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The P-Block Elements

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Published in: Chemistry
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Concept Of Class -XII

Neha M / Chandigarh

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Qualification: M.Sc (Kurukshetra University , Kurukshetra - 2011)

Teaches: Chemistry, Physics, Mathematics, Science, B.Tech Tuition, IIT JEE Advanced

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  1. 1. 2. 3. 4. 5. 6. 8. 9. 1 Class XII: Chemistry Chapter 7: The p-BIock Elements Top Concepts p-BIock elements: Elements belonging to groups 13 to 18 of the periodic table are called p-block elements. General electronic configuration of p-block elements: The p- 2 1-6 block elements are characterized by the ns np valence shell electronic configuration. Representative elements: Elements belonging to the s and p-blocks in the periodic table are called the representative elements or main group elements. Inert pair effect: The tendency of ns electron pair to participate in bond formation decreases with the increase in atomic size. Within a group the higher oxidation state becomes less stable with respect to the lower oxidation state as the atomic number increases. This trend is called 'inert pair effect'. In other words, the energy required to unpair the electrons is more than energy released in the formation of two additional bonds. Nitrogen family: The elements of group 15 nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi) belong to configuration is ns np Oxygen family: Group 16 of periodic table consists of five elements - oxygen (O), sulphur (S), selenium (Se), tellurium (Te) and polonium (PO). Their general electronic configuration is ns np The halogen family: Group 17 elements, fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At), belong to halogen family. Their general electronic configuration is ns np Group 18 elements: Helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn) are Group 18 elements. They are also called noble gases. Their general electronic configuration is ns np except helium which has electronic configuration Is . They are called noble gases because they show very low chemical reactivity. GROUP 15 ELEMENTS Atomic and ionic radii: Covalent and ionic radii increase down the group. There is appreciable increase in covalent radii from N to P. There is small increase from As to Bi due to presence of completely filled d or f orbitals in heavy elements.
  2. 10. 11. 12. 13. 14. 2 Ionisation energy: It goes on decreasing down the group due to increase in atomic size. Group 15 elements have higher ionisation energy than group 14 elements due to smaller size of group 15 elements. Group 15 elements have higher ionization energy than group 16 elements because they have stable electronic configuration i.e., half filled p-orbitals. Allotropy: All elements of Group 15 except nitrogen show allotropy. Catenation: Nitrogen shows catenation to some extent due to triple bond but phosphorus shows catenation to maximum extent. The tendency to show catenation decreases down the group. Oxidation states: The common oxidation states are +3, +5, -3. The tendency to show -3 oxidation state decreases down the group due to decrease in electronegativity which is due to increase in atomic size. The stability of +5 oxidation state decreases whereas stability of +3 oxidation state increases due to inert pair effect. Nitrogen shows oxidation states from -3 to +5. Nitrogen and phosphorus with oxidation states from +1 to +4 undergo oxidation as well as reduction in acidic medium. This process is called disproportionation. 3 HN02 HN03 + H20 +2 NO Reactivity towards hydrogen: All group 15 elements from 3 trihydrides, MH3. Hybridisation - sp The stability of hydrides decrease down the group due to decrease in bond dissociation energy down the group. NH3 > PI-13 > AsH3 > SbH3 > BiH3 Boiling point: PH3 < AsH3 < NH3 < SbH3 < BiH3 Boiling point increases with increase in size due to increase in van der Waals forces. Boiling point of NH3 is more because of hydrogen bonding. Bond angle: NH3 (107.80) > PI-13 (99.50) > AsH3 (91.80) SbH3 (91.30) > BiH3 (900) Electronegativity of N is highest. Therefore, the lone pairs will be towards nitrogen and hence more repulsion between bond pairs. Therefore bond angle is the highest. After nitrogen, the electronegativity decreases down the group. Basicity decreases as NH3 > PH3 > AsH3 > SbH3 < BiH3.
