Classification of Elements and Periodicity in Properties chem.pdf
- 1. Welcome to Classification of Elements and Periodicity in Properties
- 2. Why do we need to classify? How do you find a book of your interest in a library? How difficult would it be if all the books were mixed up? We know that when books are organised in the library, it is very easy to identify the book but it may be a tedious job when they are randomly placed.
- 3. Why do we need to classify? 1…………….…………………………...………………………………………………118 ….………… Systematic way of organizing knowledge by classifying elements. Properties of elements repeat themselves after regular intervals
- 4. Developments in Periodic Classification Newlands’ Law of Octaves Mendeleev Periodic Table Modern Periodic Table Dobereiner’s Triads
- 5. Dobereiner’s Law of Triads Triads Arrangement of the elements in increasing order of their atomic mass Group of three elements possessing similar properties Element Atomic Mass K 39 Li 7 Na 23 7 + 39 2 = 23 Mean atomic mass of 1st & 3rd element Na Li K This rule seemed to work only for a few elements. So, it was dismissed.
- 6. Newlands’ Law of Octaves Arrangement of the elements in increasing order of their atomic masses Properties of every eighth element are similar to the first element
- 7. Sa Re Ga Ma Pa Dha Ni Element Li Be B C N O F Atomic mass 7 9 11 12 14 16 19 Element Na Mg Al Si P S Cl Atomic mass 23 24 27 29 31 32 35.5 Element K Ca Atomic mass 39 40 Newlands’ Law of Octaves Drawback of Newlands’ Octaves Applicable only upto Ca
- 8. Mendeleev's Periodic Table Elements with similar properties occupy same vertical column Arranged 63 known elements in increasing order of atomic weight in a tabular form Periodic table Vertical column (8) Group Horizontal rows (7) Period
- 9. Mendeleev's Periodic Law Physical and chemical properties of elements are a periodic function of their atomic masses Predicted some undiscovered elements and their properties Eka-Aluminium (Gallium) & Eka-Silicon (Germanium) Atomic mass of Be was corrected Considering its valency 2 (2 × 4.5 = 9)
- 10. Drawbacks of Mendeleev’s Periodic Table Hydrogen did not have a fixed position No regular trend in increasing order of atomic mass Isotopes failed to comply with the periodic table
- 11. Moseley’s Modification Bombarded high speed electrons on different metal surfaces Plot of √𝜈 (frequency) v/s Z (atomic number) gave a straight line √𝜈 a (Z − b) = a, b are costants
- 12. Modern Periodic Law Physical and chemical properties of elements are a periodic function of their atomic number
- 13. Periodicity Repetition of properties of elements after regular intervals when elements are arranged in order of increasing atomic number. Periodic law is the consequence of the periodic variation in electronic configuration.
- 15. Modern Periodic Table Elements belonging to same period have same valence shell. Periods Horizontal rows (7) Elements of same group have similar valence shell electronic configuration. Group Vertical column (18)
- 16. Blocks of the Periodic Table s-block f-block p-block d-block Based on the type of orbitals which receives the differentiating electron ● H is placed outside the periodic table though it has last electron in s-orbital. ● He belongs to s-block but placed in p-block with noble gases. Exceptions:
- 17. s-block
- 18. s-block s-block Last electron enters the ‘ns’ subshell Outer electronic configuration ns1-2 Groups 1-2 & all are metals
- 19. s-block Group 1 Alkali metals Group 2 Alkaline earth metals Outer electronic configuration ns1 Outer electronic configuration ns2
- 20. p-block
- 21. p-block p-block Last electron enters the ‘np’ subshell Outer electronic configuration ns2 np1-6 Groups 13-18 & includes some metals, all nonmetals and metalloids
- 22. p-block 13 14 15 16 Group 17 18 Boron family Carbon family Pnictogens Chalcogens Name Halogens Noble gases
- 23. d-block
