron gain enthalpy of group 17 elements 

Halogens 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 configurations. Electron gain enthalpy of the elements of the group becomes less negative down the group. However, the negative electron gain enthalpy of fluorine is less than that of chlorine. It is due to small size of fluorine atom. As a result, there are strong interelectronic repulsions in the relatively small 2p orbitals of fluorine and thus, the incoming electron does not experience much attraction.

 

Physical Properties of group 17 elements  

 

Halogens display smooth variations in their physical properties. Fluorine and chlorine are gases, bromine is a liquid and iodine is a solid. Their melting and boiling points steadily increase with atomic number. All halogens are coloured. This is due to absorption of radiations in visible region which results in the excitation of outer electrons to higher energy level. By absorbing different quanta of radiation, they display different colours. For example, F2, has yellow, Cl, greenish yellow, Br2, red and I2, violet colour. Fluorine and chlorine react with water. Bromine and iodine are only sparingly soluble in water but are soluble in various organic solvents such as chloroform, carbon tetrachloride, carbon disulphide and hydrocarbons to give coloured solutions.

 

Chemical Properties of group 17 elements

Oxidation states and trends in chemical reactivity of group 17 elements  

 All the halogens exhibit –1 oxidation state. However, chlorine, bromine and iodine exhibit + 1, + 3, + 5 and + 7 oxidation states also as explained below:

The higher oxidation states of chlorine, bromine and iodine are realised mainly when the halogens are in combination with the small and highly electronegative fluorine and oxygen atoms. e.g., in interhalogens, oxides and oxoacids. The oxidation states of +4 and +6 occur in the oxides and oxoacids of chlorine and bromine. The fluorine atom has no d orbitals in its valence shell and therefore cannot expand its octet. Being the most electronegative, it exhibits only –1 oxidation state.

 All the halogens are highly reactive. They react with metals and non-metals to form halides. The reactivity of the halogens decreases down the group.

The ready acceptance of an electron is the reason for the strong oxidising nature of halogens. F2 is the strongest oxidising halogen and it oxidises other halide ions in solution or even in the solid phase. In general, a halogen oxidises halide ions of higher atomic number.

F+ 2X → 2F– + X2 (X = Cl, Br or I)

Cl+ 2X → 2Cl– + X(X = Br or I)

Br+ 2I → 2Br– + I2

The decreasing oxidising ability of the halogens in aqueous solution down the group is evident from their standard electrode potentials (Table 7.8) which are dependent on the parameters indicated below:

The relative oxidising power of halogens can further be illustrated by their reactions with water. Fluorine oxidises water to oxygen whereas chlorine and bromine react with water to form corresponding hydrohalic and hypohalous acids. The reaction of iodine with water is non-spontaneous. In fact, I can be oxidised by oxygen in acidic medium; just the reverse of the reaction observed with fluorine.

 

Anomalous properties of fluorine     

 Like other elements of p-block present in second period of the periodic table, fluorine is anomalous in many properties. For example, ionisation enthalpy, electronegativity, and electrode potentials are all higher for fluorine than expected from the trends set by other halogens. Also, ionic and covalent radii, m.p. and b.p., enthalpy of bond dissociation and electron gain enthalpy are quite lower than expected. The anomalous behaviour of fluorine is due to its small size, highest electronegativity, low F-F bond dissociation enthalpy, and non availability of d orbitals in valence shell.

Most of the reactions of fluorine are exothermic (due to the small and strong bond formed by it with other elements). It forms only one oxoacid while other halogens form a number of oxoacids. Hydrogen fluoride is a liquid (b.p. 293 K) due to strong hydrogen bonding. Hydrogen bond is formed in HF due to small size and high electronegativity of fluorine. Other hydrogen halides which have bigger size and less electronegativity are gases.

 

Reactivity of group 17 elements towards hydrogen       

1.     Acidic strength: HF <HCl<HBr< HI

2.     Stability: HF >HCl>HBr> HI. This is because of decrease in bond dissociation enthalpy.

3.     Boiling point: HCl<HBr< HI < HF. HF has strong intermolecular H bonding. As the size increases van der Waals forces increases and hence boiling point increases.

4.     % Ionic character: HF >HCl>HBr> HI Dipole moment: HF >HCl>HBr> HI. Electronegativity decreases down the group.