  3. 15. 16. sp hybridisation , pyramidal shape Pentahalides - sp d hybridisation, T BP shape They are lewis acids because of the presence of vacant d - orbitals. PC15 + Cl- [PC16]- PC15 is ionic in solid state and exist as [PC14]+ [PC16]- In PC15, there are three equatorial bonds and two axial bonds. The axial bonds are longer than equatorial bonds because of greater repulsion from equatorial bonds. Nitrogen does not form pentahalides due to absence of d- orbitals. Reactivity towards metals: All elements react with metals to form 17. binary compounds in -3 oxidation state. Anomalous behaviour of nitrogen: The behaviour of nitrogen 18. differs from rest of the elements. Reason : It has a small size. It does not have d - orbitals ii . 3 This is because the lone pair of electrons are concentrated more on nitrogen and hence the basicity will be maximum in the case of NH3. It will decrease down the group as the electronegativity decreases down the group. The reducing power of hydrides increases down the group due to decrease in bond dissociation energy down the group. Reactivity towards oxygen: All group 15 elements from trioxides (M203) and pentoxides (M205). Acidic character of oxides decreases and basicity increases down the group. This is because the size of nitrogen is very small. It has a strong positive field in a very small area. Therefore, it attracts the electrons of water's O-H bond to itself and release H ions easily. As we move down the group, the atomic size increases. Hence, the acidic character of oxides decreases and basicity increases as we move down the group. Reactivity towards halogen: Group 15 elements form trihalides and pentahalides. Trihalides - covalent compounds and become ionic down the group.
  4. iii. iv. 4 It has high electronegativity It has high ionization enthalpy 19. 20. 21. Dinitrogen: Preparation: heat NH4 Cl(aq) + NaN02 (aq) N2 (g) + 21-12 0(1) + NaCI(aq) heat (NH 4 ) 2 C r 2 07 N2 + 41-12 + 03 Properties: It is a colouless, odourless, tasteless and non - toxic gas. It is chemically un-reactive at ordinary temperature due to triple bond in N N which has high bond dissociation energy. Ammonia: Ammonia molecule is trigonal pyramidal with nitrogen 3 atom at the apex. It has 3 bond pairs and 1 lone pair. N is sp hybridised. Preparation: Haber's process: N2 (g) + 31-12 (g) YZZZZZZX2 NH3 (g) Pr essure =200x105 Pa Temperature -773 K fHO - 46.1 kJ mol-I Catalyst is Feo with small amounts of K20 and A1203 Nitric Acid: a. Ostwald Process: 4NH3 + 502 Pt /Rhgauge — 4N0 + 61-12 0 500 K,9bar 2NO + 02 YZZZZZZZX2N02 3N02 (g) + 1-12 0(1) N03 (aq) + NO(g) ......(iii) NO thus formed is recycled and the aqueous HN03 can be concentrated by distillation upto 68% by mass. Further
  5. 5 concentration to 98% can be achieved by dehydration with concentrated H2S04. Nitric acid is strong oxidizing agent in the concentrated as well as in the dilute state. Phosphorus: 22. Has an iron grey luster a. b. It shows the property of catenation to maximum extent due to most stable P- P bond. It has many allotropes, the important ones are i. White phosphorus ii. Red phosphorus iii. Black phosphorus White phosphorus Discrete tetrahedral P4 molecules Very reactive Glows in dark Translucent waxy solid Soluble in CS2 but insoluble in water It has low ignition temperature, therefore, kept under water Red phosphorus Polymeric structure consisting of chains of P4 units linked together Less reactive than white P Does not glow in dark Insoluble in water as well as Black phosphorus Exists in two forms - a black P and ß black P Very less reactive Has an opaque monoclinic or rhombohedral crystals
  6. Prepa ration : White P High 573 K in an inert atmosphere for several days 6 Red p pressure 473 K Black P 23. Phosphine: Preparation : Ca3P2 calcium phosphide 61-12 0 water calcium hydroxide 803 K In a sealed tube + 2PH3 phosphine 3CaC12 + 2PH3 ( phosphine ca3 P2 + 6HCI ii. + 3NaOH + 3H20 3NaH2PO 2 sodium hypophosphite ( phosphine) Phosphine is highly poisonous, colourless gas and has a smell of rotten fish .