- 24. d-block d-block Last electron enters the (n - 1)d subshell Outer electronic configuration (n-1) d1-10 ns0-2 Groups are from 3 to 12 & all are metals
- 25. d-block 3-d series 4-d series Sc(Z = 21) - Zn(Z = 30) Y(Z = 39) - Cd(Z = 48) 5-d series 6-d series La(Z = 57) - Hg(Z = 80) Ac(Z = 89) - Cn(Z = 112) Transition series Elements 1st 2nd 3rd 4th Forms a bridge between s-block (chemically active) & groups 13 & 14 (less active) Transition elements
- 26. d-block Zn, Cd and Hg are not transition elements because they do not have partially filled d-orbitals in neutral or in any stable oxidation state
- 28. f-block f-block Last electron enters the ‘(n - 2)f’ subshell Outer electronic configuration (n-2) f1-14 (n-1) d0–1 ns2 Group 3 & all are metals
- 29. 2nd Inner transition series Actinides Lanthanides 1st Inner transition series Th(Z = 90) - Lr(Z = 103) Ce(Z = 58) - Lu(Z = 71) f-block f-block
- 30. f-block Elements after Uranium in the periodic table are called Transuranic elements
- 31. Main group elements Elements of groups 1, 2 and 13–18 (except H) Typical elements 2nd period elements except Ar Remember! Coinage metals Group 11 elements are Cu, Ag & Au
- 32. Remember! Representative elements s-block elements along with p-block elements Transition elements Elements which have partially filled d-orbitals in neutral state or in any stable oxidation state Inner transition elements Elements in which all the three shells i.e. n, (n– 1) & (n – 2) shells are incomplete
- 33. Remember! 2 Atomic number of Noble Gases 10 18 36 54 86 He Ne Ar Kr Xe Rn
- 34. How are elements named in periodic table? May be on the name of the scientist? May be based on the names of planets, Greek mythology…? May be on the place of discovery? May be on the meaning of elements with its behavior?
- 35. How Elements Got their Names? Name of Elements Place of discovery Behaviour Planet names Scientist names Americium Hydrogen Mercury Bohrium
- 36. Roots are put together in order of digits from 0 to 9 which make the atomic number “ium” is added at the end Corresponding symbol has three letters IUPAC Nomenclature for Elements Having Z > 100
- 37. Digit 0 Name nil 1 un 9 enn Abbreviation n u e 2 bi 3 tri b t 4 quad 5 pent q p 6 hex 7 sept h s 8 oct o IUPAC Nomenclature for Elements Having Z > 100 1 un 0 nil 5 pent unp Z = 105 unnilpentium
- 38. Subshell in which the last electron enters is the block Eg: 1s2 2s2 2p4 ⇒ p - block Write the electronic configuration Finding Position of an element in periodic table Find block
- 39. s - block Number of electrons in the outer s-orbital is its group number p - block (Number of electrons in outermost shell) + 10 = group number Find group Finding Position of an element in periodic table
- 40. d - block Number of electrons in outermost shell + penultimate d subshell f - block Group 3 Find group Finding Position of an element in periodic table
- 41. Z ≥ 103 Z = 100 + x , x = group number = 3 Example: 6 Z = 106 Group Number Atomic number Finding Position of an element in periodic table Find group when Z ≥ 100
- 42. Highest value of Principal Quantum Number (n) Write the electronic configuration Find period Finding Position of an element in periodic table
- 43. On the basis of nature Metals Metalloids Non Metals Classification of elements
- 44. Metals Non-metals and Metalloids Metals Non - Metals Metalloids
- 45. Classification of elements Metal Good conductor Lose electron(s) easily High lustre Ductile & Malleable
- 46. Non-metal Poor conductors Gain electron(s) easily Brittle Classification of elements
- 48. Penetration The proximity of electrons in an orbital to the nucleus s > p > d > f Order of Penetration power Consequence of Penetration effect Shielding Effect
- 49. Comparison of Penetration Power Measure of orbital’s closeness to the nucleus r 𝚿 R 2 4𝜋r 2 3p 3s 3d
- 50. Shielding Effect Outer electrons experience an attractive force from the nucleus & repulsive forces from the other electrons +3 Repulsion Attraction