5.     Reducing power: HF <HCl<HBr< HI

 

Reactivity of group 17 elements towards oxygen

Halogens react with oxygen to form oxides. However, it has been found that the oxides are not steady. Beside oxides, halogens also form a number of halogen oxoacids and oxoanions.

Reactivity of group 17 elements towards halogens        

1.     Binary compounds of two different halogen atoms of general formula X  are called interhalogen compounds where n = 1, 3, 5, or 7. All these are covalent compounds.

2.     Interhalogen compounds are more reactive than halogens because X-X is a more polar bond than X-X bond.

3.     All are diamagnetic.

4.     Their melting point is little higher than halogens.

5.     XX’ (CIF, BrF, BrCl, ICl, IBr, IF) (Linear shape) XX’3 () (Bent T- shape) XX’5, (square pyramidal shape) XX’7 (Pentagonal bipyramidal shape)

 

Reactivity of group 17 elements towards metals

As halogens are very reactive, they react with most of the metals instantly and form the resulting metal halides. For example, sodium reacts with chlorine gas and forms sodium chloride. This process is an exothermic one and gives out a bright yellow light and a lot of heat energy.

2Na(s) + Cl2(g) → 2NaCl(s)

Metal halides are ionic in nature. This is because of the high electronegative nature of the halogens and high electropositivity of the metals. This ionic character of the halides reduces from fluorine to iodine.

Chlorine  

Preparation:

Chlorine can be prepared by any of the following processes:

 

Properties:

•     It is a greenish yellow gas with pungent and suffocating odour.

•     It is soluble in H2O

•     Reaction with metals and non-metals: Chlorine reacts with a number of metals and non-metals to form chlorides.

      For example:

                                                2 Al + 3Cl2 → 2AlCl3

                                                S8 + 4Cl2 → 4S2Cl2

•     Reaction with ammonia: When treated with excess ammonia, chlorine gives nitrogen and ammonium chloride whereas when excess chlorine reacts with ammonia, nitrogen trichloride is formed.

                         8NH3   + 3Cl2 → 6NH4Cl + N2

                        Excess

                        NH3     +    Cl2  → NCl3 + 3 HCl

                                      Excess

•      Reaction with NaOH: Chlorine reacts differently with cold dilute NaOH and hot concentrated NaOH.

                        Cl2    +   2NaOH        →     NaCl + NaOCl + H2O

                                    Cold dil.

                        Cl2 +      6NaOH        →     NaCl + NaOCl + H2O

                                     Hot conc.

Reaction with slaked lime: Cl2 when treated With dry slaked lime it gives bleaching powder:

                                      2Ca(OH)2 + 2Cl2 → Ca(OCl)2 + CaCl2 + 2H2O

Cl2 acts as a powerful bleaching agent and its bleaching action is due to its oxidizing nature.

                                                Cl2 + H2O → 2HCl + O

Uses:

(i) Chlorine is used for bleaching woodpulp.

(ii) It is used in the extraction of gold and platinum

(iii) It is used in in sterilising drinking water.

(iv) It is used in the manufacture of dyes, drugs and organic compounds like CCl4, DDT, refrigerants, etc.

 

Hydrogen halide                            

Preparation:

It is prepared by heating sodium chloride with concentrated sulphuric acid.

                                             2 NaCl + H2SO4 + Heat → Na2SO4 + HCl

Properties:

HCl is a colourless gas with pungent odour.

It is extremely soluble in water, HCl + H2O → H3O+ + Cl

It decomposes salts of weaker acids, Na2CO3 + 2HCl → 2NaCl + H2O + CO2

When treated with NH3, it gives white fumes of NH4Cl, NH3 + HCl → NH4Cl

3HCl : 1HNO3 is called aqua regia, which is used for dissolving noble metals.

            Au + 4 H+ + NO3 + 4Cl → AuCl4 + NO + 2 H2O

Uses:

(i) Hydrogen chloride is used in medicine and as a laboratory reagent.

(ii) It is used in the manufacture of chlorine, NH4Cl and glucose.

 

Oxoacids of Halogens                  

Fluorine due to its small size and high electronegativity forms only one oxoacid HOF (Hypofluorous acid).

Other halogen form several oxoacids as given in the following table:

 

Interhalogen Compounds                      

Binary compounds of two different halogen atoms of general formula XX’n  are called interhalogen compounds where n = 1, 3, 5, or 7. All the interhalogen compounds are covalent in nature.

Some properties of interhalogen compounds are given in the following table: