P Block Elements Class 12 Notes | Part 3 | Group 17 and 18 | Chapter 7 | Chemistry |
P Block Elements Class 12 Notes (Part 3)
Group 17 of periodic table consists of five elements, i.e., fluorine, chlorine, bromine, iodine and astatine.
These elements are called Halogens.
Fluorine is most reactive and hence is also called super halogen.
Astatine is radioactive and hence occurs in nature only in traces.
[He] 2s2 2p5
[Ne] 3s2 3p5
[Ar] 3d10 4s2 4p5
[Kr] 4d10 5s2 5p5
[Xe] 4f14 5d10 6s2 6p5
1.Atomic and Ionic Radii
Question: Why Halogens have smallest atomic radii in their respective period?
Answer: This is due to the reason that atomic radii decreases as the nuclear charge increases.
Question: Why atomic and ionic radii increases down the group?
Answer: This is due to increase in number of shells as we move down the group.
Question: Why halogens have very high amount of ionization enthalpy?
Answer: Due to smaller size and greater nuclear charge, halogens have very less tendency to loose electrons. Thus, halogens have very high value of ionization enthalpy.
Question: Why ionization enthalpies decreases down the group?
Answer: On moving down the group, ionization enthalpies progressively decreases due to increase in size from F to I.
3.Electron Gain Enthalpy
Question: Why halogens have large negative electron gain enthalpies?
Answer: Halogens have one electron less than the nearest noble gas configuration. Therefore, they have a strong tendency to accept an additional electron and hence halogens have large negative electron gain enthalpies.
Question; The negative value of electron gain enthalpy is less for fluorine than for chlorine?
Answer: Due to small size the electron – electron repulsion in 2p-subshell of F are comparatively large and hence the incoming electron is not accepted with the same ease as in case with relatively bigger 3p subshell of Cl atom. Thus, negative electron gain enthalpy of F is lower than that of Cl.
Note: Amongst halogens, chlorine has the most negative electron gain enthalpy.
Question: Why halogens are highly electronegative elements?
Answer: Due to small size and higher nuclear charge, halogens have high electronegativity in their respective period.
Question: Why electronegativity decreases from fluorine to iodine?
Answer: As we move down the group from fluorine to iodine, electronegativity decreases due to corresponding increase in size of atom.
Question: Why halogens exists as diatomic molecule?
Answer: Each halogen has one electron less than the nearest noble gas and hence has a strong tendency to share its unpaired electron with another atom of the same halogen to form diatomic molecules.
Melting and Boiling Points
Question: Why F2 and Cl2 are gases while I2 is solid?
Answer: The strength of the van der waals forces increases as the size of halogen increases from fluorine to iodine. As a result, F2 and Cl2 are gases at room temperature, Br2 is a liquid whereas I2 is a solid.
Question: Why m.pt and b.pt increases down the group?
Answer: As we move down the group, van der waal force increases due to increase in molecular mass. Thus, melting and boiling points increases down the group.
Enthalpy of Dissociation
Question: F2 has lower bond dissociation enthalpy than Cl2 . Why?
Answer: Due to smaller size, the lone pair of electrons on the F atoms repel the bond pair of F—F bond. In contrast, because of comparatively larger size of Cl atoms, the lone pairs on Cl atoms do not repel the bond pair of Cl—Cl bond. As a result, F—F bond energy is lower than that of Cl—Cl bond energy.
Thus, bond dissociation enthalpy follows the sequence: Cl2 > Br2 > F2 > I2
Question: Why halogens are coloured?
Answer: The colour is due to the fact that their molecules absord light in the visible region as a result of which their electrons are excited to the higher energy levels while the remaining light is transmitted. The colour of halogen is actually the colour of this transmitted light.
Answer: The amount of energy required for the excitation decreases progressively from fluorine to iodine as size increases. Conversely, energy of transmitted light increases progressively. Fluorine has small size and force of atttraction between nucleus and electrons is high. So it requires high excitation energy and thus absorb violet light and appears pale yellow while iodine absorbs yellow and green light and hence appears deep violet.
Nature of Bonds
In order to attain noble gas configuration, a halogen atom either
gain a electron from an electropositive element forming the halide ion X¯
shares an electron with another atom to form a covalent bond as in halogen molecule.
Question: Fluorine does not exhibit any positive oxidation states. Explain.
Answer: Since F is the most electronegative atom known, therefore, the possibility of sharing its unpaired electron with an element more electronegative than itself does not arise and hence it cannot show an oxidation state of +1. Further, F does not have d orbitals in its valence shell. Therefore, it cannot expand its valence shell and hence does not show higher positive oxidation state of +3, +5 and +7. Thus, in nutshell, F cannot show positive oxidation state.
Exception: In HOF fluorine shows an oxidation state of +1.
Answer: Halogens can share their unpaired electrons with more electronegative elements and thus can also show an oxidation state of +1.
Question: Why halogens show -1 oxidation state?
Answer: All halogens show an oxidation state of -1 by accepting an electron from highly electropositive elements or by sharing an electron with less electronegative elements .
Question: Halogens are highly reactive. Explain?
Answer: Halogens have seven electrons in their respective valence shells and thus need only one more electron to complete their respective octets. Moreover they have low bond dissociation energy. That is why halogens are very reactive.
The relative oxidising power of halogens can be illustrated by their reaction with water as given below:
2F2 + 2H2O → 4H+ + 6F¯ +O2
3F2 + 3H2O → 6H+ + 6F + O3
Cl2 +H2O → HCl + HOCl (Hypochlorous acid)
Br2 + H2O → HBr + HOBr (Hypobromous acid)
Question: Among Halogens F2 is the strongest oxidising agent while I2 is weakest oxidising agent. Explain?
Answer: Since fluorine has highest standard electrode potential, it is most easily reduced. Thus it acts as strongest oxidising agent. However, I2/ I¯ has minimum standard electrode potential and thus weakest oxidising agent.
The order of oxidising power is F2 > Cl2 > Br2 > I2
Trends in Chemical Reactivity
Reactivity towards Hydrogen
The reactivity of halogens towards hydrogen decreases from fluorine to iodine.
HF and HCl are prepared by heating fluorides and chlorides with concentrated sulphuric acid.
CaF2 + H2SO4 (conc.) → CaSO4 + 2HF
2NaCl + H2SO4 (conc.) → Na2SO4 + 2HCl
HBr and HI cannot be prepared by action of conc. H2SO4 on bromide and iodide because HBr and HI are strong reducing agent and hence they reduce sulphuric acid to sulphur dioxide and are themselves also oxidised. Thus, they are prepared by heating bromide or iodide in presence of non oxidising agent.
3 NaBr + H3PO4 → Na3PO4 +3 HBr
3 NaI + H3PO4 → Na3PO4 +3 HI
Properties of Hydrides
Physical State : Hydrogen fluoride is a low boiling liquid whereas HBr, HCl and HI are gases. The anamalous behaviour of HF is due to presence of intermolecular hydrogen bonding in its molecules.
Melting points and boiling points : HF has highest m.pt and b.pt due to extensive intermolecular hydrogen bonding. However, as we move from HCl to HI, their melting and boiling points show a regular increase due to corresponding increase in the magnitude of van der waals force of attraction as the size of halogen increases.
Dipole moment : As the electronegativity of halogen decreases, the dipole moment of hydrogen halides decreases in same order, i.e., HF > HCl > HBr > HI
Acidic strength : Hydrogen fluoride is a weaker acid than hydrogen chloride in aqueous solution . Because of smaller size of F as compared to Cl, the bond dissociation energy of H—F bond is much more higher than H—Cl bond. As a result H—Cl bond can break more easily to release H+ ion than H—F bond. Thus acidic strength increases in order : HF < HCl < HBr < HI
Thermal Stability : It is directly proportional to bond dissociation energy and thus follows the order: HI < HBr < HCl < HF
Reducing Power: Greater the bond dissociation energy, more stable is the halogen acid and hence weaker is reducing agent. It increases in the order : HF < HCl < HBr < HI
Reactivity towards Oxygen (Formation of Oxides)
Oxides of Fluorine
Fluorine form two binary compounds with oxygen, i.e., OF2 and O2F2. However, OF2 is thermally stable at 298 K and O2F2 is highly unstable . They are called as oxygen fluorides rather than oxides of fluorine since the electronegativity of F is higher than that of O.
Structure: The structure of OF2 is similar to that of H2O while that of O2F2 is similar to that of H2O2 involving sp3 hybridization of O atom.
Question: Compare bond angles of H2O and OF2.
Answer: In H2O , due to greater electronegativity of O than H, the bond pair lie nearer the O atom. Consequently, bond pair – bond pair repulsion in H2O are stronger than bond pair – bond pair repulsions in OF2 and hence F—O—F bond angle is little shorter (103º) than H—O—H bond angle (104.5º).
Question: Compare bond lengths in H2O2 and O2F2.
Answer: Since F is much more electronegative than O, therefore, it attracts the lone pair of electrons on the O atom towards itself. Therefore, the lone pair – lone pair repulsion between the two O atoms in O—O bond is much lower in O2F2 than in H2O2. In other words, the O—O bond length in O2F2 is much shorter (122pm) than that in H2O2 (148pm).
Stability of oxides of halogens decreases in the order: I > Cl > Br. Further, the highr oxides of halogens tend to be more stable than lower ones.