- 51. Shielding Effect Attraction between the outer electron & the nucleus decreases In a multi-electron system Due to the repulsions with the inner electrons Effect in which the inner electrons shield the valence electron(s) from the attraction of the nucleus s > p > d > f Strength of Shielding effect
- 52. Effective Nuclear charge (Zeff ) Shielding decreases the nuclear attraction Refers to the actual charge experienced by the outermost electrons Z = Atomic number 𝛔 = Screening or Shielding constant Zeff Z = 𝛔 −
- 53. Effective Nuclear charge Zeff increases from left to right across a period Li < 1.3 Be < B < C < N < O < F < Ne 1.95 2.60 3.25 3.90 4.55 5.20 5.85 For s block elements Zeff increases slightly and becomes constant down a group Li < Na ≃ K ≃ Rb ≃ Cs 1.3 2.2 2.2 2.2 2.2
- 54. Trends Across the Periodic Table 01 03 04 Electron gain enthalpy 02 Electro- negativity 05 Atomic radius Ionization enthalpy Electron affinity
- 55. Atomic Radius Distance between the centre of the nucleus and the outermost electron(s) of an atom Finding the size of an atom is a lot more complicated Size of an atom is very small No sharp boundary exists
- 56. Atomic Radius Size of an atom can be measured using the internuclear distance between two identical atoms by X-ray diffraction or other spectroscopic techniques. Atomic radii Covalent radii Metallic radii van der Waals radii
- 57. Covalent radii rcov = d d 2 Generally used for non-metals Units - picometer (pm) or Angstrom (Å) Half the distance between the centres of two nuclei of the same element bonded by a single covalent bond
- 58. rA dA-B = rB + 0.09 Δ − 𝜒 ● 𝜒 = Electronegativity (Pauling scale) rA = dA-A 2 (in Å) rB = dB-B 2 (in Å) Schomaker & Stevenson equation Covalent radii of Heteronuclear Diatomic Molecule ● ●
- 59. Metallic radii d rmetallic = d 2 Half the internuclear distance between two adjacent metallic atoms in a crystalline lattice structure
- 60. van der Waals Radii van der Waals radius For Noble gas For Non-metals
- 61. For Noble gases, van der Waals Radii Half the distance between the nuclei of two non-bonded nearest neighbouring atoms of the same element in its solid state rvdw d
- 62. Half the distance between the nuclei of two non-bonded nearest neighbouring atoms of two adjacent molecules of the same element in solid state For Non-metals, ( C l 2 m o l e c u l e ) 3 6 0 p m d (Cl2 molecule) 99 pm van der Waals Radii rvdw = 180 pm rcovalent = 99 pm
- 63. Comparison of Metallic and Covalent Radii Metallic radius rNa rCl Covalent radius van der Waals radii Metallic radii Covalent radii Overlapping Radii > >
- 64. Number of shells Atomic Radius Example ∝ Cs > Rb > K > Na > Li Number of shells Factors affecting Atomic radii
- 65. Atomic Radius 1 Effective nuclear charge ∝ Li > Be > B > C > N > O > F Factors affecting Atomic radii Effective nuclear charge Example
- 66. Screening effect Atomic Radius Example ∝ Cs > Rb > K > Na > Li Factors affecting Atomic radii Screening effect
- 67. Atomic Radius Example ∝ Factors affecting Atomic radii N N N N N N > > 1 Number of Bonds Number of bonds
- 68. Variation of Atomic Radii in a Group Atomic radius Principal Quantum number (n) Distance between valence electrons & nucleus
- 69. Variation of Atomic Radii in a Group 300 250 200 150 100 50 Atomic number (Z) Atomic radius (pm) Li(152) Na(186) K(231) Rb(244) Cs(262) F(72) Cl(99) Br(114) I(133) Alkali Metals Halogens
- 70. Atomic Radius Atomic number , Zeff Attraction between valence electrons & nucleus Variation of Atomic Radii in a Period
- 71. Variation of Atomic Radii in a Period 2nd Period 2 4 6 8 10 60 80 100 120 140 160 Atomic number (Z) Atomic radius (pm) Li B Be N F C O
- 72. Atomic radius increases (in a period from left to right) Atomic Radii Atomic radius increases down the group
- 73. Irregularities in Variation of Atomic radii Generally van der Waals radius is larger than the covalent & metallic radius Atomic radius of an inert gas (18th group) is the largest in its period Represented by van der Waals radius (monatomic state) Due to the poor shielding effect of d electrons Atomic Radius of Ga Atomic Radius of Al >
