Group – 17 elements (Halogens)
→ The elements of this group are Fluorine (F), Chlorine (Cl), Bromine (Br), lodine (I), and Astatine (At). All the elements of this group are non-metallic in nature. Collectively these elements are known as halogens. Halogen is a Greek word (halo sea salts; gens = producing) whose meaning sea salt forming.
→ Among all the halogens fluorine is the most reactive element so it is known as super halogen. The last element Astatine of this group-17 is radioactive in nature hence it occurs only in traces. The name halogen was given by Schweigger in 1811.
Occurrence of Group – 17 elements
→ All the balogens are very reactive so they are not found in free state or native state in nature. They occur in combined state (except Astatine, as it is radioactive in nature). Halogens our na halide (X–) salt Flourine and chlorine are fairly abundant in the earth’s crust while bromine and lodine are lesser abundant, Fluorine is present mainly as insoluble fuorides. Por example Cryolite (Na3AlF6), Fluorspar (CAF2).
Fluorapatite CaF2 . 3Ca3 (PO4)2 or Ca5(PO4)3 F] etc.
→ Except above compounds, mall quantities of fluorine is also present in bones, soil, plants, river water and teeth of animals. Minor currence of Durine are found in Topse (Al2SiO4(OH)F2), Sellaite (MgF2), Villaumite Nat) etc.
→ Pluorine is thirteenth most abundant element of the earth’s crust Chlorine cours as chlorides to the extent of about 0.14% in earth’s crust. The main source of chlorines water (2.8% by ma x epte water it also present in salt well and salt beds.
→ It cours in Sodium Chloride (NaCl), Camellite (KCl . MgCl2 . 6H2O and Calcium Chloride (CaCl2) Bromine is found in bromocarnallite (KBr MgBr2 . 6H2O) and also in se water in the form of bromides lodine mainly occurs in sea weeds (0.5% by m e ) alkali metal iodides, cliche (crude chile salt peter) which is mainly wodium nitrate containing (0.2% by mas) of iodine s sodium iodate (NaIO3).
The relative abundance of halogens are given in table 7.15.
Table 7.15 Relative abundance of halogens in earth’s crust.
|Element||Abundance in earh’s crust by mass (ppm)||Relative abundance|
Physical properties of Group – 17 elements
The physical properties of group 17 elements are given in table 7.16
Table 7.16 Atomic and physical properties of atomie halogen
- Pauling scale
Electronie Configuration : The general electronic configuration of halogen is ns2 np5:
Table 7.17 Electronic configuration of halogens
→ Physical State : All the halogen elements occur in the form of non polar, covalent dintomic molecules. Their physical states are as follows –
→ The difference in physical state is due to difference in the extent of van der Waals forces of attraction which depend upon the atomic wine and atomie weight. Since the size and atomie weights of urine and chlorine are less bence they exist in the form of gas. Because of larger size and larger weight, bromine occurs as liquid and iodine as solid (An sine increase, van der Waal’s forces of attraction also increase).
→ Melting and boiling point : On moving down the group atomic se increases hence the strength of van der Waal’s forces of attraction increases the melting and boiling points also increase m.p and b pa van der Waal’s force atomic weight F2 < Cl2 < Br2 < I2 (m.p and b.p)
→ Atomic and lonie radil. In their respective periods halogens have the smallest atomie rudis. Because on moving left to right in period nuclear charge increases hence atomic radius decreases On moving down the coup atomic size increases because number of shells increases The size of halide ion is always greater than the size of halogena because with the addition of one electron the effective nuclear chance decreann hence the size increases.
→ Ionisation Enthalpy Halogen have very high ionisation energy due to small atomic size. So they have very very little tendency to lose electrons. On moving down the roup lonisation energy decreases because atomic size increases.
F > C > Br > I(Tonisation enthalpy)
→ The value of first lonisation energy of the elements of group-17 is higher than the elements of group-16 because on moving left to right in period, effective unclear charge incrematomie akze decrer hence the value of ionisation enthalpy also increases.
→ Electronegativity: Due to small size and higher nuclear charge, the electronegativity of the elements of group-17 is highest. On moving down the group electronegativity decreases. Fluorine has the highest electronegativity (40) in periodic table.
F > CI > Br > I (Element)
4.0 > 30 > 2.8 > 2.5 (Electronegativity)
→ Electron gain enthalpy or Electron affinity: On moving down the group the electron guin enthalpy first increases from fluorine to chlorine then decreases
F < C > Br > I > At(Electron main enthalpy)
→ The low value of electron gain enthalpy of fluorine is due to its small size i.e., having high electron density which makes the addition of electron somewhat difficult. But halogens have maximum electron gain enthalpy in their respective periods. It is due to its electronic configuration ng up which is one electron less than the stable noble gas configuration (n ap). Due to this reason halogens have the maximum tendency to accept electron.
→ Non-metallic character: All the elements of group-17 are non metala. On moving down the non metallic character dece b ut metallic character increases. Non-metallic character of the elements of group-17 is due to their high loniantion energy and high electronegativities. Lodine has some metallic character in it. For example:
- Iodine is solid.
- Iodine has metallic luature
Iodine has maximum tendency in group-17 to make eation
Colour: All the halogens are coloured. The colours of halogens are given below:
→ They show colour due to the fact that their molecules absorb light in the visible region. Due to the absorption of light they get excited and jump to higher energy level while the remaining light is transmitted This transmitted light is men in the form of colour.
→ Actually the halogens show the complementary colour. As we have seen that the colour becomes deeper and deeper as the atomic number increes from fluorine to iodine. This change in colour is called as blue shift or bathochromie shift.
→ When light falls on halogens, fluorine absorbs voilet portion of the light higher exeitation energy and thus appears as pale yellow while iodine absorbs yellow and green portion of light (lower excitation energy) and thus appears as voilet.
→ Bond energy On moving down the group the bond dissociation energy decreases because the bond length increases – bond dissociation energy is smaller than
→ Cl – Cl and even smaller than that of Br – Br. It is due to this reason that fluorine has very small size and have greater electron-electron repulsion between lope pair of electron.
→ Thus, the order of bond dissociation energy of group-17 is as follow –
Cl2 < Br2 > F2 > I2 (Bond dissociation energy)
The value of bond energies of halogens are as follows:
|Bond length (pm)||143||199||228||266|
|Bond dissociation enthalpy’ (k J/mol)||158,8||242.6||192.8||151.1|
Anomalous Behaviour of Fluorine :
→ Fluorine is the first member of group-17. It differs from other members of this group. This differences are due to the following reasons:
- It has very small size
- It has very high electronegativity
- It does not have any vacant d-orbitals in its valence shell
- Its molecular form (F2) has very low bond dissociation enthalpy.
→ The main points of differences between fluorine and other halogens are as follows:
- It show only one oxidation state i.e.-1 in their compounds while other halogens can show positive as well as negative oxidation states.
- It has very low band dissociation enthalpy hence fluorine (F2) has the maximum reactivity in halepen family
- Fluorine can form hexafluoride with sulphur (SF6) while other member of this family cannot form hexahalide with sulphur.
- HF is liquid at room temperature while other halogen acids are gases.
→ Fluorine form only one oxo seid i.e., HOF Thybofluorous acid) while other members of halogens can form a number of oxo acide ie HXO, HXO2, HXO3, HXO4.
→ Due to the presence of tendency to form hydrogen band, HF extat in Ancinted form while other halogen acids exist as monomers.
→ Fluoride ions can form complex ions as [AIF6]3-], [FeF6]3- etc. while other halide ions show this complex formation tendency in smaller extent.
Due to the absence of Vacant d-orbitals it cannot form polvfluoride ions while other members can form.
For example: Cl3–, Br3–, I3–, I5– etc.
Fluorine can react with water and form a mixture of oxygen and osone. But other members do not form oxygen and ozone.
Only fuorine ean forma jonic fluorides while others form covalent halides.
For example: AIF3, SnF4 are ionic while SnCl4 is covalent.
HF used for etching of glass it can react with glass while other halogen acids do not show this reactions.
Due to above special properties of fluorine it is called as super halogen.
Chemical Properties of Group – 17 elements (Halogens)
→ Halogens are higly reactive elements due to their high electronegativity, high negative electron gain entalphy values and low bond discelation energies. Fluorine is the most reactive element among group-17. The important chemical properties of group-17 are given below
→ Nature of bonds. To attain noble gas electronic configuration, a halogen atom either gain one electron from an electropositive element or share one electron with another atom. When it gain an electron then it converts into habide ion and it forms ionic bond with electropositive in or when it share one electron, it forms covalent bond. Thus, we can say that halogenean form lonie as well as covalent compounds.
- The halides of highly electropositive metals are one while the halides of weakly electroportive metals or non-metals are convalent in nature. Some examples are given below;
- Ionic Compounds : NaCI, KCI, KI, CaCl2, MgCl2 etc.
- Covalent Compounde: HCl, AlCl3, CCl4, SiCl4, BiCl3, SbCl3 etc.
→ Oxidation State: All the elements of group-17 have tendency to acquire noble gas electronie configuration Por getting this configuration they either accept one electron to form monovalent anion or share one electron to form covalent compound.
→ Thus in this way they can show an oxidation state of -1 0r +1 depending on the nature of combining element and when the element combined with halogens is less electronegative then halogen sbow -1 oxidation state. When the element is more electronegative, they show +l oxidation state.
→ Since fluorine is the highest electronegative element among the periodie table hence it always shows only – oxidation state. Other halogers can also show higher positive oxidation states like +3, +5, +7 due to presence of empty L-orbitals.
→ As fluorine do not have empty d orbitals so it does not show higher oxidation tates. The higher cucidation state shown by chlorine, bromine sod iodine is due to the presence of empty orbitals where outer por pelectrons can easily be promoted as shown below:
(Oxidation state = +7) (Seven unpaired electron hence O.S = +7)
→ Solubility : Halogen molecules are non polar in nature hence they are insoluble in polar solwenta like water, alcohols etc. Eut fluronie reacts with cold water in dark and form the mixture of oxygen and one as given below.
→ Cl2, Br2 and I2 are highly soluble in organic solvents like CHCl3, CCl4, CS2 and hydrocarbon etc. They dissolves in CCI, and CHCl, and give yellow, brown and purple colour respectively.
→ Chlorine and bromine dissolves in water in presence of sunlight and decomposes water. The reactivity of Br2 is lesser than chlorine.
- X2 + H2O → HX + HXO (Where X-CI, & Br.) i.e.,
- Cl2 + H2O → HCl + HCIO
- Br2 + H2O → HBr + HBrO
→ Oxidising Power : All the halogens have a strong tendency to accept an electron due to high elecgronegativity and high negative value of electron gain enthalpy, hence they act as strong oxidising agent. The oxidising power of halogens decreases from F2 to I2 is the strongest oxidising agent in group-17. F, On oxidise all the other halide ions into halagens. The involved reactions are as follow:
- F2 + 2X– → 2 + X2 (X– = Cl–, Br–, I–)
- Cl2 + 2X– → 2Cl– + X2 (X– = Br–,I–)
- Br2 + 2l– → 2Br– + I2
→ It indicates that iodine is the weakest oxidising agent amongst all the halogens. In general, we can say that a baloren of lower atomie weight can oxidise the holide ions of higher atomie weight. The oxidising power of halogens is related to the standard reduction potentinl(E) values. Higher the Evalue, higher will be the oxidising power. E’ value of halogens are as follow :
Hence, the oxidising power is as follows:
- F2 > Cl2 Br2 >I2 (Oxidising power) In contrast,
- I– > Br– > Cl– > F– (Reducing power)
The relative oxidising power can further be explained with the help of the following reactions.
→ The bond diaxatic energy of F2(g) (79.1 kJ mol-1) is smaller than the bond dissociation exergy of Cl2(g) (122 kJ mol-1)The electron gain enthalpy of Cl (-394 kJ mol-1) is higher than the electron gain enthalpy of F (-333 kj.mol-1).
→ The hydration energy ∆hyd H– of F– ion(- 460 kJ mol is higher than Clion (-343 kJ mol-) which is due to small size and high charge density of fluoride (F–) ion. On calculating the net energy released, we can calculate the value of standard reduction potential.
→ Reactivity towards hydrogen (Formation of hydrogen halide)! All the halogense react with hydrogen to form hydrogen halide.
H2 + X2 → 2HX
→ The reactivity of halogen towards hydrogen decreases down the group. For example: fluorine reacts with hydrogen violently even in the dark and cold atmosphere.
Hence, the order of activity of halogens with hydrogen is
F2 > Cl2 > Br2 > I2 (Reactivity order).
Properties of hydrogen halide : Some important characteristics of hydrogen halides are as follow:
→ Physical state: Except HF all the other halides like HCl, HBr, HI are ses at room temperature. He is liquid at room temperature with low boling point (b.p → 292K). HF A liquid at room temperature due to the presence of H-bonding.
→ Melting and boiling points: The boiling point of hydrogen fluoride is higher than the boiling points of the other hydrogen halide. It is due to the presence of intermolecular H-bonding.
→ On moving down the group the boiling point increases regularly from HCI to HI because as molecular weight increases the extent of van der Waal’s forces of attraction increase hence, boiling point also increased.
HCl < HBr < HI < HF (Boiling point)
→ Similarly, the melting point of HF is higher than other hydrogen halides Arain on moving down the group the melting point also increases gradually from HCI to HI due to increase in extent of ander waala interaction.
HCl < HBr < HI < HF (Melting point)
→ Volatility : The volatility of halide decreases down the group from HCI to HI because the extent of van der Waals interaction increases. But the HF is least volatile due to the presence of hydrogen bonding
HCl > HBr > HI > HF (Volatility)
→ Thermal stability On moving down the group the thermal stability decreas beenuse the size of halogens increases due to which the bond length increases, hence bond energy decreases.
HF > HCI > HBr > H (Thermal stability)
→ Acidic behaviour : In gaseous state hydrogen halides do not behave as acid because they are covalent in nature but in aqueous state they can ionise and behave se acid by donatine H ion. The acidic strength of these hydrogen halide incorases down the group because sine incron ,bondenergy decree, hence the tendency to donate proton (H+) increases.
HF < HCl < HBr < HI (Acidic strength)
Hence HF in the wenkeat acid while He is the strongest acid.
→ Reducing character : On moving down the group the reducing behaviour incerns because the bond dissociation energy decreases.
HF < HCl < HBr < HI (Reducing Nature)
→ Do you know: Only HF has tendency to react with glass. This tendency of HF is used in etching of glass.
All the properties of hydrogen halides are summearised in tahle given below:
Table 7.18 : Properties of hvdeogen halides
|Bond length (H – X)/pm||91.7||127.4||141.4||160.9|
→ Reactivity towards oxygen : Oxygen and halogen react with each other directly or indirectly to form a lot of compounds. Some of them are as follows:
→ Oxides of fuorine: Fluorine forms two compounds with oxygen. They are OF2 and O2F2. As the electronegativity of fluorine is higher than oyren bence they are called as fluorides not oxides.
Only OF2 stable at 296 K temperature. These oxygen fluorides can be prepared as follows:
→ Both the compounds ie. OF2 and O2F2 are strong fluorinating agents. O2F2 can oxidise PuF6 and the reaction is used in removing plutonium as Puf from spent nuclear fuel.
Oxides of other halogens : Oxides of other halogens are given below
Table 7.19: Oxides of halogens
→ From the above table, it is clear that all halogens react with oxygen and form various oxides in which their oxidation states range from +1 to +7. Chlorine forms a large number of oxides while iodine form least number of carides.
→ These oxides are binary in nature. They form covalent bonds due to small difference in their electronegativities i.e., in between oxygen and halogen It is quite interesting that the stability of oxides of iodine is greater than those of chlorine while bromine oxides are least stable lodine-oxygen bond is stable due to greater polarity of the bond.
→ Chlorine oxygen bond is stable due to multiple bond formation in which d-orbitals of chlorine atom involve. But bromine being in between lacks both characteristics ie, neither show greater bond polarity nor show multiple bond formation.
Thus, the stability of oxides of halogens show the following order
I > C > Br
→ Oxides of chlorine i.e., Cl2O, ClO2, Cl2 O6, and Cl2O7 are highly reactive and behave as oxidising agents. They have tendency to explode. ClO2 is used as a bleaching agent for paper pulp, textiles and in water treatment.
→ The oxidea of bromine le, Br2O, BrO2 and BrO3 are least stable halogen and t hey exist only at low temperature. They are powerful oxidising agents.
→ The oxides of iodine are I2O4, I2O5 and I2O7. Out of these only I2O5 is the most stable. These oxides decompose on heating. They behave as good oxidising agent and are used in estimating the carbon mono oxide Because it oxidises Co to CO2 quantitatively liberating iodine which can be titrated against sodium thiosulphate.
I2O5 + 5CO → I2 + 5CO2
→ Reactivity towards Metals All the halogens reacts with almost all metals and non-metals except He, Ne, Ar to form binary eompounds. These compounds are known as halides. These halides may either be simple or complex molecules which differ from each other only in their structures and stoichiometries.
→ Halogens reacts with highly electropositive metals like NA, K, Ca, ete, to form ionle halides. These sonic halides have high melting and boiling points. The ionic character of these haldes decreased down the group because electronegativity decreases.
M – F > M – Cl > M – Br > M – I (Tonic character)
→ Halogens reacts with three metals which have very high ionisation energy like Sn, Pb, Sb, ete, to form covalent halides. If we compare the covalent character of these halidea for similar metal having different oxidation atates then the holides of metala with high oxidation state are more covalent than the halides of metals with lower oxidation state. For example:
- SnCl4 > SnCl2 (Covalent character)
- SbCl5 > SbCl3 (Covalent character)
- PbCl4 > PbCl2 (Covalent character)
→ Halogen reacts with non-metals and form various covalent halides like PFs, PCB, PBrs. Asig.8, Cloete. For a single halogen the covalent character increases with increase in oxidation number of metal. For example
NaCl < MgCl2 < AlCl3 < ThCl2 (Covalent character) The above order is according to Fajan’s Rule.