Cl2O (Dichlorine monoxide) : The O atom is sp3- hybridized and its structure is similar to that of OF2. However, due to steric crowding of the two Cl atoms, the Cl—O—Cl bond angle is 111º.
ClO2 (Chlorine dioxide): The central atom is sp2 hybridised with O—Cl—O angle of 118º. Both the Cl—O bonds have equal (149 pm) bond lengths and are quite shorter than those in Cl2O (171 pm) Therefore, Cl—O has appreciable double bond character due to pπ—dπ bonding. The molecule is paramagnetic since it has one odd electron in a p orbital. It is sused asd bleaching agent for paper pulp , textiles and in water treatment.
Cl2O6 (Dichlorine hexaoxide) : It is a solid . It exists in equilibrium with monomer ClO3 which is a liquid. The exact structure of neither solid nor liquid is known. Both are diamagnetic and have no unpaired electrons.
Cl2O7 (Dichlorine heptoxide) : It consists of two ClO3 molecule linked by an oxygen atom.
Reactivity of Halogens toward other Halogens
Halogens react with each other to form a number of compounds called interhalogen compounds. The general formula is XX‘n where X is a less electronegative halogen (halogen of larger size) while X’ is a more electronegative halogen (halogen of smaller size ) and n is its number.
Physical State and colour
Pale brown gas
IF (very unstable)
Ruby red solid and brown red solid
Bent T- shaped
Yellow green liquid
Bent T shaped
Bent T shaped
Bent T shaped
Colourless gas but solid below 77K
Properties of Interhalogen Compounds
In interhalogen the larger halogen always serve as the central atom.
All interhalogen compounds are essentially covalent compounds because of small electronegativity difference.
All interhalogen compounds have paired electrons and hence are diamagnetic in nature.
The physical properties of interhalogen compounds are intermediate between those of constituent halogens except their melting and boiling points are little higher due to polarity associated with these molecules. Further, their melting and boiling point increases as the difference in electronegativity increases.
They are either volatile liquids or solids except ClF, BrF, ClF3 , IF5 which are gases at room temperature.
Interhalogen compounds are generally more reactive than the halogen (except F2) since X—X’ bond between two dissimilar electronegative elements is weaker than the bond between two similar atoms. This is due to reason that overlapping of orbitals of two dissimilar atoms is less effective than the overlapping of orbitals of similar atoms.
Thermal stability decreases as the size difference or electronegativity difference between the two halogen atom decreases.
Hydrolysis: ICl + H2O → HCl + HOI
They are used as non aqueous solvent.
Interhalogen compounds of fluorine is used as fluorinating agents.
Deacon’s process: In this process, hydrogen chloride gas is oxidised by atmospheric oxygen in the presence of CuCl2 as catalyst at 723 K. 4HCl + O2 → 2Cl2 + 2H2O
It is a greenish yellow gas with pungent and suffocating odour. It is about 2.5 times heavier than air. It can be liquified into greenish yellow liquid which boils at 239 K. It is soluble in water.
Reaction with metals, non metals and metalloids:
2 Na + Cl2 → 2NaCl
2 Al + 3Cl2 → 2AlCl3
2 Fe + 3 Cl2 → 2 FeCl3
P4 + 6 Cl2 → 4 PCl3
S8 + 4Cl2 → 4 S2Cl2
2 As + 3 Cl2 → 2 AsCl3
Oxidising Properties: When it is dissolved in water, it forms chlorine water. On standing, chlorine loses its yellow colour due to formation of HCl and HClO. Hypochlorous acid formed is unstable and decomposes to give nascent oxygen which is responsible for its oxidising properties.
In the preparation of poisonous gas such as phosgene, tear gas, mustard gas.
Hydrogen chloride is manufactured by salt cake method. This reaction is endothermic and is performed in two stages.
First stage: NaCl + H2SO4 → NaHSO4 + HCl
Second stage : NaCl + NaHSO4 → Na2SO4 + HCl
It is a colourless pungent smelling gas. It is easily liquified to form a colourless liquid (b.p. 189 K) and freezes to a colourless liquid (f.p. 150 K). It is extremely soluble in water and its aqueous solution is called hydrochloric acid.
Acidic Nature: It ionizes in aqueous solution
HCl + H2O → H3O+ + Cl
Reactions with salts of weak acids: HCl decomposes salts of weaker acids.
Na2CO3 + 2HCl → 2 NaCl + H2O + Cl2
NaHCO3 + HCl → NaCl + H2O + Cl2
Na2SO3 + 2HCl → 2NaCl + H2O + SO2
Reaction with metals: Active metals react with HCl to produce corresponding metal chloride with liberation of hydrogen gas.
Zn + 2HCl → ZnCl2 + H2
Mg + 2HCl → MgCl2 + H2
Fe + 2HCl → FeCl2 + H2
2Al + 6HCl → 2AlCl3 + 3H2
in dyeing, calicoprinting, tanning and sugar industry.
for extracting glue from bones and purifying bone black.
in medicine and as laboratory reagent.
for the preparation of aqua regia which is used to dissolve noble metals.
Oxoacids of Halogens
The group 18 of the periodic table consists of six monoatomic gases, i.e., helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn).
These gases do not show any chemical reactivity at ordinary temperatures and hence they were earlier called inert gases. Later, it was found that these gases are not chemically inert. Consequently, these gases are now called noble gases.
[He] 2s2 2p6
[Ne] 3s2 3p6
[Ar] 3d10 4s2 4p6
[Kr] 4d10 5s2 5p6
[Xe] 4f14 5d10 6s2 6p6
Atomic and Physical Properties
Monoatomic nature: All the noble gases are monoatomic, colourless and odourless. Their monoatomic nature is due to their stable outer electronic configuration (ns2 np6) of their atoms.
Atomic Radii: The atomic radii of noble gases are by far the largest in their respective periods. This is due to the reason that noble gases have only van der waals radii while others have covalent radii. van der Waals radii, are larger than covalent radii.
Ionisation enthalpy: The ionization enthalpies of noble gases are highest in their respective periods due to their stable gas configuration. However, as we move down the group ionisation enthalpy decreases due to increase in atomic radii and shielding effect of inner electrons.
Electron gain enthalpy: Noble gases have completely filled subshells. As a result, there is no vacant room in their valence orbitals and hence the additional electron has to be placed in an orbital of next higher shell. In other words, energy has to be supplied to add an additional electron and hence, the electron gain enthalpy of noble gases is positive.
Melting and boiling points: The melting and boiling points of noble gases are very low. This is due to the reason that the atoms of these elements are held together by weak van der waals force of attraction both in liquid as well as in solid state.
Ease of liquefaction: As the atomic size increases, the magnitude of their van der waals force of attraction increases and hence ease of liquefaction increases as we move down the group from He to Xe.
Question: What inspired N. Bartlett for carrying out reaction between Xe and PtF6?
Answer: Neil Bartlett observed that PtF6 reacts with O2 to yield an ionic solid, O2+ PtF6¯ . In this reaction, O2 gets oxidised to O2+ by PtF6¯ . Since the first ionization enthalpy of Xe (1170 KJ mol¯ ) is fairly close to that of O2 molecule (1175 KJ mol¯ ), Bartlett though that PtF6 should also oxidise Xe to Xe+. This inspired Bartlett to carry out the reaction between Xe and PtF6 . When Xe and PtF6 were mixed, a rapid reaction took place and a red solid with the formula Xe+ PtF6¯ .
Question: Why Xe form maximum compound with O2 and F2?
Answer: Atomic size of Xe is large so its ionization enthalpy is minimum. Both fluorine and oxygen being highly reactive are in position to cause electron shift from filled orbitals to vacant orbitals.
Preparation : Xenon react directly with fluorine under appropriate condition to form three binary fluorides XeF2 (xenon difluoride ), XeF4 (xenon tetrafluoride) and XeF6 (xenon hexafluoride).
All the fluorides are colourless crystalline solids and sublime readily at 298 K.
The lower fluorides form higher fluorides when heated with F2 under pressure.
Reaction with Water:
2XeF2 + 2H2O → 2Xe + 4HF + O2
XeF6 + 2H2O → XeO2F2 + 4HF
6 XeF4 + 12H2O → 4Xe + 2XeO3 + 24HF + 3O2
Question: What happens when XeF4 reacts with SbF5?
Answer: XeF4 reacts with fluoride ion acceptor such as SbF5 to form cationic species.
XeF4 + SbF5 → [XeF3]+ [SbF6]¯
Question : Xenon fluorides acts as strong fluorinating agents, Justify.
XeF4 + 2SF4 → Xe + 2SF6
XeF4 + Pt → Xe + PtF4
XeF2 + 2NO → Xe + 2NOF
Preparation: Hydrolysis of XeF4 nad XeF6 with water gives XeO3 while its partial hydrolysis give oxyfluorides.
6 XeF4 + 12H2O → 4Xe + 2XeO3 + 24HF + 3O2
6 XeF4 + 12H2O → 4Xe + 2XeO3 + 24 HF + 3O2
XeF6 + 3H2O → XeO3 + 6HF
XeF6 + H2O → XeOF4 + 2 HF
XeF6 + 2H2O → XeO2F2 + 4HF
XeO3 is colourless explosive solid and acts as powerful oxidising agent.
3 Pu2+ + XeO3 + 6 H+ → 3 Pu4+ + Xe + 3H2O
XeO4 is not as stable as XeO3 and decomposes to give xenon and oxygen.
XeO4 → Xe + 2O2
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