- 74. Radii of the elements decrease and then become constant 1st Transition series Near the end, radius increases Atomic Radii trends in d block Element Sc Ti V Cr Mn Fe Co Ni Cu Zn Atomic radius (pm) 144 132 122 117 117 117 116 115 117 125 Sc - Cr Cr - Cu Cu - Zn Zeff dominates Zeff ≃ Electronic repulsion High electronic repulsion
- 75. 4f subshell electrons exert very poor shielding effect and cancel out the normal size increase on moving down a group Lanthanoid Contraction 14 lanthanide elements Lanthanum Hafnium 5d series
- 76. Exceptions in d-block elements Due to Lanthanoid contraction Atomic radii of d block elements : 3d < 4d ≅ 5d But Sc Y La < < 4f electrons are absent in La
- 77. Covalent Radii of Transition Metals 1.8 1.7 1.6 1.5 1.4 1.3 1.2 1.1 4 (Ti) 5 (V) 6 (Cr) 7 (Mn) 8 (Fe) 9 (Co) 10 (Ni) 1 2 3 Group Number (1st row element) Covalent Radius (Å) R o w N u m b e r
- 79. Ionic Radius Cation Smaller than its parent atom Less electrons than its parent atom More Zeff e p electron proton < 1
- 80. Ionic Radius Anion Larger than its parent atom More electrons than its parent atom Less Zeff e p electron proton > 1
- 81. Atom Cation Atom Anion Ionic Radii (pm) Li 152 Na 186 K 231 Rb 244 Cs 262 Li+ 76 Na+ 95 K+ 138 Rb+ 146 Cs+ 174 F 72 Cl 99 Br 114 I 133 F- 136 Cl- 184 Br- 196 l- 216
- 82. Isoelectronic species Atoms and ions which contain the same number of electrons Zeff For isoelectronic species Ionic radii Al3+ Mg2+ Na+ F- O2- N3- Ionic radii increases as Zeff decreases
- 83. Li+ Mg2+ ≈ > For isotopes Size of anion Size of atom Size of cation > > For same element Remember !!! 35 37 rCl rCl
- 84. Smallest anion is F− and not H− Order of radii: F− < Cl− < Br− < H− < I− H− : e/p ratio = 2 As e/p ratio is very high for H− , its size is abnormally bigger. Did you know?
- 85. Ionization Energy Minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state M (g) M+ (g) + e- Unit - kJ/mol, kcal/mol, eV
- 86. Ionization Energy Quantitative measure of the tendency of an element to lose an electron Easier it is for a neutral atom to change itself to a positive ion Smaller is the ionization energy Ionization energy for an atom is always positive
- 87. First Ionization Energy Energy required to remove the outermost electron from a gaseous neutral atom M (g) M+ (g) + e- , IE1
- 88. Second Ionization Energy Energy required to remove the second most loosely bound electron M+ (g) M2+ (g) + e- , IE2
- 89. Third Ionization Energy Energy required to remove the third most loosely bound electron M2+ (g) M3+ (g) + e- , IE3
- 90. I.E.1 < I.E.2 < I.E.3 ... Removal of an electron from a cation is more difficult than from a neutral gaseous atom 2nd I.E. > 1st I.E. Therefore Which is Larger - First, Second or Third I.E.?
- 91. Successive Ionization Energies Element Electronic Configuration I.E.1 (kJ/mol) I.E.2 (kJ/mol) I.E.3 (kJ/mol) Li 1s2 2s1 520 7298 11815 Be 1s2 2s2 899 1757 14850 B 1s2 2s2 2p1 801 2427 3660 C 1s2 2s2 2p2 1086 2353 4620 N 1s2 2s2 2p3 1402 2856 4578 O 1s2 2s2 2p4 1314 3388 5300 F 1s2 2s2 2p5 1681 3374 6050 Ne 1s2 2s2 2p6 2080 3952 6122
- 92. I.E. Factors affecting Ionization Energy 1 Size of an atom 2 Effective nuclear charge ( Zeff ) I.E. 3 Screening effect I.E. 4 Penetration Effect I.E. 5 Electronic Configuration Fully filled > Half filled > Partially filled Order of I.E.
- 93. Ionization Energy in a Period Across a period Zeff I.E.
- 94. Ionization Energy in a Group Down the group Size I.E.
- 95. Ionization Energy Ionization energy increases Ionization energy increases
- 96. IE1 of the first 60 elements 10 20 30 40 50 60 500 1000 1500 2000 2500 Ne He Ar Li Na K Kr Rb Xe Cs Atomic number (Z) Δ 1 H (kJ / mol)
- 97. Li B < Be C O N F Ne Order of First Ionization Energy 2nd Period Element < < < < < < Zeff in a period Increases from left to right I.E.