  7. 24. Chlorides of Phosphorous: PC13 Colourless oily liquid P4 + 6C12 QC13 P4 + 8soc12 -+ 4PC13 + 4S02 + 2S2 C12 7 PC15 Yellowish white owder P4 + IOC12 4pC15 P4 + IOS02 c12 + IOS02 hydrolysed in moisture PC13 + 31-120 3 CH3COOH + PC13 H3P03 the Of H3P03 + 3HCl 3 cH3coa + PCI + H O 5 2 POC13 + 3 1-120 3 CH3COOH + poc13 + HCI poa + 21-lcl 3 -9 H3P04 +3HCl PC15 CH3COCl + 3 C2H50H + PCI 3 3 C2H5Cl + H3P03 Pyramidal shape, sp hybridisation GROUP 16 ELEMENTS C2H50H + PC15 C2H5Cl + poc13 + HCI T BP geometry, sp d hybridisation 2Ag + pc15 2 Aga + pc13 Sn + 2 pc15 SnC14 + 2 pc13 25. 26. Oxidation states: They show -2, +2, +4, +6 oxidation states. Oxygen does not show +6 oxidation state due to absence of d orbitals. PO does not show +6 oxidation state due to inert pair effect. The stability of -2 oxidation state decreases down the group due to increase in atomic size and decrease in electronegativity. Oxygen shows -2 oxidation state in general except in OF2 and 02F2 The stability of +6 oxidation state decreases and +4 oxidation state increases due to inert pair effect. Ionisation enthalpy: Ionisation enthalpy of elements of group 16 is lower than group 15 due to half filled p-orbitals in group 15 which are more stable. However, ionization enthalpy decreases down the group.
  8. 27. 28. 29. 30. 8 Electron gain enthalpy: Oxygen has less negative electron gain enthalpy than S because of small size of O. From S to PO electron gain enthalpy becomes less negative to PO because of increase in atomic size. Melting and boiling point: It increases with increase in atomic number. Oxygen has much lower melting and boiling points than sulphur because oxygen is diatomic (02) and sulphur is octatomic (S8). Reactivity with hydrogen: All group 16 elements hydrides. Bent shape form Bond angle: H20 > H2S < H2Se < H2Te 13K 232K 269 373 Intermolecular increase in van der Waals forces H bonding Acidic nature: H20 < H2S < H2Se < H2Te This is because the H-E bond length increases down the group. Therefore, the bond dissociation enthalpy decreases down the group. Thermal stability: 1-120 < < < H2Te < H2P0 This is because the H-E bond length increases down the group. Therefore, the bond dissociation enthalpy decreases down the group. Reducing character: 1-120 < < < H2Te < H2P0 This is because the H-E bond length increases down the group. Therefore, the bond dissociation enthalpy decreases down the group. Reactivity with oxygen: E02 and E03 Reducing character of dioxides decreases down the group because oxygen has a strong positive field which attracts the hydroxyl group and removal of H becomes easy. Acidity also decreases down the group. S02 is a gas whereas Se02 is solid. This is because Se02 has a chain polymeric structure whereas S02 forms discrete units.
  9. 31. 32. 33. 9 Reactivity with halogens: EX2 EX4 and EX6 The stability of halides decreases in the order F - > Cl- > Br- > I This is because E-X bond length increases with increase in size. Among hexa halides, fluorides are the most stable because of steric reasons. Dihalides are sp hybridised, are tetrahedral in shape. Hexafluorides are only stable halides which are gaseous and have sp3d2 hybridisation and octahedral structure. H20 is a liquid while H2S is a gas. This is because strong hydrogen bonding is present in water. This is due to small size and high electronegativity of O. Oxygen: Preparation: heat 2KC103 2 + 30 Mn02 finely divided metals 2H2 0(l) + 02(g) 21-12 02(aq) 2HgO(s) 2Pb304(s) (Red lead) 2Pb02 (s ) Oxides: heat 4Ag(s ) + 02(g) 2Hg(I) + 02 (g) 6pbO(S ) + 02(g) 2PbO(s) + 02(g) The compounds of oxygen and other elements are called oxides. Types of oxides: a. b. Acidic oxides: Non- metallic oxides are usually acidic in nature. S02 + H20 —Y H2S03 (sulphurous acid) Basic oxides: Metallic oxides are mostly basic in nature. Basic oxides dissolve in water forming bases e.g.,
  10. c. d. 10 Na20 + 1-120 2NaOH 1
  11. 36. 37. 38. 39. 40. 41. Sulphuric acid: Preparation: By contact process 1 —8 S 8 + 02 SO 2 V205 2S02 ( g) + 02 ( g) 2S03 ( g) 2bar 720K 11 HO --196.6kJ mol - 1 Exothermic reaction and therfore low temperature and high pressure are favourable SO 3 (g) + H 2 S04 1—12 S 2 07 (Oleum) H2S207 + H20 2 H2S04 (96 -98%) It is dibasic acid or diprotic acid. It is a strong dehydrating agent. It is a moderately strong oxidizing agent. GROUP 17 ELEMENTS Atomic and ionic radii: Halogens have the smallest atomic radii in their respective periods because of maximum effective nuclear charge. Ionisation enthalpy: They have very high ionization enthalpy because of small size as compared to other groups. Electron gain enthalpy: Halogens have maximum negative electron gain enthalpy because these elements have only one electron less than stable noble gas configuration. Electron gain enthalpy becomes less negative down the group because atomic size increases down the group. Eelctronegativity: These elements are highly electronegative and electronegativity decreases down the group. They have high effective nuclear charge. Bond dissociation enthalpy: Bond dissociation enthalpy follows the order C12 > Br2 > F2 > 12 This is because as the size increases bond length increases.