- 98. Order of I.E. for 2nd Period Elements 1st I.E. 2nd I.E. Be C < B N F O Ne Li < < < < < < Li B < Be C O N F Ne < < < < < <
- 99. For 13th group Order of First Ionization Energy Lanthanoid Contraction Poor shielding of d electron In Al < Ga Tl < < < B
- 100. Points to Remember I.E. of non-metals > I.E. of metals Noble gases have the highest I.E. values in any period Caesium has the lowest I.E., hence it is used in photoelectric cells Helium has the highest I.E.
- 101. Metallic Character Metallic character I.E. Metallic character means how easily an element lose electrons in chemical reactions
- 102. Ionization energy increases Metallic Character and Ionization Energy Metallic character increases
- 103. Non-Metallic Character Accept electrons during chemical reactions Non-Metallic Character Non-Metallic character I.E.
- 104. Metallic character decreases Period 3 Group 5A P e ri o d s Metallic character increases
- 105. Electron Gain Enthalpy (Δ H) X (g) + e- X- (g) kJ/mol Enthalpy change when an electron is added to a neutral gaseous atom to convert it into a gaseous anion eg
- 106. Electron Gain Enthalpy Electron Gain Enthalpy Positive Negative When energy is absorbed When energy is released
- 107. Factors affecting Electron Gain Enthalpy Zeff Half filled or fully filled subshell Atomic size Magnitude of electron gain enthalpy Δeg H Screening effect
- 108. Electron Gain Enthalpy Across a Period Across a period Zeff Magnitude of Δeg H
- 109. Ne Be < N < Li < B < C < O < F < Trend in 2nd Period Fully filled subshell Half filled subshell
- 110. Electron Gain Enthalpy Down the Group Down the group Size Magnitude of Δeg H
- 111. Electron Gain Enthalpy Down a Group Cl > F S > O P > N Si > C Al > B Magnitudes of Δeg H of 3rd period elements are greater than corresponding 2nd period ‘p’ block elements It happens because 2nd period elements (F and O) have smallest size in their respective group and addition of extra electron causes greater repulsion as compared to 2nd period elements.
- 112. Why Such Exceptions…? Cl has the highest negative Δeg H value It is easier to add an electron to Cl and S due to its larger size which can accommodate extra electron more easily compared to 3rd period elements. Δeg H of noble gases is positive because it already have stable electronic configuration and after addition of an electron, it becomes comparatively unstable.
- 113. Magnitude of Electron Gain Enthalpy trends Δeg H increases Δ eg H increases
- 114. Successive Electron Gain Enthalpy Δeg H for the addition of a second electron to a neutral atom is positive Electron repulsion outweighs the nuclear attraction O (g) + e− → O− (g) ; Δeg H1 = − 141 kJ/mol O− (g) + e− → O2− (g) ; Δeg H2 = + 744 kJ/mol Second and subsequent electron gain enthalpies are always positive.
- 115. Electron Affinity (E.A.) Conventionally defined as the energy released when an electron is added to the valence shell of a neutral isolated gaseous atom X (g) + e- X- (g) eV/atom, kJ/mol
- 116. Electron Affinity (E.A.) Δeg H = − E.A. − RT 5 2 ● R : Universal Gas Constant ● T : Temperature (K) Δeg H = − E.A. For comparison purpose in Periodic Trends, we use both the terms interchangeably At 0 K
- 117. Electron Affinity (E.A.) Same as negative Δeg H Order of E.A. S Se > Te > Po > O > Order of negative Δeg H or E.A. in oxygen family
- 118. Remember !!! The smallest value of I.E. of an element (Cs) is even greater than the highest value of E.A. of an element (Cl)
- 119. Electronegativity (E.N. or 𝜒) Measure of the tendency of an atom to attract the shared electrons towards itself in a covalently bonded molecule
- 120. E.N. is not a property of an isolated atom Electronegativity 1 E.N. of an atom may be different in different molecules 2 E.N. is not a measurable property 3
- 121. Electronegativity Scales of E.N. Pauling scale Mulliken-Jaffe scale Allred-Rochow scale
- 122. Pauling scale is the most commonly used scale for E.N. Assigned value 4.0 F Highest E.N. Electronegativity
- 123. Electronegativity values (Pauling scale)
- 124. Factors Affecting Electronegativity Distance between the nucleus & the valence shell electrons Force of attraction b/w the nucleus & the valence shell electrons Electronegativity value
- 125. Factors Affecting Electronegativity Zeff Electronegativity value Force of attraction b/w the nucleus & the valence shell electrons
- 126. Factors Affecting Electronegativity Magnitude of positive charge on the atom Electronegativity value Force of attraction b/w the nucleus & the valence shell electrons
- 127. Electronegativity across a period Li Be B C N O F Elements 1.0 1.5 2.0 2.5 3.0 3.5 4.0 E.N. (Pauling Scale) From left to right in a period, Atomic radius , Electronegativity
- 128. Electronegativity down a group F Cl Br I 4.0 3.0 2.8 2.5 Elements (Group 17) Electronegativity value Down the group Atomic radius Electronegativity
- 129. F has the highest electronegativity 1 Cs has the lowest electronegativity 2 3 4 Alkali metals have the lowest E.N. in their respective period Halogens have the highest E.N. in their respective period Some important points
- 130. Electronegativity Strong tendency of non-metallic elements to gain electrons E.N. is directly related to the non-metallic properties of elements
- 131. Cannot be measured experimentally Can be experimentally measured Electron Gain Enthalpy Electronegativity Tendency of an atom in a molecule to attract the shared pair of electrons Tendency of an isolated gaseous atom to attract an electron Electron Gain Enthalpy vs Electronegativity
- 132. Electronegativity Ionization Energy Atomic Radius Electron Affinity Increases Increases Decreases Increases Decreases Increases Decreases Decreases
- 133. Number of electrons in the outermost shell Valence of representative elements 8 − (Number of electrons in the outermost shell) OR Valence
- 134. The term oxidation state is frequently used for valence Oxidation state of an element in a molecule or an ion (Imaginary) charge the atom would have, if the electron in each bond were located on the more electronegative atom Oxidation State (O.S.)
- 135. Na2 O O.S. of Na is + 1 O.S. of O is -2 O is more E.N. than Na Oxidation State (O.S.)
- 136. OF2 O.S. of F is −1 O.S. of O is +2 F is more E.N. than O Oxidation State (O.S.)
- 137. What is Valency? 1 1 1 2 2 2 13 3 1,3 14 4 2,4 Group No. of Valence Electrons Valency 15 5 3,5 16 6 2,4,6 17 7 1,3,5,7 18 8 0,8
- 138. Variation along a period Number of valence electrons from left to right increases from 1 to 8
- 139. Variation within a group Down a group, the number of valence electrons remains the same All the elements in a group exhibit the same valency
- 140. Valency and O.S. Alkali Metal Elements Valency of 1 Oxidation State of +1
- 141. Periodic Trends in Valency 1 LiH 2 CaH2 NaH KH Formula of hydride B2 H6 AlH3 Group 13 Li2 O K2 O Na2 O Formula of oxide MgO CaO SrO BaO B2 O3 Al2 O3 Ga2 O3 In2 O3
- 142. Periodic Trends in Valency 14 15 Group 16 Formula of hydride NH3 PH3 AsH3 SbH3 H2 O H2 S H2 Se H2 Te CH4 SiH4 GeH4 SnH4 Formula of oxide N2 O3 , N2 O5 P4 O6 , P4 O10 As2 O3 , As2 O5 Sb2 O3 , Sb2 O5 CO2 SiO2 GeO2 SnO2 Bi2 O3 PbO2 SO3 SeO3 TeO3
- 143. Periodic Trends and Chemical Reactivity Δeg H increases from left to right in a period I.E. increase from left to right in a period High chemical reactivity at the two extremes( left side due to less I.E. and right side due to high electron affinity)
- 144. Acidic oxides Oxides of non-metals React with bases Basic oxides Oxides of metals React with acids Oxides
- 145. Oxides Amphoteric oxides Reacts with acids as well as bases BeO, Al2 O3 , ZnO, Cr2 O3 Neutral oxides Does not react with an acid or a base CO, N2 O, NO
- 146. Oxides Oxide formed by the element on the extreme left is the most basic (Na2 O) Oxide formed by the element on the extreme right is the most acidic (Cl2 O7 )
- 147. Anomalous behaviour Absence of d orbitals (Maximum covalency = 4) Small size & high charge/radius ratio in their respective groups Highest electronegativity in their respective groups Why Anomalous behaviour for 2nd period elements?
- 148. Diagonal Relationship Li & Be show exceptional behaviour from other elements of their respective groups Form compounds with pronounced covalent character Other members of their groups predominantly form ionic compounds Li & Be are more similar to second element of their following group, i.e., Mg & Al respectively. This is called diagonal relationship in periodic properties.