  12. 42. 43. 44. 45. 46. 12 Bond dissociation enthalpy of C12 is more than F2 because there are large electronic repulsions of lone pairs present in F2. Colour: All halogens are coloured because of absorption of radiations in visible region which results in the excitation of outer electrons to higher energy levels. Oxidising power: All halogens are strong oxidisinga gents because they have a strong tendency to accept electrons. Order of oxidizing power is F2 > C12 > Br2 > 12 Reactivity with H2: Acidic strength: HF < HCI < HBr < HI Stability: I-IF > Ha > HBr > HI This is because of decrease in bond dissociation enthalpy. Boiling point: I-ICI < HBr < HI < I-IF HF has strong intermolecular H bonding As the size increases van der Waals forces increases and hence boiling point increases. % Ionic character: HF > HCI > HBr > HI Dipole moment: > HCI > HBr > HI Electronegativity decreases down the group. Reducing power: HF < HCI < HBr < HI Reactivity with metals: Halogens react with metals to form halides. Ionic character: MF > MCI > MBr > Ml Halides in higher oxidation state will be more covalent than the one in the lower oxidation state. Interhalogen compounds: Reactivity of halogens towards other halogens: Binary compounds of two different halogen atoms of general formula X X'n are called interhalogen compounds where n These are covalent compounds. All these are covalent compounds.
  13. 47. 48. 49. 50. 51. 52. 53. 13 Interhalogen compounds are more reactive than halogens because X- X' is a more polar bond than X-X bond. All are diamagnetic. Their melting point is little higher than halogens. XX' (CIF, BrF, Bra, ICI, 1Br, IF) (Linear shape) (CIF3, Bra, IF3, ICE) (Bent T- shape) XX'5 - CIF5, BrF5, IF5, (square pyramidal shape) XX'7 - IF7 (Pentagonal bipyramidal shape) Oxoacids of halogens: Fluorine forms only one oxoacid HOF (Fluoric (I) acid or hypofluorous acid) due to high electronegativity. Acid strength: HOCI < HC102 < HC103 < HC104 Reason: HC104—H + 004- most stable Acid strength: HOF > I-loci > HOBr > HOI This is because Fluorine is most electronegative. GROUP 18 ELEMENTS: Ionisation enthalpy: They have very high ionization enthalpy because of completely filled orbitals. Ionisation enthalpy decreases down the group because of increase in size. Atomic radii: Increases down the group because number of shells increases down the group. Electron gain enthalpy: They have large electron gain enthalpy because of stable electronic configuration. Melting and boiling point: Low melting and boiling point because only weak dispersion forces are present. XeF2 is linear, XeF4 is square planar and XeF6 is distorted octahedral. KrF2 is known but no true compound of He Ne and Ar are known. Compounds of Xe and F:
  14. 2 Xe + 3F 14 673 K, Ibar XeF2 873 K xeF4 7bar 573 K 60 -70 bar XeF4 + 02 F 2 + 02 XeF2, XeF4 and XeF6 are powerful fluorinating agents. 54. Compounds of Xe and O: 6 XeF4 + 121-120 —4Xe + 2Xe03 + 24HF XeF6 + 3 1-120 Xe03 + 6HF 302

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